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CHEM& 140 Workbook

A Collection of Worksheets and Laboratory Experiments 2019 Edition | Bellevue College Chemistry Department Faculty

CHEM& 140 WORKBOOK A Collection of Worksheets and Laboratory Experiments Bellevue College Chemistry Department Faculty 2019 Edition

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The information and material contained in this manual are provided “as is,” and without warranty of any kind, expressed or implied, including without limitation any warranty concerning accuracy, adequacy, or completeness of such information. Neither the authors, the publisher nor any copyright holder shall be responsible for any claims, attributable errors, omissions, or other inaccuracies contained in this manual. Nor shall they be liable for damages of any type, including but not limited to, direct, indirect, special, incidental, or consequential damages arising out of or relating to the use of such material or information. These experiments are designed to be used in college or university level laboratory courses, and should not be conducted unless there is an appropriate level of supervision, safety training, personal protective equipment and other safety facilities available for users. The publisher and authors believe that the lab experiments described in this manual, when conducted in conformity with the safety precautions described herein and according to appropriate laboratory safety procedures, are reasonably safe for students for whom this manual is directed. Nonetheless, many of the experiments are accompanied by some degree of risk, including human error, the failure or misuse of laboratory or electrical equipment, mis-measurement, spills of chemicals, and exposure to sharp objects, heat, blood, body fluids or other liquids. The publisher and authors disclaim any liability arising from such risks in connection with any of the experiments in the manual. Any users of this manual assume responsibility and risk associated with conducting any of the experiments set forth herein. If students have questions or problems with materials, procedures, or instructions on any experiment, they should always ask their instructor for immediate help before proceeding.

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III

CHEM& 140 WORKBOOK TABLE OF CONTENTS

CHEM& 140 Student Learning Objectives . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Scientific Method: Observation, Hypothesis, Law, or Theory? . . . . . . . . . . . . . . . . . . . . . . . . . 6

Element Symbols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7

Chemistry Matters! . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8

Molecular Level Diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15

Elements, Compounds & Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16

Scientific Notation: Powers of 10 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19

Exponents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 20

Significant Figures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21

Measurements and Significant Figures 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22

Measurements and Significant Figures 2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23

Significant Figures and Scientific Notation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 24

Significant Figures in Calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25

List of Metric Prefixes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 27

TABLE OF CONTENTS

Metric System Units & Definitions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28

Two Methods for Metric Conversions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30

Metric Conversions: Dimensional Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31

One-Step Metric Conversions Using Dimensional Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . 32

Two-Step Metric Conversions Using Dimensional Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . 33

Metric System: The Ladder Method . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34

Metric Conversions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35

Multi-Step Conversions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36

Volume Conversions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38

Ratios as Conversion Factors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39

The Jabberwocky and Googs per Mulm . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

Converting Both Numerator & Denominator . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42

Squared or Cubed Unit Conversions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43

Units Raised to a Power . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 45

Density . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46

Activity: The Nuclear Atom . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 48

Periodic Table Search . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52

Atomic Theory, Periodic Table, and Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 53

Atomic Structure: Isotopes and Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 56

Isotope Calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57

Atomic Structure Review . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58

Forming Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 60

Bohr Model Diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61

TABLE OF CONTENTS

Charges of Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 62

Tutorial: Ionic Nomenclature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 64

Polyatomic Ions to Memorize . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 66

Memorizing Polyatomic Ions: The T43 Method . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 67

Deriving Polyatomic Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68

Combining Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 70

Ionic Compound Matrix . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 71

Ionic Nomenclature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72

Nomenclature Practice Test 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74

Covalent Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 75

Covalent Nomenclature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 76

Ionic or Molecular? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 77

The Mole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 78

Mass, Moles, Avogadro’s Number, and Moles Within Moles Conversions . . . . . . . . . . . . . 79

Molar Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

Types of Substances . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 81

Moles, Grams, Number of Particles, Elements in Compounds . . . . . . . . . . . . . . . . . . . . . . . . 82

Moles: Calculate the Missing Quantities . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 84

Mole Ratios Using Chemical Formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85

Practicing Mole Ratios . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 86

Moles within Moles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 87

Percent Composition by Mass. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 89

Empirical and Molecular Formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

TABLE OF CONTENTS

Review: Moles and Chemical Composition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91

Review: Nomenclature and Chemical Composition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 93

Balancing Chemical Equations 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 94

Balancing Chemical Equations 2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95

Balancing Word Equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 96

Tutorial: Classifying Types of Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97

Stoichiometry Practice . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 98

Stoichiometry 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 99

Stoichiometry Calculations Using Balanced Equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 101

Stoichiometry: Iron Ore Smelting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102

Stoichiometry: Removing CO2 in Space . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103

Limiting Reactant by Moles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 104

Activity: Limiting Reactant . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106

Theoretical Yield and Limiting Reagents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 108

Percent Yield . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 109

Review: Stoichiometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 110

Molarity 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112

Molarity 2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113

Molarity: Applications . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114

Molarity: Fill-in-the-Blanks Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 115

Concentration: Mass Percent (% m/m) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116

VII

TABLE OF CONTENTS

Percent Concentration Practice . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117

Solution Stoichiometry 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118

Solution Stoichiometry 2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120

Activity: Electromagnetic Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121

Energy, Wavelength, and Frequency . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 123

Activity: The Bohr Model . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 125

Atomic Orbital Energy Levels . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 129

Electron Configurations 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 130

Electron Configurations 2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132

Electron Configurations 3 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 133

Orbital Diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 134

Review: Electron Configurations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 135

Electron Configurations: Puzzle . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 136

Lewis Dot Structures of Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 137

Tutorial: Lewis Dot Structures of Covalent Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 139

Lewis Dot Structures 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 140

VSEPR / Molecular Shape . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 141

Lewis Dot Structures, Geometry, and Shape . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 142

Review: Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 144

Laboratory Equipment: Names and Pictures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 146

VIII

TABLE OF CONTENTS

Experiment 1: Rubber Band Blast! . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 149

Experiment 2: Measurements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 155

Measurements Practice/Reference Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 159

Experiment 3: Density . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 165

Experiment 4: Chemical Changes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 175

Experiment 5: Empirical Formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 183

Experiment 6: Chemical Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 189

Experiment 7: Synthesis of Alum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 197

Experiment 8: Stoichiometric Analysis of an Antacid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 205

Experiment 9: Titration of Vinegar . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 215

Experiment 10: Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 227

Experiment 11: Molecular Model Building . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 239

Experiment 12: Separation of a Ternary Mixture . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 245

CHEM& 140 STudENT LEARNINg OBjECTIVES

CHEM& 140 Student Learning Objectives

SECTIONS FROM CHEMISTRY LIBRETEXTS, A FREE ONLINE RESOuRCE (CHEMISTRY TEXTBOOK) https://chem.libretexts.org/Textbook_Maps/Introductory_Chemistry/map%3A_Introductory_ Chemistry_(Tro)

unit 1: The Science of Chemistry (Pages 6–18)

1. Define chemistry as a science and explain how it is useful and relevant to everyday life. 1.1

2. Define and distinguish the major components of the “scientific approach to knowledge”: observations, hypothesis, experiments, scientific law, and theory. 1.3–1.4

3. Distinguish between what is considered science and what is a belief. Workbook

4. State and apply the law of conservation of mass. 3.7

5. Memorize chemical symbols for elements with atomic numbers 1 through 56 and 78 through 84 and locate them on the periodic table.

6. Define matter, atom, molecule, pure substance, element, compound, mixture, heterogeneous mixture, and homogeneous mixture, and classify a given substance as an element/compound/heterogeneous or homogeneous mixture based on its chemical composition.

3.2, 3.4

7. Interpret and draw atomic/molecular scale drawings. Workbook

8. Explain what is meant by physical state of matter and describe the differences among a solid, liquid, and gas in terms of particle position and strength of attraction.

3.3

9. Determine whether changes to matter are chemical or physical in nature, and which properties of matter are chemical versus physical. 3.5–3.6

10. Write the chemical formula for a compound given the number of atoms, and given the formula of a compound, state the number of atoms (including the use of parentheses in chemical formulas).

5.3

unit 2: Measurements and units (Pages 19–47)

11. Use scientific notation for a given number and apply this in multiplication, division, addition, subtraction, powers and roots. Review exponent rules (10a × 10b = 10a+b, etc).

2.2

12. Distinguish between exact and measured numbers; count the number of significant figures in a measured number, record measurements using proper number of significant figures, andreport correct number of significant figures in single and multi-step calculations, including temperature conversions—do not memorize °C/°F conversion.

2.1, 2.3–2.4

13. Describe the difference between accuracy and precision and calculate percent error. (Workbook, Lab)

CHEM& 140 STudENT LEARNINg OBjECTIVES

14. State the base units of measurements in SI and metric system. Memorize the metric prefixes: tera, giga, mega, kilo, deci, centi, milli, micro, nano, and pico, and their relationship to base unit and apply them in metric conversions.

2.5

15. Apply dimensional analysis by identifying and using conversion factors to solve single-step and multi-step calculation problems, including density. 2.6–2.9

16. Perform square and cubic unit conversions (units raised to a power). Interconvert area/volumes involving metric (memorized prefixes) and English system (provided). Memorize 1 mL = 1 cm3 = 1 cc.

2.8; Workbook;Lab

unit 3: Atoms and Elements (Pages 48–61)

17. Identify the main points of Dalton’s atomic theory (early 1800s) and some of the ideas and natural laws leading up to it. Explain what is wrong with Dalton’s atomic theory according to our current knowledge of atomic structure.

4.2

18. Explain the difference between Thomson’s plum pudding model of the atom and Rutherford’s nuclear model. Describe the observations which led to Rutherford’s conclusions about atomic structure.

4.3

19. State the relative masses and charges of the subatomic particles: protons, neutrons, and electrons. Explain where most of the atomic mass is located in an atom.

4.4

20. Use the atomic number, mass number, and charge given the AZ X notation of isotopes to determine the number of protons, neutrons, and electrons in neutral atoms.

4.5–4.6, 4.8

21. For ions, explain the difference between atoms versus ions, and cations versus anions, in reference to the subatomic particles. 4.7

22. Identify regions in periodic table as metals, nonmetals, semimetals (metalloids), main group (representative) elements, transition metals (elements), inner transition metals or elements (lanthanides and actinides), and also by families/groups in periodic table—alkali metals, alkaline earth metals, halogens, and noble gases.

4.6

23. List the common properties of metals, nonmetals, and metalloids. 4.6

24. Calculate average atomic mass for an element from the given natural abundance and mass of each isotope of that element. For elements with two or more naturally occurring isotopes, calculate the natural abundance of each isotope or isotope masses from provided data.

4.9

unit 4: Molecules and Compounds (Pages 62–79)

25. Explain this statement by providing specific examples: When elements combine to form compounds, an entirely new substance is created with properties different than the constituent elements.

5.1

26. Use the Bohr model to predict the charge of ions for ionic compounds. Contrast with covalent compounds which share electrons rather than form ions. Workbook

27. Identify the basic unit of an element or compound as atoms, molecules, formula units, or molecules. Given a chemical formula, determine if ionic bonding or covalent bonding is present.

5.4

CHEM& 140 STudENT LEARNINg OBjECTIVES

28. Predict the correct formula for an ionic compound. Write the name of binary ionic compound from a given formula that has fixed charge metals and variable charge metals (using Roman numerals), and vice versa, and for ionic compounds containing polyatomic ions. Memorize seven required polyatomic ions (sulfate, nitrate, carbonate, chlorate, phosphate, hydroxide and ammonium) and derive others based on these.

5.5–5.7

29. Write the binary molecular compound name using Greek prefixes (memorize 1 to 10) if given the formula, and vice versa. 5.8

30. Given a chemical formula, classify it as an ionic or a covalent (molecular) compound and name it (or given the name, provide its chemical formula). 5.10

unit 5: The Mole and Chemical Composition (Pages 80–95)

31. Explain the concept of the mole in chemistry and why it is useful. 6.1–6.3

32. Distinguish between formula masses (in amu) and molar masses (in g/mol).

33. Convert grams to moles (and vice versa) and moles to number of particles such as atoms, molecules, formula units (and vice versa), using molar mass and Avogadro’s number as conversion factors.

6.3–6.4

34. Given a chemical formula, identify possible mole ratios and use the ratio as a conversion factor. 6.5

35. Given a chemical formula, calculate the mass percent of each element in the compound. Use mass percent as a conversion factor to convert between mass of an element and mass of a compound.

6.6–6.7

36. Define the empirical formula and molecular formula, and determine the empirical formula if the molecular formula is given. Determine the empirical formula of an unknown compound from mass or mass percent data.

6.8

unit 6: Chemical Reactions (Pages 96–100)

37. Describe or identify observations that indicate a chemical change (reaction) has occurred. 7.2

38. Distinguish between reactants and products, and the meaning of the chemical arrow. Write and balance chemical equations. Translate sentences describing a chemical reaction into a balanced chemical equation. Apply the abbreviations for the four physical states in the balanced equations: solid, liquid, gas, aqueous.

7.3–7.4

39. Classify reactions as synthesis (combination), decomposition, single/double replacement, and combustion. 7.10

40. Define combustion reaction, and write the chemical equation for the combustion of a fuel (given its formula).

7.9, end of section only

unit 7: Stoichiometry (Pages 101–114)

41. Explain what coefficients in chemical equations represent on a molecular level and macroscopic (molar) scale. Use balanced chemical equations to write out mole ratios, and use them to determine moles of a substance reacted or formed when moles of another substance are known, i.e., moles A → moles B.

8.2–8.3

42. Apply the concept of mole ratios in a balanced chemical equation and use molar mass conversions to determine the theoretical yield of a product when the mass of any one reacant is known, assuming the other reactant(s) is in excess. g A → moles A → moles B → g B

8.4

CHEM& 140 STudENT LEARNINg OBjECTIVES

43. Define limiting reactant and determine the limiting reactant and theoretical yield of a product for a chemical reaction given the moles or mass of two or more reactants.

8.5

44. Define actual yield and determine the percent yield for a chemical reaction given the moles or mass of one or more reactants. 8.6

unit 8: Solution Concentrations (Pages 115–125)

45. Define the following terms and identify specific examples of each in every day solutions: solution, solvent, solute, aqueous solution, concentration, dilute solution, and concentrated solution.

13.2–13.3

46. Define % (m/m) and molarity as conversion factors. Calculate the % (m/m) or molarity of a solution or use % (m/m) or molarity as a conversion factor to solve problems involving mass, moles, and volumes.

13.6

47. Distinguish concentration from density and identify when to use each concept, molarity or density.

48. Solve solution stoichiometry problems to determine the volume or molarity of one reactant needed to react with a given volume or molarity of a second reactant. Given V and M of solution A → moles A → moles B → M or V of solution B. Apply the concept of limiting reactant to reactions in aqueous solutions when necessary.

13.8

49. Define equivalence point and explain why it is useful. Define indicator and end point and explain how these are used in titrations.

Titration lab; 14.6

unit 9: Electrons in Atoms (Pages 126–142)

50. Describe the wave and particle nature of electrons and light. Identify the properties of a wave (frequency, wavelength, amplitude), recognize the Greek letters used to represent these properties, and determine which of these are directly proportional and inversely proportional using the relationship c = λν. Discuss photons in relation to energy, frequency and wavelength using the relationship E = hν.

9.2

51. Name the seven major regions of electromagnetic radiation in order from lowest to highest energy (mnemonic: RMIVUXG). State the range of wavelengths of visible light (in nm) and name the seven majors colors of visible light (mnemonic: ROY G BIV).

9.3

52. Describe the Bohr model of the atom and how this model explains atomic emission spectra. Distinguish between continuous and atomic line spectra, and how it relates to quantized energy. Explain how the atomic emission spectrum of an element can be used to identify an element.

Spectroscopy lab; Workbook;

9.4

53. Describe the quantum mechanical model of the atom and why it is an improvement on the Bohr model. Define probability density and use it to describe where an electron exists in an orbital. Define principal level (or principal shell) and sublevel (or subshell).

9.5

54. Write the full and Noble gas electron configuration and orbital box diagram for atoms, cations, and anions formed from main group elements. 9.6

55. Use the electron configuration to determine the valence and core electrons for main group elements and relate this to general shape of periodic table: s-block, p-block, d-block, and f-block elements.

9.7

CHEM& 140 STudENT LEARNINg OBjECTIVES

unit 10: Chemical Bonding (Pages 143–153)

56. Draw Lewis dot structures for an atom, and use electron configurations to review valence electrons. Describe the concept of the octet (or duet) rule to determine the Lewis dot structures for monatomic ions. Show how ionic compounds are formed by a transfer of electrons

10.2–10.3

57. Draw the Lewis dot structures to represent covalent bonds in molecular compounds and polyatomic ions. Show lone pairs (unshared or non-bonded pairs) and single, double, and triple bonds. Distinguish between covalent bonding and ionic bonding

10.4–10.5

58. Describe how valence shell electron pair repulsion (VSEPR) theory is used to determine molecular shape and count the number of “electron groups” around a central atom (one electron group = a single bond or a double or a triple bond or a lone pair on a central atom).

10.7

59. Given a molecule, predict its electron geometry (linear, trigonal planar, and tetrahedral) and the bond angle around the central atom. Determine molecular geometry (linear, bent, trigonal planar, trigonal pyramidal, tetrahedral) for single center molecules.

10.7

SCIENTIFIC METHOd: OBSERVATION, HYPOTHESIS, LAW, OR THEORY?

Scientific Method: Observation, Hypothesis, Law, or Theory?

■ Observation: Qualitative descriptions or quantitative measurements of what happened in an experiment.

■ Hypothesis: A tentative statement such as “if A happens then B must happen” that can be tested by direct experiment or observation and is falsifiable (can be proven false). A hypothesis supported by evidence can be expressed as a law or a theory. A disproven hypothesis can sometimes be re-tested and found correct as measurements improve.

■ Law: A logical relationship (generalized rule) between two or more things that is based on a series of observations and supported hypotheses. It is often a mathematical state- ment of how two or more quantities relate to each other.

■ Theory: A well-substantiated and reliable explanation for some aspect of nature. Explains why (the underlying reasons) certain laws exist and predicts natural behav- iors beyond the observations and or laws that it is based on. May be modified based on further experiments. Theories are not said to be true, instead they are said to be sup- ported by evidence.

■ Belief: A statement that is not scientifically provable in the same way as facts, laws, hypotheses or theories. Scientifically disproven beliefs can still be held (by a person) to be true.

1. For every action, there is an equal and opposite reaction.

2. F = ma

3. Water freezes at 32 °F.

4. The Earth is a sphere.

5. Matter is composed of small particles called atoms that obey the law of conservation of mass.

6. Humans were created separately from all other life on Earth.

7. Humans and gorillas evolved from a common ancestor species.

8. Light is an electromagnetic phenomenon of energy described by Maxwell’s Laws.

9. The sun will likely die in 7.5 billion years.

10. Sunspots are colder than the surface of the Sun.

11. There are such things as ghosts.

12. Matter can be converted into energy, and they are related by Einstein’s equation, E = mc 2.

13. The positions of the planets can cause humans to act in specific ways.

14. Momentum is the product of a body’s mass and its velocity.

15. We will never know how life started on Earth.

16. The Milky Way is a spiral-type galaxy.

17. The sun will rise tomorrow morning.

18. The Earth is older than 10,000 years.

19. Some numbers are luckier than others.

20. Planetary orbits are ellipses and are governed by the inverse-square law for gravity and Newton’s laws of motion.

Adapted from http://image.gsfc.nasa.gov/poetry

ELEMENT SYMBOLS

Element Symbols

Memorize the names and element symbols for elements #1–56, 78–84. You do not need to memorize the element numbers or order on the periodic table, just the names and ele- ment symbols!

1. Write the symbol for the following elements. Remember that the first letter is capital- ized and the second letter is not capitalized.

a. oxygen _____

b. hydrogen _____

c. chlorine _____

d. mercury _____

e. fluorine _____

f. barium _____

g. helium _____

h. uranium _____

i. radon _____

j. sulfur _____

k. plutonium _____

l. americium _____

m. radium _____

n. germanium _____

o. zinc _____

p. arsenic _____

q. lead _____

r. iron _____

s. calcium _____

t. cobalt _____

2. Write the name of the element that corresponds to each of the following symbols.

a. Kr _______________________

b. K _______________________

c. C _______________________

d. Ne _______________________

e. Si _______________________

f. Zr _______________________

g. Sn _______________________

h. Pt _______________________

i. Na _______________________

j. Al _______________________

k. Cu _______________________

l. Ag _______________________

m. P _______________________

n. Mn _______________________

o. I _______________________

p. Au _______________________

q. Mg _______________________

r. Ni _______________________

s. Br _______________________

t. Hg _______________________

Elements with unique Symbols

The symbols for some elements do not match the English name for the element. This is because the symbols for some elements are based on their ancient names, coming from languages other than English, like Latin, Greek, or even German.

Write the element symbol for each of these elements.

a. lead _________

b. mercury _____

c. tin ___________

d. copper ______

e. potassium ___

f. tungsten _____

g. gold _________

h. silver ________

i. iron __________

j. sodium ______

CHEMISTRY MATTERS!

Chemistry Matters!

Welcome to chemistry! What is chemistry? Chemistry is the study of matter and how it changes. Matter is anything that has mass and takes up space. All matter is composed of very small particles called atoms.

PART 1: ELEMENT, COMPOuNd, MOLECuLE? An atom is the smallest particle that has the characteristics of the element. We will get into more details about atoms (and what they’re made of) in a future chem lesson.

An element is made up of only one type of atom. There are approximately 118 elements known at this time. These are combined in various ways to make up ALL the matter in the universe. How can you tell if something is an element? All of the known elements are listed on the periodic table.

Notice that all the elements are usually abbreviated by their one- or two-letter chemi- cal symbol. The first letter is always capitalized. Here is a list of the elements, by name and symbol: http://en.wikipedia.org/wiki/List_of_elements_by_name (of course, it’s also in your textbook!).

The first 36 elements (plus a few others) in the periodic table are fairly common. You should recognize whether something is an element or not. For example, you should learn that Cu is copper and is on the periodic table which tells us it is an element. CuS is a combination of copper and sulfur (more than one element), so it is a compound. You will always be able to use your periodic table so feel free to take your time and look up the element if you’re not sure!

Checkpoint 1: Name the following elements based on the chemical symbol given. You may use your book or the internet:

a. O ___________________________________

b. Ca __________________________________

c. Na __________________________________

d. Ag __________________________________

Checkpoint 2: Give the chemical symbol for the following elements. You may use your book or the internet:

a. neon _______________________________

b. chlorine ____________________________

c. potassium __________________________

d. iron _________________________________

CHEMISTRY MATTERS!

Compounds are formed when atoms of different elements undergo a CHEMICAL change.

A very important idea in chemistry: Compounds are chemically and physically different than the elements that make them up. For example, water is made up of hydrogen, a flam- mable gas, and oxygen, a gas that we breathe. Water, however, is very different. Water is a liquid at room temperature, it is denser than air, not flammable, and we cannot breathe it.

Atoms can bond with each other and form molecules. Molecules can be considered ele- ments or compounds. See if you can figure out when a molecule is an element and when it is a compound:

Checkpoint 3: A molecule of oxygen gas is composed of two oxygen atoms bonded together. It is considered a(n) _______________ (element or compound)?

Checkpoint 4: A molecule of carbon dioxide gas is composed of one carbon atom and two oxygen atoms bonded together. It is considered a(n) _______________ (element or compound)?

Checkpoint 5: What is the chemical difference between Co and CO?

PART 2: CHEMICAL FORMuLAS Chemists use special notation called “chemical formulas” to describe elements and com- pounds. A molecule of oxygen described in Question 3 above is written as O2. The sub- script 2 indicates the number of O atoms bonded together.

NOTE: Are you wondering why there are two ways oxygen atoms can make molecules— as O2 (oxygen gas) or O3 (ozone). Which is considered an element? They both are! Oxygen gas and ozone are called “allotropes” of oxygen (different forms).

Try answering these questions:

Checkpoint 6: What is the chemical formula of carbon dioxide? _______________

Checkpoint 7: How many atoms in an ozone molecule, O3? _______________

Checkpoint 8: How many atoms are in CuSO4? _______________

Checkpoint 9: How many atoms are in Ba(OH)2? HINT: The ( )2 means that everything inside the ( ) is multiplied by two. _______________

CHEMISTRY MATTERS!

PART 3: MOLECuLAR-LEVEL dIAgRAMS Are you a visual learner? Sometimes it helps to draw pictures of what the atoms and mol- ecules might look like. These are called molecular-level diagrams. In these diagrams, each circle is one atom. The shading of the atom (empty, solid, etc. …) indicates a particular ele- ment. Bonded atoms are shown as touching each other. Here are some examples of molec- ular-level diagrams for terms you just learned above:

A B C

Checkpoint 10: Which boxes show molecules? List all that apply.

Checkpoint 11: Which boxes show elements? List all that apply.

PART 4: MIXTuRES Mixtures contain two or more separate components. Each component retains its properties.

EXAMPLES

Salt water. This mixture is made up of two compounds, water and salt. The water and salt do not undergo a chemical change. How do you know? Because you can separate the water and salt by evaporating off the water.

Sand is a mixture. It contains the components shell, rock, glass, etc.

Checkpoint 12: Which one of these is considered a mixture, A or B? Explain.

A B

CHEMISTRY MATTERS!

There are two types of mixtures, homogeneous and heterogeneous.1

Homogeneous mixtures: The components of the mixture are uniformly distributed through- out the mixture. These are sometimes called solutions. Salt water is an example of a homo- geneous solution. Air is also an example of a homogeneous mixture of various gases (oxy- gen, nitrogen, etc.).

Heterogeneous mixtures: A difference can be seen in the different parts of the mixture. Some examples of heterogeneous mixtures are sand, a chocolate chip cookie, and carbon- ated soda (you can see the difference between the liquid and the bubbles).

How do you tell the difference between a compound, element, and a mixture? Mixtures can be separated into their components by exploiting differences in physical properties. For example, you can take advantage of a difference in phase by using filtration. You prob- ably do this every morning when you make coffee. The coffee filter separates the solid grounds from the liquid coffee.

Checkpoint 13: Which of the following is a pure element (no mixtures)? _____

Checkpoint 14: Which of the following is a pure compound (no mixtures)? _____

Checkpoint 15: Which of the following are mixtures? _____

Checkpoint 16: Which of the following is a homogeneous mixture? _____

Checkpoint 17: Which of the following is a heterogeneous mixture? _____

A B DC

IMPORTANT!

Notice in Checkpoint 12 that the elements hydrogen and oxygen are drawn in Box B as diatomic mol- ecules. There are seven elements that occur as diatomic molecules under standard condition (you should remember this):

Br2, I2, N2, Cl2, H2, O2, and F2 (Mnemonic: Brinclhof or Hofbrincl)

Another note: H atom is an element; H2 molecule is also called hydrogen and is an element, not a com- pound. (Sometimes it’s called “dihydrogen.”)

1 SPECIAL NOTE: It’s also possible to have a colloid—a homogenous mixture in which larger particles (> 1 nm) are dispersed in different phases. One example of a colloid is fog, in which water droplets are suspended in air. Colloids display the Tyndall Effect: http://silver-lightning.com/tyndall/

CHEMISTRY MATTERS!

PART 5: PHYSICAL VS. CHEMICAL Physical properties describe characteristics of matter. Examples are color odor, melting and boiling points, density, and phase (solid, liquid, gas). In a physical change the iden- tity of the chemical remains the same. For example, when ice melts, the chemical is still water. Just the phase has changed. NO NEW SUBSTANCES ARE FORMED IN A PHYSICAL CHANGE.

Chemical properties describe the ability of matter to react or BECOME ANOTHER SUBSTANCE. An example is flammability (the ability to react with oxygen and burn). Most chemical changes involve color changes (as one substance with one color changes into a substance with another color), bubbling as a gas is formed, cloudiness as a precipitate is formed, or the exchange of heat (the release of heat is called an exothermic reaction, the absorption of heat is called an endothermic reaction).

Checkpoint 18: Open YouTube and search for the two videos below. For each one, deter- mine if you are observing a physical change or a chemical change.

a. In YouTube, search for “sublimation of iodine” video (length is 0:52). Or here’s the URL: http://www.youtube.com/watch?v=0_LWBgeQrvk

b. In YouTube, search for “ammonium dichromate volcano” video (length is 0:59). Or here’s the URL: http://www.youtube.com/watch?v=Ula2NWi3Q34

Checkpoint 19: Are each of the following statements chemical or physical descriptions?

a. alcohol evaporates easily _______________

b. iron rusts over time _______________

c. natural gas is burned for heat _______________

d. salt is crushed into a fine powder _______________

Checkpoint 20: Lightning converts O2 (oxygen we breathe) to O3 (ozone). Is this a physical or chemical change? _______________

CHEMISTRY MATTERS!

EXERCISES 1. Define each of these terms and give an example from everyday life, using either words

or chemical symbols/formulas:

a. element

b. compound

c. pure substance

d. mixture

2. Air is composed of roughly 20% oxygen and 80% nitrogen. Make a molecular-level drawing to represent this statement using 10 atoms or molecules to show the relative proportions. Don’t forget the diatomic elements (BrINClHOF)!

3. Is one water molecule considered a pure substance or a mixture? Explain.

4. Can a glass of water be a pure substance, a mixture, or both? Explain.

5. If you have a container with hydrogen gas and oxygen gas in it do you have water? Explain why or why not.

6. Categorize each of the following as a pure substance, a homogeneous mixture, or het- erogeneous mixture.

a. white grape juice ________________________________________________________________

b. open bottle of beer ______________________________________________________________

c. soda with ice cubes _____________________________________________________________

d. fresh garden salad _______________________________________________________________

CHEMISTRY MATTERS!

7. Suppose salt water (consisting of NaCl and water) is found to contain several grains of pure, quartz sand (sand is composed of silicon dioxide, SiO2). The mixture is filtered and the solid sand placed in a vial labeled A. Half the remaining liquid is placed in vial B. The other half of the liquid is heated until it boils. The gas that escapes is collected and cooled. The liquid that reforms is placed in vial C. The solid that remains after the boiling is placed in vial d.

Use the following four choices to categorize each sample below:

a. pure element

b. pure compound

c. homogeneous mixture

d. heterogeneous mixture

i. The original sample _____

ii. The sample in vial A _____

iii. The sample in vial B _____

iv. The sample in vial C _____

v. The sample in vial d _____

8. In each of the three boxes below, draw the following. Use different shading to show different types of atoms.

a. Pure substance of molecules (that are compounds) in the liquid phase

b. Homogeneous mixture of atoms in the solid phase

c. Heterogeneous mixture of molecules in the gas phase

MOLECuLAR LEVEL dIAgRAMS

Molecular Level diagrams

Refer to these drawings for the questions that follow.

A B D EC

1. Which of these drawings contains elements only (no compounds)? List all. _________

2. Which of these drawings represent pure substances? List all. ___________________

3. Which of these drawings contains molecules? List all. _________________________

4. Which drawings contain at least two chemical substances? List both. ____________

5. Drawing C shows a (circle one):

pure substance / homogeneous mixture / heterogeneous mixture

6. Which drawing includes two different diatomic molecular elements? One answer: ___

7. Which drawings contain substances in their gaseous state? List all. ______________

8. Which drawing has only a single monatomic element? _______________

9. Which drawings include a compound? Circle all that apply:

A / B / C / D / E

10. Which drawing(s) show a mixture of two substances? _________________________

11. Drawing E shows (circle all that apply):

two diatomic elements / two molecular elements / a compound

12. Which drawing appears to contain substances in two different physical states? _______________

13. The outer layer of a penny is composed of copper. Which drawing best represents the penny’s outer layer? _______________

14. Which drawing best represents a sample of liquid water? _______________

15. Air is composed of gaseous O2 and N2 molecules. Which diagram best represents air? _______________

ELEMENTS, COMPOuNdS & MIXTuRES

Elements, Compounds & Mixtures

PART 1 Read the following information on elements, compounds and mixtures. Fill in the blanks where provided.

ELEMENTS

■ A pure substance contains only one _______________. ■ An element is always uniform all the way through (homogeneous). ■ An element _______________ be separated into simpler materials (except during

nuclear reactions).

■ Over 100 existing elements are listed and classified on the _______________.

COMPOuNdS

■ A compound contains two or more _______________. ■ The atoms are _______________ combined in some way. Often times (but not always)

they come together to form groups of atoms called molecules.

■ A compound is always homogeneous (uniform). ■ Compounds _______________ be separated by physical means. Separating a com-

pound requires a chemical reaction.

■ The properties of a compound are usually _______________ the properties of the ele- ments it contains.

MIXTuRES

■ In a mixture two or more _______________ or _______________ are NOT chemically combined.

■ There is no reaction between substances. ■ Mixtures can be uniform (called _______________ mixtures) and are known as

solutions.

■ Mixtures can also be non-uniform (called _______________ mixtures). ■ Mixtures _______________ be separated into their components by chemical or physi-

cal means.

■ The properties of a mixture are _______________ the properties of its components.

ELEMENTS, COMPOuNdS & MIXTuRES

PART 2 Classify each of the following as elements (E), compounds (C) or Mixtures (M). Write the letter X if it is none of these.

_____ Diamond (C)

_____ Air

_____ Krypton (K)

_____ Water (H2O)

_____ Ammonia (NH3)

_____ Wood

_____ Dry Ice (CO2)

_____ Sugar (C6H12O6)

_____ Sulfuric Acid (H2SO4)

_____ Bismuth (Bi)

_____ Alcohol (CH3OH)

_____ Salt (NaCl)

_____ Bronze

_____ Baking Soda (NaHCO3)

_____ Milk

_____ Gasoline

_____ Uranium (U)

_____ Pail of Garbage

_____ Energy

_____ Ink

_____ Titanium (Ti)

_____ Iron (Fe)

_____ Electricity

_____ Popcorn

_____ A dog

_____ Gold (Au)

_____ Pizza

_____ Concrete

PART 3 Match each diagram with its correct description. Diagrams will be used once.

A B D EC

1. Pure element—only one type of atom present _____

2. Mixture of two elements—two types of uncombined atoms present _____

3. Pure compound—only one type of compound present _____

4. Mixture of two compounds—two types of compounds present _____

5. Mixture of a compound and an element _____

ELEMENTS, COMPOuNdS & MIXTuRES

PART 4 Identify whether the substance is an element (E), a compound (C), a heterogeneous mix- ture (HM), or a homogenous mixture (hm).

1. Chocolate chip cookie _______

2. Steam _______

3. Toothpaste _______

4. Chicken soup _______

5. Tap water _______

6. Salt water _______

7. Pencil lead (graphite) _______

8. A burrito _______

9. Lemonade _______

10. Distilled water _______

11. Dirt _______

12. Soda/cola _______

13. Italian salad dressing _______

14. Pure gold _______

15. Smoothie _______

SCIENTIFIC NOTATION: POWERS OF 10

Scientific Notation: Powers of 10

In the sciences, sizes of atoms and molecules are extremely small compared to the sizes of cells and small organisms, and even smaller than the distance to planets and galaxies. Many zeros in a number can lead to error, so a compact way of indicating size is to use exponents/powers of 10. guideline: Use scientific notation generally when numbers are larger than 103 (1000) or smaller than 10–3 (0.001).

Here are some examples of how scientific notation is used:

decimal Notation Scientific Notation

0.000134 1.34 × 10–4 (coefficient is 1.34)

895000 8.95 × 105 (coefficient is 8.95)

2100.7 2.1007 × 103 (coefficient is 2.1007)

1. For numbers smaller than 1, do you move the decimal to the left or the right to convert the number into scientific notation?

2. For numbers larger than 1, do you move the decimal to the left or the right to convert the number into scientific notation?

3. True or False? The coefficient in scientific notation is usually between 1 and 10.

4. Express the two numbers below in decimal form. Verify that they are the same number.

a. 450 × 10–9 __________________________ b. 4.5 × 10–7 ___________________________

5. Practice converting these numbers into scientific notation:

a. 0.33 ________________________________

b. 45,000 ______________________________

c. 3,450,000 ___________________________

d. 0.009708 ____________________________

6. Practice converting these numbers into decimal form:

a. 4 × 102 ______________________________

b. 5.3 × 10–3 ___________________________

c. 1.01 × 10–5 __________________________

d. 8.9345 × 106 ________________________

7. Practice converting these to “standard” scientific notation:

a. 39 × 10–1 ____________________________

b. 3859 × 107 __________________________

c. 0.00075 × 108 _______________________

d. 53.498 × 10–34 ______________________

EXPONENTS

Exponents

How to perform simple mathematical operations on exponents without a calculator:

10a × 10b = 10(a+b)

10a ÷ 10b = 10(a–b)

(10a)b = 10(a*b)

100 = 1

Practice doing these calculations without a calculator. Put your answer in scientific nota- tion when appropriate (smaller than 10–4 or larger than 104):

a. (1 × 109) × (1 × 10–4) = _______________

b. (1 × 104)/(1 × 10–6) = _________________

c. (3.50 × 107)2 = _______________________

d. 0.45 × 100 = _________________________

uSINg A SCIENTIFIC CALCuLATOR

Try this calculation on your calculator: 60 ÷ 2.5 × 104 = _____________

The answer should be 2.4 × 10–3. If you are not careful, your calculator will tell you it’s 240,000! This type of error tends to occur when you divide by powers of 10.

There are several ways to avoid the wrong answer:

■ You can use parentheses: 60 ÷ (2.5 × 10^4) ■ You can use the EE or EXP function on your calculator as a shortcut for “10^.” Try 60 ÷

2.5E4. (Ask your instructor if you have trouble locating these buttons on your calculator!)

Practice on these calculations:

1. (3.5 × 104) ÷ (2 × 109) = _____________ (Answer: 1.75 × 10–5 or 1.75E-5)

2. (4.3 × 102) × (5 × 10–3) = _____________ (Answer: 2.15)

3. (3.2 × 104) – (2.0 × 103) = _____________ (Answer: 3.0 × 104)

4. (5.9 × 102) + (1.6 × 103) = _____________ (Answer: 2.2 × 103)

5. (6.4 × 102) ÷ (3.0 × 10–4) = _____________ (Answer: 2.1 × 106)

SIgNIFICANT FIguRES

Significant Figures

WHAT ARE SIgNIFICANT FIguRES? In a recorded measurement, the significant figures (sig. figs. or s.f.) include all of the certain digits, plus one last uncertain digit (estimated or guessed digit).

# of sig figs = # certain digits + 1 last uncertain digit

WHY dO WE uSE SIgNIFICANT FIguRES? Significant figures are used as a way of communicating the amount of precision that a recorded measurement has. When you record a measurement to a specific decimal place you are communicating the precision of the measuring device that was used. If you write too few or too many digits you are misrepresenting the precision of the measurement, sug- gesting that the measurement was more approximately made or more finely made than it actually was. Below are examples of volume measurements in which the units are millili- ters (mL).

EXAMPLE

3.0 mL The 3 is certain and the 0 is uncertain. There is a total of 2 significant figures (2 s.f.).

Try these:

1) 3.1 mL _____ 2) 3.05 mL _____ 3) 3.00 mL _____ 4) 3.000 mL _____

ANOTHER EXAMPLE

0.31 mL The 3 is certain and the 1 is uncertain. There is a total of 2 s.f. The 0 is neither, it is a place holder to indicate magnitude.

Now, try these:

5) 0.30 mL _____ 6) 0.300 mL _____ 7) 0.00300 mL _____

A METHOd FOR COuNTINg SIgNIFICANT FIguRES Reading left to right, start counting at the first non-zero digit.

■ If the number has a decimal point, count all the way to the end. ■ If there is no decimal point included, stop counting at the last non-zero digit.

Try a few more:

8) 3 mL _____

9) 31 mL _____

10) 301 mL

11) 310 mL _____

12) 30000. mL _____

13) 30000 mL _____

14) 30000.00 mL ____

MEASuREMENTS ANd SIgNIFICANT FIguRES 1

Measurements and Significant Figures 1

1. Accuracy refers to how close a measurement is to an accepted or true value. Suppose you are a technician measuring a 1 gram weight on four different balances and obtain the following four measurements. Which measurement is from the most accurate balance?

a. 0.99 g b. 1.20 g c. 1.05 g d. 0.89 g

2. Precision refers to the fineness of a measurement and is often indicated by how many digits are given in a measurement. Suppose you are a nurse using three different ther- mometers. Which measurement is from the highest precision instrument?

a. 99 °F b. 98.6 °F c. 98.584 °F

3. Measure the arrow using the ruler provided below it using the number of digits allowed by this ruler.

1 2 3 4 5 6 7 8 9 10 cm

4. Provide a measurement for the temperature displayed on this Celsius thermometer (a portion of the thermometer is shown, the dashed line indicates the temperature).

5. Beaker

a. Report a measurement for the volume in this beaker.

b. Would this beaker be a good choice for measuring a vol- ume of 5.0 mL?

21 °C

20 °C

100 200 300 400 500

MEASuREMENTS ANd SIgNIFICANT FIguRES 2

Measurements and Significant Figures 2

In the lab you will practice using various instruments to measure length, volume, or mass. Here are some common measuring tools. Determine the appropriate number of decimal places to record for each measuring tool based on the level of graduation.

Decimal places:

hundreds tens ones decimal point tenths hundredths thousandths

100 10 1 . 1/10 1/100 1/100

1. Ruler

a. In centimeters, record to __________

b. In millimeters, record to __________

2. 100-mL graduated cylinder

a. Record to __________

3. Buret

a. Record to __________

b. This reading is about __________ mL __________

(NOTE: A buret is marked in reverse!)

4. 10-mL graduated cylinder

a. Record to __________

b. This reading is about __________ mL

5. 100-mL beaker

a. Record to __________

SIgNIFICANT FIguRES ANd SCIENTIFIC NOTATION

Significant Figures and Scientific Notation

1. How many significant figures are in:

a. 1.204 × 10–2 g ______

b. 3.160 × 108 Å ______

c. 0.00281 g ______

d. 810 mL ______

e. 12.82 Liters ______

f. 3.19 × 1015 atoms ______

g. 4.300 × 10–6 cm ______

h. 0.00641 g ______

i. 8.2354 × 10–19 m ______

j. 0.0559 g ______

k. 2.92 × 102 g ______

l. 4.1 Liters ______

m. 0.0002 cm ______

n. 0.045 g ______

2. Convert the standard (decimal) notation numbers above into standard scientific nota- tion. Remember to carry all significant figures into the coefficient.

3. Assign ± error readings to the following measurements (HINT: Which digit has uncer- tainty?) Assume units of 1 (0.1, 0.01, etc.)

a. 3.412 g

b. 45 mL

c. 0.00498 g

d. 8.2 cm

e. 559 L

f. 1.00 × 102 m

SIgNIFICANT FIguRES IN CALCuLATIONS

Significant Figures in Calculations

ROuNdINg If the digit to the right of the final digit you want to keep is

■ less than 5, round down. ■ 5 or more, round up.

EXAMPLES

Round 0.067556 to 2 sig figs. Answer: 0.068

Round 3.145 × 103 to 3 sig figs. Answer: 3.15 × 103

MuLTIPLICATION/dIVISION Round your calculated answer to have the same number of significant figures as the least precise measurement (fewest number of s.f.).

EXAMPLES

5.02 × 89.665 × 0.10 = 45.00681 = 45 (3 s.f. × 5 s.f. × 2 s.f. → round to 2 s.f.) 5.892/6.10 =

AddITION/SuBTRACTION Round your calculated answer to have the same number of decimal places as the least pre- cise measurement.

EXAMPLES

5.74 + 0.823 + 2.651 = 9.214 = 9.21 (rounded to 1/100s place, results in 3 s.f.)

0.987 + 125.1 – 1.22 =

SIgNIFICANT FIguRES IN CALCuLATIONS

MuLTI-STEP/COMBINATION CALCuLATIONS Keep track at each step based on multiply/divide or add/subtract, but only round at the very end!

EXAMPLE

6.78 × 5.903 × (5.489 – 5.01) = 19.17070086 = 19

(1/100 place for subtraction step results in 2 s.f, then in multiply step 2 s.f. is fewest → round to 2 s.f.)

19.667 – (5.4 × 0.916) =

PRACTICE

1. Use the equation for converting 545.0 kelvins to °C (degree Celsius) and round using significant figures. °C = K + 273.15 =

2. Use the equation for °C/°F to convert body tem- perature, 37.0 °C to units of °F.

(°C * 1.8) + 32 = °F [NOTE: The 1.8 and 32 are considered exact numbers.]

3. Convert 114°F to °C. (Answer: 45.5°C)

4. Convert –10.0°C to °F and Kelvin. (Answer: 14°F, 263 K)

212 °F

32 °F

Fahrenheit

100 °C

0 °C

Celsius

373 K

273 K

Kelvin

LIST OF METRIC PREFIXES

List of Metric Prefixes

Prefix Symbol Mathematical Meaning In Words …

Yotta- Y 1024

Zetta- Z 1021

Exa- E 1018

Peta- P 1015

Tera- T 1012 trillion

giga- g 109 billion

Mega- M 106 million

Kilo- k 103 thousand

Hecto- h 102 hundred

Deca- da 101 ten

1 = 100 one

deci- d 10–1 tenth

Centi- c 10–2 hundredth

Milli- m 10–3 thousandth

Micro- μ 10–6 millionth

Nano- n 10–9 billionth

Pico- p 10–12 trillionth

Femto- f 10–15

Atto- a 10–18

Zepto- z 10–21

Yocto- y 10–24

Memorize the prefixes, symbols, and mathematical meanings in bold (tera, giga, mega, kilo, deci, centi, milli, micro, nano, pico).

COMMONLY uSEd BASE uNITS IN THE METRIC SYSTEM

dimension unit

Mass grams (g)

Length meters (m)

Volume liters (L)

Time seconds (s)

METRIC SYSTEM uNITS & dEFINITIONS

Metric System units & definitions

1. Fill in the following blanks for units of length.

a. 1 terameter (Tm) = 1012 m or 1 m = 10–12 Tm

b. 1 gigameter (Gm) = __________________ m or 1 m = _______________ Gm

c. 1 megameter (Mm) = _________________ m or 1 m = _______________ Mm

d. 1 kilometer (km) = ___________________ m or 1 m = ________________ km

e. 1 decameter (dam) = _________________ m or 1 m = _______________ dam

f. 1 centimeter (cm) = __________________ m or 1 m = ________________ cm

g. 1 decimeter (dm) = __________________ m or 1 m = ________________ dm

h. 1 millimeter (mm) = __________________ m or 1 m = _______________ mm

i. 1 micrometer (µm) = _________________ m or 1 m = ________________ µm

j. 1 nanometer (nm) = __________________ m or 1 m = ________________ nm

k. 1 picometer (pm) = __________________ m or 1 m= ________________ pm

2. Give the full names of each of the following units:

a. mL __________________________________

b. Ds __________________________________

c. mg _________________________________

d. cL __________________________________

3. Write symbols for each of the following units:

a. Liter _________________

b. decimeter ____________

c. kilosecond ___________

d. megameter __________

e. microliter ____________

f. milligram ____________

4. Express each of the following in terms of their base units (e.g., 1 mL = 0.001 L)

a. 1 cm _________ b. 1 mg _________ c. 1 km _________ d. 1 µL __________

5. Write the correct prefixes for the following measurements:

a. 1000 g = ___________ g

b. 0.001 m = __________ g

c. 1,000,000 s = _______ g

d. 0.01 g = ____________ g

e. 0.0046 g = _________ g

f. 3300 s = ___________ g

6. Which is the larger unit?

a. centimeter or decimeter

b. milliliter or microliter

c. milligram or megagram

d. deciliter or decaliter

METRIC SYSTEM uNITS & dEFINITIONS

7. Complete the following blanks:

a. 1 L = ____________________________ mL

b. 1 km = ___________________________ m

c. 1 mg = ____________________________ g

d. 1 µL = _____________________________ L

e. 1 cm = ___________________________ m

f. 1 g = _____________________________ µg

g. 1 dL = _____________________________ L

h. 1 µm = ___________________________ m

i. 1 Tg = _____________________________ g

j. 1 m = ___________________________ cm

k. 1 mL = ____________________________ L

l. 1 L = _____________________________ µL

m. 1 mm = __________________________ m

n. 1 ML = ____________________________ L

o. 1 L = _____________________________ µL

p. 1 m = __________________________ dam

q. 1 mL = ____________________________ L

r. 1 L = ____________________________ GL

TWO METHOdS FOR METRIC CONVERSIONS

Two Methods for Metric Conversions

There are several ways to perform metric conversions. Choose the method that works best for you. For example, let’s convert 0.000500 meters to micrometers.

Write the relationship between the base unit and the prefixed unit.

The micro (µ) prefix mathematically means 10–6. The equality between the base unit meter and micrometer is 1 µm = 1 × 10–6 m, since µ means 10–6 (you are replacing the µ symbol with 10–6). Or, you can write this as 1 m = 106 µm if you think in terms of 1 of the larger unit equals a lot of the smaller unit.

1 µm = 1 × 10–6 m or 1 m = 106 µm

METHOd #1: uSINg dIMENSIONAL ANALYSIS

POSSIBLE CONVERSION FACTORS

A conversion factor is a relationship between two units written as a fraction. The relation- ship between micrometers and meters can be written two different ways:

1 μm or

10–6 m 10–6 m 1 μm

SETTINg uP THE dIMENSIONAL ANALYSIS CALCuLATION

Start with what you are trying to convert, the given information, and write that down. Choose which version of the conversion factor to use so that the units you are trying to convert cancel out.

0.0005000 m × 1 μm

= 0.0005000 × 106 = 5.00 × 102 μm 10–6 m

METHOd #2: MOVINg THE dECIMAL POINT For this method, it is helpful to use a drawing of the metric prefix scale.

106 k Base

Unit103 d

10–1 c

10–2 m

10–3 μ

10–6

Find the base unit (marked as –) then count spaces over to the µ symbol. There are six spaces, so there is a difference of 106 in scale. Since micrometers are smaller than meters, converting from meters to micrometers results in a more of those smaller units, so the dec- imal point moves to the right six times. (0.000500 becomes 500.)

0.000500 m = 500 µm = 5.00 × 102 µm

METRIC CONVERSIONS: dIMENSIONAL ANALYSIS

Metric Conversions: dimensional Analysis

1. Fill in the number which represents the meaning of each prefix. Use scientific notation (or exponents) when appropriate, usually for numbers smaller than 0.001 and larger than 1000.

a. m _______

b. k _______

c. M _______

d. c _______

e. μ _______

f. n _______

g. d _______

h. da _______

i. G _______

2. Complete each conversion as a fraction with the correct number.

a. 1 cm m

b. 1 m

c. 1 kg g

d. g 1 kg

e. 1 mL L

f. 1 L

g. 1 ns s

h. ns

3. Can you flip each conversion? Write these next to the conversions above. Check your answers with others.

4. Using any of the conversions above, fill in the blanks below and solve:

a. 30 cm × ⎛ ⎜ ⎝

⎞ ⎟ ⎠

= _________ m

b. 45 m × ⎛ ⎜ ⎝

⎞ ⎟ ⎠

= _________ cm

c. 150 g × ⎛ ⎜ ⎝

⎞ ⎟ ⎠

= kg

d. 10 mL × ⎛ ⎜ ⎝

⎞ ⎟ ⎠

= _________ L

e. 0.005 ks × ⎛ ⎜ ⎝

⎞ ⎟ ⎠

× ⎛ ⎜ ⎝

⎞ ⎟ ⎠

= _________ ns

5. Now try these conversions from scratch. Use the set up provided above and cancel out your units.

a. 900 g to kg

b. 3.4 cm to m

c. 800 mm to m

d. 0.0450 kg to cg

e. 31 mm to nm

f. 2.87 kL to μL

ONE-STEP METRIC CONVERSIONS uSINg dIMENSIONAL ANALYSIS

One-Step Metric Conversions using dimensional Analysis

Perform the following metric system conversions that involve converting between a base unit and a prefixed unit. Put all answers in both standard and scientific notation. Use sig- nificant figures and use dimensional analysis to show your work.

1. 2.14 kg to grams

2. 6172 mm to meters

3. 0.0256 L to milliliters

4. 3.14 × 10–5 g to micrograms

5. 1.01 × 10–8 s to deciseconds

6. How many milliliters are in 5 liters?

7. How many kilograms are in 598 grams?

8. How many centiseconds are in 12 seconds?

TWO-STEP METRIC CONVERSIONS uSINg dIMENSIONAL ANALYSIS

Two-Step Metric Conversions using dimensional Analysis

Perform the following metric system conversions that involve converting between two prefixed units.

Use the relationships between each prefixed unit and the base unit to generate two con- version factors that will allow you to convert in a two-step process from the given prefixed unit to the base unit and then to the desired prefixed unit.

Put all answers in both standard and scientific notation. Use significant figures and use dimensional analysis to show your work.

EXAMPLE

How many centimeters are in 0.5 kilometers?

0.5 km × 103 m

× 1 cm

= 0.5 × 105 cm = 5 × 104 cm (or 50,000 cm) 1 km 10–2 m

1. How many millimeters are in 29 micrometers?

2. How many kilograms are in 100 gigagrams?

3. Convert 1316 mg to kilograms

4. Convert 5.88 mL to deciliters

5. Convert 2.19 cm to micrometers

METRIC SYSTEM: THE LAddER METHOd

Metric System: The Ladder Method

Count the number of “jumps” from start to end. Move the decimal that number of jumps in the correct direction.

EXAMPLES

35 milligrams (mg) = ___________ grams (g)

5 kilometers (km) = ___________ meters (m)

1. Write the correct abbreviation for each metric unit. Identify the base unit in each.

Abbreviation Base Unit

a. kilogram ________ ________

b. milliliter ________ ________

c. decimeter ________ ________

d. meter ________ ________

e. megameter ________ ________

Abbreviation Base Unit

f. centimeter ________ ________

g. gram ________ ________

h. liter ________ ________

i. decagram ________ ________

j. microliter ________ ________

2. Try these conversions using the ladder method. Use scientific notation for numbers bigger than 1000 and smaller than 0.001.

a. 2000 mg = _______________________ g

b. 5 L = ___________________________ mL

c. 16 cm = _______________________ mm

d. 104 km = ________________________ m

e. 2500 s = ________________________ ns

f. 480 cm = _______________________ dm

g. 8 mm = _________________________ cm

h. 65 daL = _______________________ mL

i. 198 g = __________________________ kg

j. 120 mg = _________________________ g

3. Complete these equalities or inequalities with the symbols =, > (greater than), or < (less than).

a. 63 cm _________________________ 6 m

b. 5 g _________________________ 508 mg

c. 536 cm ____________________ 53.4 dm

d. 1500 mL ______________________ 1.5 L

e. 500 mg ________________________ 5 g

f. 3.6 ns ______________________ 360 ps

g. 4 × 10–7 m __________________ 400 nm

h. 3.5 × 10–3 mm ______________ 3.5 μm

i. 100 pL ____________________ 10–6 mL

j. 500 Gb ________________________ 5 Tb

METRIC CONVERSIONS

Metric Conversions

Use significant figures. Practice with dimensional analysis or the ladder method.

Use scientific notation for numbers larger than 104 and smaller than 10–4.

1) 1 mm = __________ m

2) 1 km = __________ m

3) 1 m = __________ cm

4) 1 m = __________ km

5) 1 cm = __________ μm

6) 5.0 mm = __________ m

7) 15 km = __________ cm

8) 6.00 × 102 m = __________ cm

9) 0.45 km = __________ m

10) 104 km = __________ μm

11) 1 × 109 Mm = __________ cm

12) 6 × 10–4 cm = __________ nm

13) 1 μm = __________ km

14) 4.9 × 10–10 mm = __________ km

15) 8 × 10–100 mm = __________ nm

Answers: (1) 0.001 m (2) 1000 m (3) 100 cm (4) 0.001 km (5) 1 × 104 μm (6) 5.0 × 10–3 m (7) 1.5 × 106cm (8) 6.00 × 104 cm (9) 450 m (10) 1.04 × 1011 μm (11) 1 × 1017 cm (12) 6 × 103 nm (13) 1 × 10–9 km (14) 4.9 × 10–16 km (15) 8 × 10–94 nm

MuLTI-STEP CONVERSIONS

Multi-Step Conversions

1. For the following multi-step conversion problem, use the provided unit relationships to fill in the conversion factors in the unit analysis set up.

A gas can holds 2.0 gallons of gasoline. What is this quantity in teaspoons?

1.057 quarts = 1 liter

4 quarts = 1 gallon

1 teaspoon = 4.93 milliliters

1,000 milliliter = 1 liter (you should know this!)

20 gallons × × × × =

2. A swimming pool is 50 meters long. How many feet is this? (1 in = 2.54 cm, exactly) (12 in = 1 ft) (Answer: 200 ft)

3. It has been reported that the distance from Los Angeles to New York City is 177,000,000 inches. What is this distance in miles? (Check unit conversions listed on the inside front cover.) (Answer: 2.79 × 103 m)

4. A unit of length called the “furlong” is used in horse racing. The units of length called the “chain” and the “link” are used in surveying. What is the length of 1 link (in inches) if: There are 8 furlongs in 1 mile (mi), 10 chains in 1 furlong, and 100 links in 1 chain. Report your answer to three significant figures. (Answer: 7.92 in)

MuLTI-STEP CONVERSIONS

5. Use these relationships for the problems that follow.

1 nautical mile = 6076.11549 feet

1 inch = 2.54 cm

3 feet = 1 yard

1 league = 5280 yards

1 cable = 120 fathoms

1 fathom = 6 feet

1 degree = 69.047 miles

1 mile = 5280 feet

1 hand = 4 inches

1 foot = 12 inches

1.609 km = 1 mile

a. Your cruise ship is leaving for a 610-league adventure. How many nautical miles is this? (Answer: 1600 na mi)

b. Later the ship is discovered at 38 fathoms deep under water. Convert this to meters. (Answer: 69 m)

c. Fortunately you survived! You are stranded on a deserted island that is located 12.5 degrees north of the equator. How many kilometers is this? (Answer: 1.39 × 103 km)

d. To reach the top of a palm tree for a coconut you will have to climb 7.4 meters. How many hands is this? (Answer: 73 hands)

VOLuME CONVERSIONS

Volume Conversions

LIQuId-BASEd VOLuME uNITS RELATEd TO CuBEd LENgTH VOLuME uNITS How many cubic millimeters (cm3) are in a 2-L bottle of soda? At first glance, it may seem like the units are not related because the unit “cm3” does not contain “L” and instead con- tains “cm” which is a length, not a volume. But, the unit of cm3 is indeed a volume and does relate to a volume in liters.

The relationship between a liter-based volume and a cubed-length-based volume is a defined equality.

1 cm3 = 1 mL or 1000 cm3 = 1 L or 1 m3 = 106 mL

The most commonly memorized of these is

1 cm3 = 1 mL

PRACTICE

1. How many milliliters are in 853 cm3?

2. How many cubic centimeters are in 3597 mL?

3. Convert 0.020 L to cm3.

4. How many cubic centimeters are in 8500 gallons? (Answer: 3.2 × 107 cm3)

5. How many cubic meters are in 8500 gallons? (Answer: 32 m3)

RATIOS AS CONVERSION FACTORS

Ratios as Conversion Factors

When you see a ratio given with different units, it can be written as a conversion factor.

For example, the density of ethanol is 0.789 grams per milliliter (same as 0.789 g/mL). If you are using it to convert grams to mL (or vice versa), then write it as a fraction:

0.789 g or

1 mL 1 mL 0.789 g

Think about the chain of conversions you’ll need to do, linking the units together: If you have 4.00 kg of ethanol, and you want to find out how many liters of ethanol you have, pro- vided the density is 0.789 grams/mL, what conversions do you need?

Unit Map: kg g mL L

Set it up here:

Use dimensional analysis to solve all of these problems. Use significant figures and unit abbreviations.

1. The density of gasoline is 0.70 kg/L.

a. Write this as a fraction.

b. What is the mass in kilograms of 15.6 liters of gasoline?

2. When one gram of liquid ethanol, C2H5OH, is burned, 29.7 kJ of heat are released.

a. Write this as a fraction.

b. How much heat in kilojoules is released when 4.274 grams of liquid ethanol are burned?

3. A sample of iron has a mass of 242.6 grams. What is the volume in cubic centimeters of this iron? Iron has a density of 7.86 g/cm3.

a. Write out a unit map to solve this problem.

b. Solve the problem.

RATIOS AS CONVERSION FACTORS

4. Suppose an adult male has about 11 kg of fat. Each gram of fat can provide the body with about 38 kJ of energy. If this person requires 8.0 × 103 kJ of energy per day to sur- vive, how many days could he survive on his fat alone? (Answer: 52 days)

5. A doctor orders Medrol to be given 1.5 mg/kg of body weight. Medrol is an anti-inflam- matory administered as an intramuscular injection. If a child weighs 72.6 pounds and the available stock of Medrol is 20. mg/mL, how many milliliters does the doctor pre- scribe? (Answer: 2.5 mL)

a. Write the two ratios in this problem as fractions.

b. Solve the problem.

6. A 5.00 mL block of aluminum contains how many aluminum atoms? The density of alu- minum is 2.71 grams per mL, the molar mass of aluminum is 26.98 grams per mole, and one mole of aluminum is 6.022 × 1023 atoms. (Answer: 3.01 × 1023 atoms)

7. Given that there are four quarts in a gallon, 453.59 g in exactly 1 pound, and 946.3 mL in exactly 1 quart, calculate the mass of one gallon of mercury, hence verifying the state- ment that a gallon of mercury weighs more than 100 pounds. (The density of mercury is 13.55 g/mL.) (Answer: 113.1 lbs)

THE jABBERWOCKY ANd gOOgS PER MuLM

The jabberwocky and googs per Mulm

Use dimensional analysis to solve these nonsense conversions.

1. Nonsense words taken from the poem Jabberwocky (from Lewis Carroll’s Through the Looking Glass)

■ There are 20 tumtum trees in the tulgey wood. ■ In each tulgey wood is one frumious Bandersnatch. ■ There are 5 slithy toves in 2 borogoves. ■ There are 2 mome raths per Jabberwock. ■ There are 2 Jubjub birds in 200 tumtum trees. ■ There are 200 mome raths in each borogove. ■ There are 5 Jubjub birds per slithy tove.

If there are 5 frumious Bandersnatches, how many Jabberwocks should there be? (Answer: 8 JW)

HINT: What is your “given” info (other than conversion factors)? Use this to start the set up. See “A Curious Exercise in Backchaining” for help.

2. A spaceship from another planet travels at a speed of 4.27 googs per mulm. There are 256 googs in a plotz and 12.3 plotz in a wraslm. If 3.4 tpocks equal one mulm, what is the ship’s speed in wraslm per tpock? Show your work. In this case, assume numbers with decimal places are measured but whole numbers are exact (we’ll pretend that “googs” are creatures!). (Answer: 4.0 × 10–4 wraslm/tpock)

CONVERTINg BOTH NuMERATOR & dENOMINATOR

Converting Both Numerator & denominator

EXAMPLE

The speed of light is 2.998 × 108 m/s. How many feet per minute is this?

2.998 × 108 m ×

102 cm ×

1 in ×

1 ft ×

60 s = 5.902 × 1010 ft/min

s m 2.54 cm 12 in 1 min

1. Convert 2.7 × 10–8 J/g to cal/oz. (Answer: 1.8 × 10–7 cal/oz)

2. Convert 14 m/s2 to ft/hr2 to 2 significant figures. (Conversion: 1 hr2 = 1.296 × 107 s2) (Answer: 6.0 × 108 ft/hr2)

3. Average gasoline mileage for a particular sports utility vehicle is 13 miles per gallon of gasoline. Express this miles/gal value in km/L. (Answer: 5.5 km/L)

4. The kidneys of a normal adult male filter 125 milliliters of blood per minute. How many gallons of blood are filtered per day? (Answer: 47.6 gal/day)

5. A patient in the hospital is given an intravenous fluid that must deliver 1.0 × 103 cc of a dextrose (sugar) solution over 8.0 hours. The intravenous fluid tubing delivers 15 drops/cc. What is the drop rate (in units of drops/min) that must be administered to the patient? (Answer: 31 drops/min)

SQuAREd OR CuBEd uNIT CONVERSIONS

Squared or Cubed unit Conversions

SQuAREd uNITS Consider a piece of paper that is 8.5 inches by 11 inches. The area of the paper is 8.5 in × 11 in = 93 in2. If you wanted to report this area in units of square centimeters (cm2), you will need to convert the units.

8.5 in × 2.54 cm

= 21.59 cm 11 in × 2.54 cm

= 27.94 cm 1 in 1 in

21.59 cm × 27.94 cm = 603.2246 cm2 = 6.0 × 102 cm2

This is the same as

8.5 in × 11 in × 2.54 cm

× 2.54 cm

= 603.2246 cm2 = 6.0 × 102 cm2 1 in 1 in

And, that is the same as

93 in2 × ⎛ ⎜ ⎝

2.54 cm ⎞ ⎟ ⎠

= 603.2246 cm2 = 6.0 × 102 cm2 1 in

Often you may not be given the measurements for the length or width but only the total area, so converting the squared units of an area by squaring the length-based conversion factor is a skill you need to master.

When starting with a squared unit and trying to convert it to another square unit, you will need to square the conversion factor.

original squared unit × (conversion factor)2 = another squared unit

EXAMPLE

Convert 87 ft2 to yd2. Given: 3 ft = 1 yd.

87 ft2 × ⎛ ⎜ ⎝

1 yd ⎞ ⎟ ⎠

= 9.6667 yd2 = 9.7 yd2 3 ft

SQuAREd OR CuBEd uNIT CONVERSIONS

CuBEd uNITS Unless a volume is a liquid-based unit like liters or gallons, volume units are based on lengths in three dimensions, such as cubic meters (m3) or cubic yards (yd3). The concept for performing cubed unit volume conversions is very similar to the squared unit area con- versions shown above, except that instead of working in two dimensions and squaring the length-based conversion factor, volumes involve three dimensions and cubing the length- based conversion factor.

EXAMPLE

Convert 246 m3 to mm3.

246 m3 × ⎛ ⎜ ⎝

103 mm ⎞ ⎟ ⎠

= 246 × (103)3 mm3 = 246 × 109 mm3 = 2.46 × 102 × 109 mm3 1 m

= 2.46 × 1011 mm3

uNITS RAISEd TO A POWER

units Raised to a Power

(Units raised to a power are also called squared or cubed unit conversions.)

EXAMPLE

My house is 1044 ft2. How many in2 is this? (12 in = 1 ft, exactly)

Just because 12 in = 1 ft DOESN’T mean 12 in2 = 1 ft2!

Try these problems for more practice on units raised to a power. Show your work!

1. Convert 1018 ft2 to in2 (Answer: 1.466 × 105 in2)

2. There are 3 ft in a yard. How many cubic feet are in one cubic yard?

3. Convert 1 m3 to cm3.

4. Convert 1 m3 to mm3.

5. Convert 1 m3 to μm3.

6. The density of platinum is 21.4 grams per cm3. What is the density of platinum in units of pounds (lb) per ft3? (Conversions: 2.205 lb = 1 kg, 1 in = 2.540 cm, 12 in = 1 ft). (Answer: 1340 lb/ft3)

dENSITY

density

■ density is the ratio of the mass of the substance to the volume of the substance at a given temperature.

■ Density has units of g/cm3 or g/cc or g/mL for liquids and solids, and g/L for gases. ■ Density is an intensive property. It does not depend on the amount of substance. ■ Density varies with change in temperature.

1. A gold-colored ring has a mass of 18.9 grams and a volume of 1.12 mL. Is the ring pure gold? (The density of gold is 19.3 g/mL.)

2. What volume would a 0.871 gram sample of air occupy if the density of air is 1.29 g/L?

3. Pumice is volcanic rock that contains many trapped air bubbles. A 225 gram sample occupied 236.6 mL. a. What is the density of pumice? (Answer is 0.951 g/mL.)

b. Will pumice float on water? (The density of water is 1.0 g/mL.)

4. Which has the greater mass, 1 liter of water or 1 liter of gasoline? The density of water is 1.00 g/mL and that of gasoline is approximately 0.68 g/mL.

dENSITY

5. From their density values, decide whether each of the following substances will sink or float when placed in sea water, which has a density of 1.025 g/mL.

a. Gasoline 0.66 g/mL

b. Mercury 13.6 g/mL

c. Asphalt 1.2 g/mL

d. Cork 0.26 g/mL

6. Mercury is a liquid metal having a density of 13.6 g/mL. What is the volume of 1.00 lb of mercury metal? (33.4 mL)

7. A sample of lead is found to have a mass of 32.6 g. A graduated cylinder contains 2.8 mL of water. After the lead sample is added to the cylinder the water level reads 5.7 mL. Calculate the density of the lead sample. (11g/mL)

8. A piece of magnesium is in the shape of a cylinder with a height of 5.62 cm and a diam- eter of 1.34 cm. If the magnesium sample has a mass of 14.1 g, what is the density of the sample? (1.78 g/mL)

ACTIVITY: THE NuCLEAR ATOM

Activity: The Nuclear Atom

This activity is modified from Chemistry: A Guided Inquiry (3/e) by R.S. Moog and J.J. Farrell, Wiley, 2006.

Model: Schematic diagrams for various atoms and ions.

1 amu = 1.6606 × 10–24 g

Subatomic Particles: Electron (–) Proton (+) Neutron (no charge)

Hydrogen Ion

1.0083 amu

1H–

Hydrogen Atom

1.0078 amu

Hydrogen Atom

2.0140 amu

Exactly 12 amu

6 Protons

12C

Carbon Atom

13.0034 amu

6 Protons

13C

Carbon Atom

7 Neutrons

13.0039 amu

6 Protons

13C–

Carbon Ion

7 Neutrons

15.9960 amu

8 Protons

16O2–

Oxygen Ion

8 Neutrons 11 Protons

22.9893 amu

23Na+

Sodium Ion

12 Neutrons

6 Neutrons

1H and 2H are isotopes of hydrogen.

12C and 13C are isotopes of carbon.

The nucleus of an atom contains the protons and neutrons.

ACTIVITY: THE NuCLEAR ATOM

You will need a periodic table for some of these questions.

1. How many protons are found in

a. 12C? b. 13C? c. 13C–?

2. How many neutrons are found in

a. 12C? b. 13C? c. 13C–?

3. How many electrons are found in

a. 12C? b. 13C? c. 13C–?

4. Based on the model, a. What do all carbon atoms and ions have in common?

b. What do all hydrogen atoms and ions have in common?

5. Look at the periodic table. What does the atomic number (above each atomic symbol) represent in terms of an atom’s subatomic particle?

There are two possible answers; choose the BEST one!

6. Based on your answer to Question 6, what do all nickel (Ni) atoms have in common? Be as specific as possible. Use your periodic table. HINT: Nickel is element 28.

7. 12C and 13C are isotopes of carbon. 1H and 2H are isotopes of hydrogen. What structural feature is different in isotopes of a particular element?

8. In the model, the mass number is shown as the left-hand superscript next to the atomic symbol. How can you determine the mass number from the structure of an atom (its subatomic particles)? In this case, explain how the number 2 is determined for 2H iso- tope based on its subatomic particles. Does that work for all the other isotopes as well?

9. Summary (fill-in-the-blanks): The subatomic particles in the nucleus of an atom are the ___________ and ___________. The ___________ (atomic or mass) number represents the number of protons in an atoms nucleus. The ___________ (atomic or mass) number includes protons and neutrons. Isotopes of an element are the same in the number of ___________ but differ in the number of ___________.

Atomic Number 1

1.0079

Mass Number 2

ACTIVITY: THE NuCLEAR ATOM

Now let’s look at atoms (neutral particles) versus ions (charged particles).

10. Select either the hydrogen or the carbon series and find an example of a symbol which contains a charge (in the upper right of the symbol, usually a – or a +). Compare this to an isotope of the same element which doesn’t have a charge.

a. Write them down here (include the element symbol, mass number, and charge):

b. Using the drawings in Model 1, which kind of subatomic particle (proton, neutron, or electron) is responsible for making atoms into ions (charged particles, shown by – or + sign in its symbol)?

c. In terms of subatomic particles, explain how 13C– got a negative charge compared to 13C (neutral, no charge).

d. In terms of subatomic particles, explain why 23Na+ has a positive charge.

e. How many protons, neutrons, and electrons would be in 1H+? (NOTE: 1H+ is not in the model! See if you can predict what it would be …)

f. In what way are 16O2– and 23Na+ similar in atomic structure?

11. An atomic mass unit (amu) is the unit of mass used for small particles like atoms and ions. The average atomic masses are given in the model for various atoms and ions below the element symbol (e.g., hydrogen’s average atomic mass is 1.0079 amu).

It’s an average mass because it takes into account all the isotopes of hydrogen and their relative amounts in nature.

a. Use the 1H and 2H and 1H– (first row) in the model to determine the approximate mass of each individual sub- atomic particle.

A neutron weighs about ________ amu

An electron weighs about ______ amu

A proton weighs about _________ amu

b. Is most of the mass of the atom inside or outside of the nucleus? Explain.

12. Bonus question: There are 3 isotopes of hydrogen, 1H, 2H (deuterium), and 3H (tritium). If the average mass of all hydrogen isotopes is 1.0079, which isotope is in greatest abun- dance (amounts) in nature?

1.0079(average) Atomic Mass

ACTIVITY: THE NuCLEAR ATOM

EXERCISES (START THESE WHEN YOu FINISH THE ACTIVITY) 1. Complete the following table:

Symbol Atomic Number Mass Number Number of

Protons Number of Neutrons

Number of Electrons

40K

15 17 18

Zn2+ 38

35 81 36

2. Indicate whether the following statement is true or false and explain your reasoning.

An 18O atom contains the same number of protons, neutrons, and electrons.

3. e– is the symbol for electrons. Electrons are gained if present on the left-side of the equation and electrons are lost if present on the right side of the equation. See if you can complete the following equations by filling in the correct ion or the correct number of electrons involved:

a. Mg → ______ + 2e– (HINT: Mg2+ or Mg2–?)

b. F + e– → ______

c. Al → Al3++ ______ (HINT: e– or 2e– or 3e–?)

d. Ca2+ + ______ → Ca

4. How many protons, neutrons, and electrons are found in each of the following?

a. 24Mg

b. 23Na+

c. 35Cl

d. 35Cl–

e. 56Fe3+

f. 15N

g. 16O2–

h. 27Al3+

5. Using grammatically correct sentences, describe what the isotopes of an element have in common and how they are different.

6. Define the terms: atomic number, mass number, average atomic mass, isotope, and ion. Learn them (be able to use them in a sentence).

PERIOdIC TABLE SEARCH

Periodic Table Search

On the periodic tables provided, locate the following and label them using colored pencils, cross-hatching, or arrows.

ON THE FIRST PERIOdIC TABLE 1. Identify the elements of

the periodic table that are

a. metals

b. nonmetals

c. metalloids (semi-metals)

2. Identify the following families (groups):

a. alkali metals (NOTE: Hydrogen is NOT an alkali metal)

b. alkaline earth metals

c. noble gases

d. halogens

3. Identify the location of all the members of Group 3A.

4. Identify the location of all the members of the 4th period.

5. Identify the location of all the members of Group 5A.

6. Identify the location of all the members of the 6th period.

ON THE SECONd PERIOdIC TABLE 7. Identify the transition

elements (B group elements).

8. Locate the inner transition elements.

9. Identify the members of the

a. Lanthanide series

b. Actinide series

10. Find the representative elements (main group elements, A group elements).

ATOMIC THEORY, PERIOdIC TABLE, ANd IONS

Atomic Theory, Periodic Table, and Ions

Some of these questions may not be covered in lecture. Use the internet to search for some of the answers.

1. a. List at least 3 observations or laws that allowed John Dalton to formulate his atomic theory.

b. What are at least 3 postulates of Dalton’s atomic theory?

2. How did Mendeleev construct the first periodic table without knowing how many pro- tons were in each element?

3. How did Ernest Rutherford’s experiments lead to the nuclear theory of the atom?

4. How did J.J. Thomson’s cathode ray experiments illustrate the existence of the electron?

5. Which element in row 1A has the fewest number of protons?

ATOMIC THEORY, PERIOdIC TABLE, ANd IONS

6. The element with 15 protons is in what row?

7. Most element symbols are derived from their names in English. For example, Argon is Ar, Lithium is Li, Phosphorus is P and so on. There are 6 common elements with symbols derived from their Latin name. Write the symbol of each element next to the names below

a. Iron ______

b. Silver ______

c. Tin ______

d. Copper ______

e. Mercury ______

f. Lead ______

8. Where are the noble gases on the periodic table? Why are they called noble gases?

9. Where are the alkali metals on the periodic table? Why are they called alkali metals?

10. Where on the periodic table are the elements that exist as gases at standard tempera- ture and pressure?

11. How are metals different from metalloids? How do we exploit the properties of metal- loids for our computing needs?

ATOMIC THEORY, PERIOdIC TABLE, ANd IONS

12. Determine the charge of each atom described below:

a. A gallium ion with 28 electrons and 38 neutrons

b. An iron ion with 23 electrons

c. An iron ion with 24 electrons and 30 neutrons

d. A sulfur ion with 18 electrons and 16 neutrons

13. An atom with 15 protons, 16 neutrons, and 15 electrons will have what charge?

14. How many protons, neutrons and electrons will be in each atom below:

a. 138Ba2+

b. 138Ba

c. 25Mg2+

d. 37Cl1–

15. How can we predict the charge of ions formed from the main group elements (1A-8A)?

16. What is the common charge for ions of the group 1A elements?

17. What is the common charge for ions of the group 6A elements?

18. Where on the periodic table is the atomic mass listed? What does the atomic mass represent?

ATOMIC STRuCTuRE: ISOTOPES ANd IONS

Atomic Structure: Isotopes and Ions

■ An atom is made up of protons and neutrons (both found in the nucleus) and electrons (in the surrounding electron cloud).

■ Chemical symbol: a one or two letter abbreviation for the element. ■ The atomic number (Z) is equal to the number of protons in the nucleus of an atom. ■ In a neutral atom, the number of protons equals the number of electrons. ■ The mass number (A) is equal to the number of protons plus neutrons. ■ The (average) atomic mass of an element is the weighted average mass of an element

considering its isotopes and their natural abundances.

■ The charge on an ion indicates an imbalance between protons and electrons. Too many electrons produces a negative charge, too few, a positive charge.

■ Isotopes of the same element have the same number of protons but different numbers of neutrons.

Isotope Symbol Notation: AZ X Isotope Name: element-A

1. The atomic number of Zn is _______ and represents the number of _______ in a Zn atom.

2. The mass number of bromine-79 is _______ and it contains _______ protons and _______ electrons.

3. It has _______ (how many?) fewer neu- trons than bromine-81.

4. The average atomic mass of bromine is _______ amu.

5. 39K is an atom, but when it loses an electron, it will have a charge of _______.

6. In an atom of 24Mg, the mass number is _______, the atomic number is _______ and the charge is _______.

Complete the following chart:

Element/Ion Atomic Number Mass Number Charge Protons Electrons Neutrons

14C

39K

26Mg2+

75As3–

109Ag1+

235U

ISOTOPE CALCuLATIONS

Isotope Calculations

1. Estimate the atomic mass of copper. Copper has two isotopes, copper-63 (69.15% abundance) and copper-65 (30.85% abundance). (HINT: What is the approximate mass of Cu-63 in amu?)

2. An element exists as four different isotopes. Refer to the data table for the masses and abundances. Calculate the weighted average atomic mass of this element. (Don’t for- get units.)

Isotope Mass (amu) Percent Abundance

49.946 4.35

51.941 83.79

52.941 9.50

53.929 2.37

3. Gallium has two isotopes with masses of 68.926 amu and 70.925 amu. What are the nat- ural abundances for these two isotopes? (Answer: 60.13% and 39.87%)

4. Magnesium has three isotopes. One isotope has a mass of 24.986 amu and a natu- ral abundance of 10.00%. Another isotope has a mass of 25.983 amu and a natural abundance of 11.01%. Determine the mass and natural abundance of the third isotope. (Answer: 23.98 amu, 78.99%)

ATOMIC STRuCTuRE REVIEW

Atomic Structure Review

1. Match the following models/discoveries/theory/Law with the related person/experi- ment:

i. Idea of the atom _______

ii. Plum Pudding Model _______

iii. Nuclear Atom _______

iv. Atomic Theory _______

v. Law of Conservation of Mass _______

a. Gold foil experiment

b. Antoine Lavoisier

c. Cathode Ray Tube/Discovery of the Electron

d. Democritus

e. John Dalton

2. Each of these drawings includes subatomic particles, which represents an atom or ion. The total mass of the atom or ion is shown in units of atomic mass units in the box (amu).

Legend: Electron (–) Proton (+) Neutron (no charge)

Hydrogen Ion

1.0083 amu

1H–

Hydrogen Atom

1.0078 amu

Hydrogen Atom

2.0140 amu

a. Which two drawings will help you estimate the mass of an electron? Calculate it and include units.*

b. Which two drawings will help you estimate the mass of a neutron? Calculate it and include units.*

c. Which drawing can you use (with the answer to part a or b) to estimate the mass of a proton? Calculate it and include units.*

d. Based on your answers, where is most of the mass of the atom, inside or outside the nucleus?

e. Which of these boxes represents an ion? Explain.

*NOTE: The precise masses cannot be determined this way due to the binding energy of the nuclear particles.

ATOMIC STRuCTuRE REVIEW

3. Which of the following descriptions of a subatomic particle is correct?

a. A proton has a positive charge and a mass of approximately 1 amu.

b. An electron has a negative charge and a mass of approximately 1 amu.

c. A neutron has no charge and its mass is negligible.

d. A proton has a positive charge and a negligible mass.

4. Write the chemical symbol or name associated with the element and list the numbers of protons, neutrons and electrons in the neutral atom

Name Symbol # of Protons # of Neutrons # of Electrons

Phosphorus 16

81Br−

5. Describe each element as a metal, semimetal, or non metal (in the first blank). In the second blank, also classify each as one of the following: alkali metal, alkaline earth metal, halogen, noble gas, transition metal.

a. Ag ______________________ also a ______________________

b. Cl ______________________ also a ______________________

6. Name an element in Group 6A: ____________. Name an element in Period 3: ____________.

7. Name an element that might have the same chemical properties as calcium. ____________

8. a. What observation in the Gold Foil experiment led to the conclusion that the atom contains a dense nucleus?

b. What observation in the Gold Foil experiment led to the conclusion that the atom is composed of mostly empty space and the nucleus is very small?

9. Magnesium exists in nature as three different isotopes. Calculate the atomic mass of magnesium from the data given. You must show work for credit. Use significant figures.

Magnesium

Atomic Mass (amu) Mg-24 23.985045 Mg-25 24.985839 Mg-26 25.982595

% Abundance 78.90 10.00 11.10

FORMINg IONS

Forming Ions

1. Noble gases are in Group _______ and are fairly stable compared to other elements. (Why?)

2. Chemical reactivity arises from the organization of the electrons. A simplified model (developed by Niels Bohr) has electrons filling orbits around the nucleus. The number of electrons which fill each orbit is 2, 8, 8, 18 (follows the periods in the periodic table). Use the periodic table to fill in the Bohr diagrams on the next page.

3. Which element on the next page is a noble gas? What is unique about its Bohr diagram compared to the others?

4. In order to gain stability, atoms gain or lose electrons to fill their outermost electron shells. This shell is called the “valence shell.”

5. For each of the elements on the next page, determine how many electrons would be gained or lost to have a filled valence shell. Which is more likely, a loss or gain of electrons?

6. If atoms lose electrons to fill their valence shell, will the resulting ion be positive or negative?

7. If atoms gain electrons, will the resulting ion be a cation or an anion?

8. For each of the elements on the next page, determine the symbol for the most likely ion to result (in order to have a filled valence shell). For example, sodium will likely become a Na+ cation.

PRACTICE

1. Draw the Bohr diagram for phosphorus.

2. What ion does phosphorus become?

3. Compare phosphorus with nitrogen (on the next page). In what ways are they similar?

4. What type of ions (on the next page) do you think will bond with phosphorus? (HINT: Opposite charges attract!)

5. Is there a way to predict for any element if it becomes a positively charge or a nega- tively charged ion? (HINT: Look at its location on the periodic table)

BOHR MOdEL dIAgRAMS

Bohr Model diagrams

Sodium Neon

Chlorine Oxygen

Aluminum Nitrogen

CHARgES OF IONS

Charges of Ions

Get familiar with the ionic charge of various elements when they become ions.

1. Draw demarcation line for metals vs. non-metals (exception: hydrogen is a non-metal).

2. Number columns for groups IA–VIIIA (1A–8A).

3. Show fixed ionic charges (for metals and non-metals):

NAMES OF NON-METAL ANIONS (NEgATIVELY-CHARgEd IONS) For non-metal ions (charged atoms) the –ide ending is used after the first syllable of the element.

H– hydride

N3– nitride

O2– oxide

F– fluoride

P3– phosphide

S2– sulfide

Cl– chloride

As3– arsenide

Se2– selenide

Br– bromide

Te2– telluride

I– iodide

NOTE: H can have a negative charge (in which case it is called “hydride”) OR a positive charge (in which case it is called “hydrogen ion”). The elements in Group IVA (4A) can also make anions, C4– is carbide and Si4– is silicide.

CHARgES OF IONS

NAMES OF METAL CATIONS (POSITIVELY-CHARgEd IONS) For metal ions (charged atoms) the name of the element is followed by “ion.”

EXAMPLE

Atom/Element Ion Name of Ion

Mg Mg2+ magnesium ion

Ag Ag+ silver ion

Some elements can exist in several forms. For example, chlorine can exist as neutral atoms, diatomic molecules, or ions.

chlorine atom Cl chlorine Cl2 chloride ion Cl –

Ionic charge is always shown as a SUPERSCRIPT for ions, otherwise the atom is assumed to be neutral. When the charge is neutralized, the superscripts are not written (e.g., Na+Cl– becomes NaCl).

PRACTICE

Write the formula and the name of the ion that is formed for each element.

Atom/ Element

Ion Formula (show charge) Cation? Anion? Name of Ion

TuTORIAL: IONIC NOMENCLATuRE

Tutorial: Ionic Nomenclature

PART I: FORMuLAS ANd NOMENCLATuRE OF IONIC COMPOuNd COMPOSEd OF CATIONS ANd ANIONS

TYPES OF CATIONS (POSITIVE IONS)

1. Metals lose electrons to form positive ions. These ions are called monoatomic because they are made up of only ONE ion. They can be of two types: Constant charge or Variable charge.

a. Constant charge

Group IA, IIA, IIIA

A few transition metals Ag+, Zn2+ and Cd2+

The names of these ions are the same as the name of the atom.

Examples: Na+ = sodium ion, Zn2+ = zinc ion

b. Variable charge

Most transition metals (except for silver, zinc and cadmium)

A few representative metals: Sn, Bi, and Pb

Fe2+ Iron (II)

Fe3+ Iron (III)

Cu+ Copper (I)

Cu2+ Copper (II)

Sn4+ Tin (IV)

Sn2+ Tin (II)

2. A polyatomic ion: Consists of more than one atom. The most common positive one is ammonium, NH4

TYPES OF ANIONS (NEgATIVE IONS)

1. Nonmetals gain electrons to form negative ions. These ions are called monoatomic because they are made up of only ONE ion. The names of these ions end in –ide.

Examples: S2– = sulfide, Cl1– = chloride, N3– = nitride, O2– = oxide

2. A polyatomic ion: Consists of more than one atom. The names of polyatomic anions often end in –ate or –ite.

Examples: NO3 1– = nitrate, NO2

1– = nitrite, SO4 2– = sulfate, SO3

2– = sulfite

Two important polyatomic ions end in –ide: OH1– = hydroxide, CN1– = cyanide

TuTORIAL: IONIC NOMENCLATuRE

Determine if the compound is ionic: Contains a cation and an anion—often a metal and a nonmetal.

HINT: Formula begins with a metal or NH4 +

IONIC: TO WRITE THE FORMuLA

1. Write the symbol of the cation with a SUPERSCRIPT charge.

2. Write the symbol of the anion with a SUPERSCRIPT charge.

3. If the charges are NOT balanced, CRISS-CROSS to find the number of each atom nec- essary to balance the charges.

4. Use a parenthesis if more than ONE polyatomic ion is necessary.

EXAMPLES

■ Aluminum oxide Al3+ O2– Al2O3 ■ Iron (II) phosphide Fe2+ P3– Fe3P2 ■ Calcium nitrate Ca2+ NO31– Ca(NO3)2 ■ Copper (I) carbonate Cu1+ CO32 – Cu2CO3 ■ Barium oxide Ba2+ O2– BaO

NOTE: The charges are already balanced so the ratio is 1:1.

■ Calcium hydroxide Ca2+ OH1– Ca(OH)2

IONIC: TO WRITE THE NAME

1. Write the name of the cation.

If the cation has a variable charge, determine the Roman numeral.

2. Follow with the name of the anion.

EXAMPLES

■ Na2O Sodium oxide (Group IA—no Roman numeral needed) ■ Cr2S3 Chromium (III) sulfide (Most transition metals need Roman numerals.) ■ BaSO4 Barium sulfate (Group IIA—no Roman numeral needed) ■ Pb(OH)2 Lead (II) hydroxide (Pb is a main group metal with variable charge.) ■ Cu2CO3 Copper (I) carbonate (Most transition metals need Roman numerals.) ■ FeSO4 Iron (II) sulfate (Most transition metals need Roman numerals.)

POLYATOMIC IONS TO MEMORIZE

Polyatomic Ions to Memorize

Formula Name Formula Name

NO2 – nitrite HCO3

– hydrogen carbonate (or bicarbonate)

NO3 – nitrate ClO– hypochlorite

SO3 2– sulfite ClO2

– chlorite

HSO3 – hydrogen sulfite

(or bisulfite) ClO3 – chlorate

HSO4 – hydrogen sulfate

(or bisulfate) ClO4 – perchlorate

SO4 2– sulfate BrO– hypobromite

PO3 3– phosphite BrO2

– bromite

PO4 3– phosphate BrO3

– bromate

HPO4 2– hydrogen phosphate BrO4

– perbromate

H2PO4 – dihydrogen phosphate IO– hypoiodite

OH– hydroxide IO2 – iodite

H3O + hydronium IO3

– iodate

NH4 + ammonium IO4

– periodate

CO3 2– carbonate

Other polyatomic ions you may encounter (not required to memorize)

Formula Name Formula Name

CrO4 2– chromate C6H5O7

3– citrate

Cr2O7 2– dichromate BO3

3– borate

CN– cyanide MnO4 – permanganate

C2H3O2 – (or

CH3COO –) acetate O2

2– peroxide

C2O4 2– oxalate Hg2

2+ mercury (I)

MEMORIZINg POLYATOMIC IONS: THE T43 METHOd

Memorizing Polyatomic Ions: The T43 Method

Many of the polyatomic ion charges and number of oxygens can be memorized using the “T43 method.” If your instructor does not cover this, try a search on “T43 method” in YouTube to see how it’s used.

uSINg A PERIOdIC TABLE 1. Draw a “T” over the B, C, N and Si.

The “T” elements indicate 3 oxygens and a charge of –3, –2, –1 respectively (left to right).

For example, the polyatomic ions you get are: BO3

3–, CO3 2–, and NO3

–. These end in “ate”: borate, carbonate, and nitrate. Ignore silicon—you will not see it very often in this course.

2. Draw a “4” over the P, As/S, Se/Te.

The “4” elements contain 4 oxygens and a charge of –3 and –2, respectively (left to right). These end in “ate”: phos- phate and sulfate.

3. Draw a “3” connecting the Cl, Br, and I.

These “3” elements contain 3 oxygens and a –1 charge: ClO3 –, BrO3

–, and IO3 –. These

end in “ate”: chlorate, bromate, iodate.

PRACTICE

Provide the formulas for these polyatomic ions using the T43 method (don’t forget the charge!)

1. Bromate

2. Carbonate

3. Nitrate

4. Phosphate

5. Sulfate

6. Iodate

Thallium Lead Bismuth Astatine RadonPolonium

Aluminum Silicon Phosphorus Chlorine ArgonSulfur

Gallium

Indium

Germanium Arsenic Bromine KryptonSelenium

Tin Antimony Iodine XenonTellurium

Boron Carbon Nitrogen Fluorine Neon

Helium

Oxygen

26.981538 28.0855 30.973761 32.066 35.4527 39.948

10.811 12.0107 14.00674 15.9994 18.9984032 20.1797

4.003

114.818 118.710 121.760 127.60 126.90447 131.29

69.723 72.61 74.92160 78.96 79.904 83.80

204.3833 207.2 208.98038 (209) (210) (222)

dERIVINg POLYATOMIC IONS

deriving Polyatomic Ions

You need to know the names and formulas (including charges) for all 27 of these polyatomic ions.

You can do this by memorizing only 7 of them and then deriving the rest! Here’s how:

Memorize these two:

Halogen oxyanions: XO# –

(X can be Cl, Br, or I; names will use “chlor” or “brom” or “iod”)

→ per …ate perchlorate increase –ate formula by one more O atom ClO4

–

Memorize this one:

chlorate ClO3 –

→ –ite chlorite reduce –ate formula

by one O atom ClO2 –

→ hypo …ite hypochlorite reduce –ite formula by one more O atom ClO

–

Formula Name

NO2 – nitrite

NO3 – nitrate

SO3 2– sulfite

HSO3 – hydrogen sulfite (bisulfite)

HSO4 – hydrogen sulfate (bisulfate)

SO4 2– sulfate

PO3 3– phosphite

PO4 3– phosphate

HPO4 2– hydrogen phosphate

H2PO4 – dihydrogen phosphate

OH– hydroxide

H3O + hydronium

NH4 + ammonium

CO3 2– carbonate

HCO3 – hydrogen carbonate (bicarbonate)

ClO– hypochlorite

ClO2 – chlorite

ClO3 – chlorate

ClO4 – perchlorate

BrO– hypobromite

BrO2 – bromite

BrO3 – bromate

BrO4 – perbromate

IO– hypoiodite

IO2 – iodite

IO3 – iodate

IO4 – periodate

O H H

N+

Hydroxide OH– Ammonium NH4 +

dERIVINg POLYATOMIC IONS

Memorize These, use to derive Related Ions

Nitrate Carbonate Sulfate Phosphate

NO3 – CO3

2– SO4 2– PO4

3–

O N+

O O

O S O

Related Ions

→ –ite nitrite — sulfite phosphite reduce –ate formula

by one O atom NO2 – — SO3

2– PO3 3–

→ hydrogen …ate — hydrogen

carbonate (or bicarbonate)

hydrogen sulfate (or bisulfate)

hydrogen phosphate

add H+ to –ate formula (add H and increase charge by 1) — HCO3

– HSO4 – HPO4

2–

→ dihydrogen …ate — — — dihydrogen phosphate add two H+ to –ate formula — — — H2PO4

–

→ hydrogen …ite — — hydrogen sulfite — reduce –ate formula by one O atom, add one H+ — — HSO3

– —

Related Acids

→ acid form of …ate –ate ending changes

to –ic acid nitric acid carbonic acid sulfuric acid phosphoric acid

add enough H+ to –ate formula until neutral charge HNO3 H2CO3 H2SO4 H3PO4

→ acid form of …ite –ite ending changes

to –ous acid nitrous acid — sulfurous acid phosphorous acid

add enough H+ to –ate formula until neutral charge HNO2 — H2SO3 H3PO3

COMBININg IONS

Combining Ions

1. Give the charges on the ions of the following elements:

a. Ca

b. Rb

c. Al

d. N

e. Se

f. Br

2. What will the charge be on elements of the following groups?

a. Group 1

b. Group 13

c. Group 15

d. Group 16

3. Write the names of the following metal cations: (NOTE: Use roman numerals in paren- thesis to indicate charge on metals with a variable charge. Roman numerals are not used for metals with a constant charge.)

a. Li+ __________________

b. Ca2+ _________________

c. Fe2+ _________________

d. Cu2+ _________________

e. Cu+ _________________

f. Zn2+ _________________

4. Write the names of the following nonmetal cations: (NOTE: Use the root name of the nonmetal and the ending -ide.)

a. O2– __________________

b. Cl– __________________

c. N3– _________________

d. S2– __________________

e. F– ___________________

f. P3– __________________

5. Give the formula of the ionic compounds formed between the following ions and name them.)

a. Na+ and Br– ________________________

b. Sr2+ and I– _________________________

c. Li+ and O2– _________________________

d. Cs+ and N3– ________________________

e. Aluminum and sulfur _______________

f. Calcium and oxygen ________________

g. Barium and chlorine ________________

h. Potassium and phosphorus _________

IONIC COMPOuNd MATRIX

Ionic Compound Matrix

Fill in the matrix with the chemical formula derived from the pairing of each cation and anion. Write its name underneath the chemical formula.

anion → cation ↓ chloride oxide nitrate sulfite phosphate

sodium

calcium

aluminum

iron (II)

iron (III)

copper (I)

copper (II)

chromium (II)

silver

zinc

ammonium

IONIC NOMENCLATuRE

Ionic Nomenclature

1. Binary Ionic Compounds with Constant Charge Metal Cations

a. NaBr _________________________

b. Zn3P2 _________________________

c. CaS __________________________

d. AlP ___________________________

e. Na2O _________________________

f. Rb3N _________________________

g. AlF3 __________________________

h. cesium iodide __________________

i. zinc nitride _____________________

j. potassium sulfide _______________

k. silver oxide ____________________

l. calcium bromide ________________

m. strontium chloride _______________

n. sodium hydride _________________

2. Binary Ionic Compounds with Variable Charge Metal Cations

a. Cu2O _________________________

b. SnS2 __________________________

c. Pb3N2 _________________________

d. NiCl2 _________________________

e. FeCl3 _________________________

f. SnCl4 _________________________

g. CoS __________________________

h. Cr2S3 _________________________

i. manganese (VII) arsenide _________

j. platinum (II) sulfide ______________

k. vanadium (V) nitride _____________

l. gallium (III) oxide ________________

m. tin (II) fluoride __________________

n. copper (II) chloride ______________

o. chromium (III) oxide ______________

p. titanium (II) selenide _____________

3. 7 Polyatomic Ions

a. NO3 – ________________________________

b. CO3 2– _______________________________

c. SO4 2– _______________________________

d. PO4 3– _______________________________

e. ClO4 – ________________________________

f. OH– _________________________________

g. NH4 + ________________________________

4. Mixed Polyatomic Ions

a. NO2 – ________________________________

b. PO4 3– _______________________________

c. SO3 2– _______________________________

d. ClO3 – ________________________________

e. NH4 + ________________________________

f. HSO4 ________________________________

g. BrO4 – _______________________________

h. IO2 – _________________________________

i. SO4 2– _______________________________

j. BrO– ________________________________

k. PO3 3– _______________________________

l. ClO– ________________________________

m. BrO2 – _______________________________

n. NO3 – ________________________________

o. HPO4 2– ______________________________

p. IO3 – _________________________________

q. H2PO4 – ______________________________

r. BrO3 – _______________________________

IONIC NOMENCLATuRE

s. OH– _________________________________

t. IO– __________________________________

u. HCO3 – _______________________________

v. IO4 – _________________________________

w. ClO2 – ________________________________

x. CO3 2– _______________________________

y. HSO3 – _______________________________

z. ClO4 – _______________________________

5. Ionic Compounds with Polyatomic Ions & Constant Charge Ions

a. CaCO3 _______________________________

b. KClO4 _______________________________

c. CdSO3_______________________________

d. Li2SO3_______________________________

e. Ag3PO4 _____________________________

f. NH4F ________________________________

g. Al(NO2)3 ____________________________

h. Ba(HCO3)2 ___________________________

i. potassium hydroxide ________________

j. beryllium hydroxide _________________

k. strontium acetate ___________________

l. ammonium oxide ____________________

m. silver cyanide _______________________

n. zinc carbonate ______________________

o. aluminum bicarbonate _______________

p. magnesium phosphate ______________

6. Ionic Compounds with Variable Charge Metal Cations & Polyatomic Ions

a. Fe(HCO3)2 ___________________________

b. NiPO4 _______________________________

c. V2(SO4)3 ____________________________

d. Ti(CN)4 ______________________________

e. Sc(OH)3 _____________________________

f. Cu(NO2)2 ____________________________

g. CoCO3 ______________________________

h. Sn(HPO4)2 ___________________________

i. copper (II) chlorate __________________

j. cobalt (III) chromate _________________

k. lead (IV) sulfate _____________________

l. iron (II) phosphate ___________________

m. chromium (II) carbonate _____________

n. iron (II) cyanide _____________________

o. nickel (II) sulfite _____________________

p. tin (IV) chlorate _____________________

7. Mixed Ionic Compounds

a. tin (II) phosphate ___________________

b. zinc oxide __________________________

c. barium sulfate ______________________

d. lithium nitride _______________________

e. silver fluoride _______________________

f. barium hydroxide ___________________

g. lead (II) iodide ______________________

h. mercury (II) chloride ________________

i. cobalt (II) nitrate ____________________

j. lithium bromide _____________________

k. AlBr3 _______________________________

l. FeS ________________________________

m. (NH4)3PO4 __________________________

n. Hg(NO3)2 ___________________________

o. Ag2S _______________________________

p. KMnO4 _____________________________

q. MgCl2 ______________________________

r. Cr2O3 _______________________________

s. K3P _________________________________

t. TiCl2 ________________________________

NOMENCLATuRE PRACTICE TEST 1

Nomenclature Practice Test 1

1. Write the formulas of the following ionic compounds:

a. tin (II) phosphate ___________________

b. zinc oxide __________________________

c. barium sulfate ______________________

d. lithium nitride _______________________

e. silver fluoride _______________________

f. barium hydroxide ___________________

g. lead (II) iodide ______________________

h. mercury (II) chloride ________________

i. cobalt (II) nitrate ____________________

j. lithium bromide _____________________

2. Write the names of the following ionic compounds:

a. AlBr3 _______________________________

b. FeS _________________________________

c. (NH4)3PO4 __________________________

d. Hg(NO3)2 ___________________________

e. Ag2S _______________________________

f. KMnO4 _____________________________

g. MgCl2 ______________________________

h. Cr2O3 _______________________________

i. K3P _________________________________

j. TiCl2 ________________________________

Nomenclature Practice Test 2

1. Write the formulas of the following ionic compounds:

a. iron (II) arsenide ____________________

b. lead (II) sulfate _____________________

c. lead (IV) hydroxide __________________

d. copper (II) acetate __________________

e. beryllium chloride __________________

f. ammonium chromate _______________

g. silver oxide _________________________

h. potassium sulfide ___________________

2. Write the names of the following ionic compounds:

a. KI __________________________________

b. Mn2(SO3)7 __________________________

c. SnBr4 _______________________________

d. Mg3P2 ______________________________

e. NaF ________________________________

f. Sr(MnO4)2 __________________________

g. Cr3(PO4)2 ___________________________

h. Al2Se3 ______________________________

COVALENT COMPOuNdS

Covalent Compounds

Covalent compounds (also called “Molecular” compounds)

Now we will learn how to name COVALENT compounds of this form:

Non-Metal + Non-Metal

Covalent Compounds do not form from ions. They are from atoms which share electrons in order to mutually fill their valence shells. You will learn more about this toward the end of the quarter.

For example, covalent bonding explains why some elements are diatomic (BrINClHOF).

Use these templates to complete a Bohr model diagram for Cl2. Treat each of these two templates as a separate chlorine atom.

Notice that both chlorine atoms do not have a filled shell. They can obtain a filled shell by sharing electrons with each other. There are two electrons shared between them.

COVALENT NOMENCLATuRE

Covalent Nomenclature

Covalent compounds do not form from ions. Therefore, you simply have to state the chem- ical formula using words and prefixes, and continue using binary endings.

EXAMPLE

SiCl4 is silicon tetrachloride (NOTE: Mono is usually omitted from the first element.)

Prefixes for Covalent Compounds (memorize)

Number Prefix Number Prefix

1 mono 6 hexa

2 di 7 hepta

3 tri 8 octa

4 tetra 9 nona

5 penta 10 deca

You should also know the common names for three compounds:

■ H2O (water) ■ CH4 (methane) ■ NH3 (ammonia)

1. XeO3

2. OF2

3. SO3

4. P2O5

5. Cl2O

6. NI3

7. NH3

8. dihydrogen monoxide

9. xenon tetrafluoride

10. diiodine pentoxide

11. chlorine monoxide

12. carbon tetrafluoride

13. tetraphosphorus triselenide

14. methane

IONIC OR MOLECuLAR?

Ionic or Molecular?

For each of the following, indicate whether it is an ionic compound (I) or a molecule (M), and give the corresponding formula or name as appropriate.

I or M Chemical Formula/Name

1. aluminum acetate _____ _____________________________________________

2. silver phosphate _____ _____________________________________________

3. ammonia _____ _____________________________________________

4. copper (II) hydroxide _____ _____________________________________________

5. ammonium sulfate _____ _____________________________________________

6. calcium perchlorate _____ _____________________________________________

7. phosphorus trichloride _____ _____________________________________________

8. titanium (IV) oxide _____ _____________________________________________

9. zinc hypochlorite _____ _____________________________________________

10. diarsenic pentasulfide _____ _____________________________________________

11. Na2O2 _____ _____________________________________________

12. P4O7 _____ _____________________________________________

13. Hg2CO3 _____ _____________________________________________

14. Ni(OH)3 _____ _____________________________________________

15. N2O5 _____ _____________________________________________

16. Sn(CrO4)2 _____ _____________________________________________

17. SO3 _____ _____________________________________________

18. Mg3N2 _____ _____________________________________________

19. GeO2 _____ _____________________________________________

20. Al2(SO4)3 _____ _____________________________________________

THE MOLE

The Mole

PART 1 Here are some common counting units relating to the numbers of things.

1 dozen = 12 1 score = 20 1 ream = 500 1 gross = 144

Practice converting between units of measure and the items they represent. Use dimen- sional analysis.

1. 24 eggs = __________ dozen eggs

2. 4 score of years = __________ years

3. 750 sheets = __________ reams of paper

4. 2 gross of pencils = __________ pencils

PART 2 In chemistry, it is convenient to group small particles (atoms, molecules, ions) into a count- ing unit called the mole.

1 mole = 6.022 × 1023 objects (This number is called Avogadro’s number.)

When you use your calculator, try using the EE, E or EXP function as a shortcut for the exponent in scientific notation “ × 10 ^ ” (times 10 to the power of). Use dimensional anal- ysis for each one.

5. 24 atoms of sodium = __________ moles of sodium atoms

6. 15.0 moles of Pb2+ ions = __________ Pb2+ ions

7. 950 atoms of silver = __________ moles of silver atoms

8. 5.36 moles of NH3 = __________ molecules of NH3 molecules

9. 5 molecules of chlorine (Cl2) gas = __________ moles of chlorine (Cl2) molecules

10. 4.00 moles of barium atoms = __________ atoms of barium

11. Why is a mole such a big counting unit compared to ones we use every day (such as the dozen)?

MASS, MOLES, AVOgAdRO’S NuMBER, ANd MOLES WITHIN MOLES CONVERSIONS

Mass, Moles, Avogadro’s Number, and Moles Within Moles Conversions

g of atomic element “A”

Moles of element “A”

Moles within moles

Chemical formula of

compound “B”

Number of ions (ionic compound) or

number of atoms (molecular compound)

of element “A” MMA

MMA

g of compound “B”

Moles of compound “B”

Number of molecules (molecular compound) or number of formula units

(ionic compound) MMB

MMB

Atoms within a molecule or formula unit

Chemical formula of

compound “B”

ABBREVIATIONS uSEd

MM = Molar mass

NA = Avogadro’s number = 6.02 × 10 23 items

Moles = Mol

Molecules = mc/mcs

MOLAR MASS

Molar Mass

■ The molar mass of a substance = the mass (in grams) of one mole of the substance. ■ One mole of an element = the atomic mass of that element (on the periodic table)

in grams.

■ One mole of a compound = the sum of the atomic masses of the atoms present in the compound in grams.

■ Example: H2O = (2 × 1.008) + 16.00 = 18.016 g/mol. ■ The units of molar mass are always grams per mole (g/mol). ■ NOTE: “Mole” may be abbreviated “mol,” but not “m” (“m” means meter).

1. What is the mass of one mole (molar mass) of Ar?

2. What is the molar mass of Na?

3. What is the mass of one mole (molar mass) of H2O?

4. What is the molar mass of NaCl?

5. How many moles of H2O are in 22.5 g of H2O?

6. How many grams are in 0.250 moles of NaCl?

7. What is the molar mass of C2H5OH (ethanol)?

8. How many moles are in 25.0 mL of ethanol, C2H5OH (the density of ethanol is 0.785 g/mL)? (Answer: 0.426 mol)

TYPES OF SuBSTANCES

Types of Substances

Most elements occur as monatomic elements (single atoms, such as Cu or Na). Some ele- ments are diatomic and exist as molecules of two atoms each (BriNClHOF). Other elements may be molecular as well (S8 or C60)—but you don’t need to know those.

Compounds are either ionic or covalent. If ionic, their basic unit is the “formula unit” because they consist of extending arrays of ions in a crystal lattice (diagram below, left). Covalent or molecular compounds have a basic unit of “molecules” because they are dis- crete or separated (diagram below, right).

NaCl crystal of Na+ and Cl– ions exist as formula units of “NaCl” (simplest ratio).

Water molecules which exist as “H2O.”

The mass of one atom, molecule, or formula unit is most appropriate in units of atomic mass units (amu).The molar mass of these substances is in units of grams per mole. The numbers are from the atomic masses on the periodic table.

Find the formula and molar mass of the following. Use at least two decimal places to ensure enough significant figures for most problems. Indicate the simplest unit as atom, molecule, or formula unit. The first is completed as an example.

Substance Formula Mass (amu) Molar Mass (g/mol) Simplest unit

CO2 44.01 amu 44.01 g/mol Molecule

Al2(SO4)3

C6H10O5

Br2

CaCl2

H2O

CaCl2 · 2 H2O (a hydrate; dot means to add 2 H2O)

MOLES, gRAMS, NuMBER OF PARTICLES, ELEMENTS IN COMPOuNdS

Moles, grams, Number of Particles, Elements in Compounds

Complete this concept map by filling in “Avogadro’s number” or “Molar Mass” over the arrows.

# of particles (atom, ions, molecules, formula units) ←→ moles ←→

mass (grams)

1. How many moles of magnesium are there in 10.00 g?

2. How many grams of iron are there in 34.77 moles of iron?

3. How many moles are there in 3.493 grams of lye, NaOH?

4. How many grams are there in 0.275 moles of Na3PO4?

MOLES, gRAMS, NuMBER OF PARTICLES, ELEMENTS IN COMPOuNdS

5. How many molecules are in 7.47 moles of NH3? How many total atoms are there? (Answer: 1.80 × 1025 atoms)

6. What is the mass in grams of 1.000 × 1012 (1.000 trillion) atoms of gold?

7. How many grams of ammonium nitrate, NH4NO3, contain 8.3 × 10 25 formula units?

8. How many molecules of water are there in 0.034 grams of water?

MOLES: CALCuLATE THE MISSINg QuANTITIES

Moles: Calculate the Missing Quantities

Fill in the blanks.

You need to use 1–2 extra digits (beyond # of sig figs in given quantity) in your molar mass to not cause a rounding error.

Chemical Formula

Molar Mass (g/mol) Mass Moles

# of Particles (atoms, molecules,

or formula units)

Type of Substance (monatomic element, molecular element, covalent compound,

ionic compound)

N2 2.6 × 10–4

moles

CH3OH 1.25 g

Na 8.170 × 10 25

Na atoms

CaF2 0.0844 moles

KNO3 19.3 kg

0.4395 g 2.75 × 10 –3

moles diatomic molecular

element

314 g 2.26 × 10 24

atoms monatomic

element

MOLE RATIOS uSINg CHEMICAL FORMuLAS

Mole Ratios using Chemical Formulas

An elephant has one trunk and four legs.

A molecule of methane, CH4 contains one carbon and four hydrogens.

A ratio can be constructed for these two objects in the following ways:

4 legs 4 H atoms 1 elephant 1 CH4 molecule

1. Can you think of another ratio for each of the above? Write it next to the ones above.

2. Now picture a dozen elephants and a dozen methane molecules. What ratios can you construct using the term ”dozen” in the numerator and denominator? Fill in the numer- ators below:

1 dozen elephants 1 dozen CH4 molecules

3. Now picture you have a mole (rather than a dozen) elephants and methane molecules. Fill in the numerator now:

1 mole of elephants 1 mole of CH4 molecules

4. Practice constructing a mole ratio between the indicated element in one mole of the each of the following compounds(first one is completed as an example):

a. Oxygen atoms (O) in CO2 Answer: 2 mole O atoms/1 mole CO2

b. Carbon in Glucose (C6H12O6)

c. Chloride ions (Cl–) in CaCl2

5. Practice using the conversions above for the given amounts. Use dimensional analysis!

a. Moles of oxygen (O) in 4 moles of CO2

b. Moles of carbon in 0.45 moles of glucose, C6H12O6

c. Moles of chloride ions in 5.00 grams of CaCl2

PRACTICINg MOLE RATIOS

Practicing Mole Ratios

1. One mole of carbon dioxide contains _____ moles of C and _____ moles of O.

Rewrite this statement as two mole ratios:

2. 88.02 grams of carbon dioxide is _____ mole(s) of CO2, which is _____ mole(s) of carbon atoms, and _____ mole(s) of oxygen atoms.

3. a. Perform this calculation, using some information you already provided in the previ- ous question.

88.02 grams of carbon dioxide is also _____ grams of carbon atoms and _____ grams of oxygen atoms.

Show your work:

b. Do the grams of carbon atoms and grams of oxygen atoms add up to the grams of carbon dioxide? Should they?

4. Indicate whether each of the following statements is true or false, and explain your reasoning.

a. One mole of NH3 weighs more than one mole of H2O.

b. There are more carbon atoms in 48 grams of carbon dioxide than in 12 grams of dia- mond (pure carbon).

c. There are equal numbers of nitrogen atoms in one mole of NH3 and one mole of N2.

d. The number of Cu atoms in 100 grams of Cu (s) is the same as the number of Cu atoms in 100 grams of copper (II) oxide, CuO.

e. The number of Ni atoms in 100 moles of Ni (s) is the same as the number of Cl atoms in 100 moles of nickel (II) chloride, NiCl2.

f. There are more hydrogen atoms in 2 moles of NH3 than in 2 moles of CH4.

MOLES WITHIN MOLES

Moles within Moles

We can interpret 1 mole of a compound or molecules in terms of moles of elements.

EXAMPLES

■ 1 mole of (NH4)2SO4 = 2 moles of nitrogen ■ 1 mole of (NH4)2SO4 = 8 moles of hydrogen ■ 1 mole of (NH4)2SO4 = 1 mole of sulfur ■ 1 mole of (NH4)2SO4 = 4 moles of oxygen

We can use the analogy of an elephant. We can say 1 mole of elephants contains 1 mole of trunks, 2 moles of ears and 4 moles of legs.

1. How many moles of oxygen atoms are there in 2 moles of KNO3? (6 mol)

2. How many hydrogen atoms are there in 2.50 moles of (NH4)2SO4 formula units? (1.20 × 1025 H atoms)

3. A mole of H2O and a mole of O2

a. have the same mass.

b. contain one molecule each.

c. have a mass of 1 g each.

d. contain the same number of molecules.

4. One molecule of phosphorus contains 4 phosphorus atoms. Then one mole of phospho- rus molecules will contain

a. 4 g of phosphorus.

b. 4 moles of phosphorus atoms.

c. 6.02 × 1023 phosphorus atoms.

d. 4 phosphorus atoms.

MOLES WITHIN MOLES

5. How many nitrogen atoms are there in 48.0 g of N and 48.0 g of N2?How many atoms of oxygen are there in 0.0327 moles of Na2SO4? (HINT: There are 4 moles of O in each mole of Na2SO4.) (Answer: 7.87 × 10

22 atoms)

6. How many atoms of nitrogen are there in 35.0 moles of (NH4)2SO4? (HINT: How many moles of N are in each mole of the compound?) (Answer: 4.21 × 1025 atoms)

7. How many grams of oxygen are there in 43.6 grams of CaCO3? (HINT: How many moles of O are in a mole of calcium carbonate?) (Answer:20.9 g)

PERCENT COMPOSITION BY MASS

Percent Composition by Mass

The percent composition is the percent by mass of each element in a given mass of compound.

% by mass = mass of elements in one mole of the compound

× 100% mass of one mole of the compound (molar mass)

Calculate the percent composition of hydrogen in the following compounds:

1. C2H6

2. Ca(C2H3O2)2

3. H2O

Calculate the mass of the specified element in the given mass of compound:

1. Mass of hydrogen in 350. g C2H6 (Answer: 70.5 g)

2. Mass of hydrogen in 124 g of Ca(C2H3O2)2 (Answer: 4.75 g)

3. Mass of oxygen in 100. g H2O (Answer: 88.8 g)

EMPIRICAL ANd MOLECuLAR FORMuLAS

Empirical and Molecular Formulas

REVIEW The empirical formula is the simplest whole-number ratio of elements in a compound. The molecular formula is the chemical formula of the chemical compound of interest.

1. Given the molecular formulas below, determine the empirical formula.

a. Ribose, C5H10O5, a sugar molecule in RNA. ______________________

b. Ethyl butanoate, C6H12O2, a cmpd w/ the odor of pineapple. ______________________

c. Chlorophyll, C55H72MgN4O5, part of photosynthesis. ______________________

d. DEET, C12H17ON, an insect repellent. ______________________

e. Oxalic acid H2C2O4, found in spinach and tea. ______________________

2. True or False? The empirical formula of a compound can be the same as its molecu- lar formula.

3. Determine the empirical formulas for each of the following compounds based on the information provided.

a. A sample of a compound contains 2.23 g Fe and 1.93 g S. (Answer: Fe2S3)

b. A compound is 63.11% C, 11.92% H and 24.97% F. (Answer: C4H9F)

c. A compound is 40.0% C, 6.67% H and 53.3% O. (Answer: CH2O)

Concept Question: This compound (Part c above) has a molar mass of 60 g/mol. What is its molecular formula?

REVIEW: MOLES ANd CHEMICAL COMPOSITION

Review: Moles and Chemical Composition

Use significant figures and show your work!

1. How many calcium cations are in 1.6 × 10–3 moles of Ca2+?

2. How many moles of argon atoms would be in a balloon containing 8.19 × 1024 atoms of Ar?

3. How many atoms of C are in 1.22 moles of C6H5Cl? (Answer: 4.41 × 10 24 atoms)

4. What is the mass of 1.31 × 1028 molecules of water (H2O)? (Answer: 3.92 × 10 5 g)

5. How many molecules of carbon dioxide are in 0.550 g? (Answer: 7.53 × 1021 mcs)

REVIEW: MOLES ANd CHEMICAL COMPOSITION

6. How many moles of O are in 8.0 g of N2O3? (Answer: 0.32 mol)

7. How many F atoms are there in 11.2 g of the chlorofluorocarbon, CF2Cl2? (Answer: 1.12 × 1023 atoms)

8. What is the % composition of each element in CF2Cl2?

9. A compound that contains only C, H and O is 48.38% C and 8.12% H. Determine the empirical formula. (Answer: C3H6O2)

10. Determine the empirical formula for a compound that gives the following composition: 71.65% Cl; 24.27% C; 4.07% H. The formula mass is known to be 98.96 amu. What is its molecular formula? (Answer: C2H4Cl2)

REVIEW: NOMENCLATuRE ANd CHEMICAL COMPOSITION

Review: Nomenclature and Chemical Composition

1. Identify the following as ionic or molecular compounds.

a. FeS ionic / molecular

b. SO2 ionic / molecular

c. CH3OH ionic / molecular

d. ZnSO4 onic / molecular

2. Fill in the blanks with names or formulas for the compounds.

a. CaCl2 ________________________________

b. N2H4 ________________________________

c. CuBr2 _______________________________

d. sodium phosphate ___________________

e. chromium (III) chloride ______________

f. boron trifluoride _____________________

3. Calculate the mass in grams of 5.94 × 1020 H2O2 molecules. (Answer: 0.0336 g)

4. Calculate the mass percent of each element present in aspirin, C9H8O4.

5. Vitamin C is known chemically by the name ascorbic acid. Determine the empirical for- mula of ascorbic acid if it is composed of 40.92% carbon, 4.58% hydrogen, and 54.50% oxygen. (Answer: C3H4O3)

6. What is the molecular formula of a compound given the molar mass of the compound is 30.04 grams/mole and the empirical formula is NH?

BALANCINg CHEMICAL EQuATIONS 1

Balancing Chemical Equations 1

While balancing equations, take into consideration the following:

1. Never change the formula and the subscripts in a formula.

2. Remember the seven diatomic molecules are always written as H2, N2, O2, F2, Cl2, Br2, I2. Balance the atoms on either side of the equation by changing the coefficients.

3. Make sure the coefficients are the smallest set of a whole number.

Now try balancing these equations:

1. _______ Cu + _______ O2 → _______ CuO

2. _______ S8 + _______ F2 → _______ SF6

3. _______ P4O10 + _______ H2O → _______ H3PO4

4. _______ ZnS + _______ HCl → _______ ZnCl2 + _______ H2S

5. _______ C5H12 + _______ O2 → _______ CO2 + _______ H2O

6. _______ C2H5OH + _______ O2 → _______ CO2 +_______ H2O

7. _______ NH3 + _______ O2 → _______ NO + _______ H2

8. _______ Fe2S3 + _______ O2 → _______ Fe2O3 + _______ SO2

9. _______ C6H6 + _______ O2 → _______ CO2 + _______ H2O

10. _______ NaN3 → _______ Na + _______ N2

BALANCINg CHEMICAL EQuATIONS 2

Balancing Chemical Equations 2

1. ____ N2 + ____ H2 → ____ NH3

2. ____ P2O3 → ____ P4 + ____ O2

3. ____ AgNO3 + ____ Cu → ____ Cu(NO3)2 + ____ Ag

4. ____ CF4 + ____ Br2 → ____ CBr4 + ____ F2

5. ____ CH4 + ____ O2 → ____ CO2 + ____ H2O

6. ____ MgF2 + ____ Li2CO3 → ____ MgCO3 + ____ LiF

7. ____ RbNO3 + ____ BeF2 → ____ Be(NO3)2 + ____ RbF

8. ____ HCN + ____ CuSO4 → ____ H2SO4 + ____ Cu(CN)2

9. ____ GaF3 + ____ Cs → ____ CsF + ____ Ga

10. ____ BaS + ____ PtF2 → ____ BaF2 + ____ PtS

11. ____ NaF + ____ Br2 → ____ NaBr + ____ F2

12. ____ AlBr3 + ____ K2SO4 → ____ KBr + ____ Al2(SO4)3

13. ____ Al + ____ HCl → ____ H2 + ____ AlCl3

14. ____ Pb(OH)2 + ____ HCl → ____ H2O + ____ PbCl2

15. ____ NH3 + ____ H2SO4 → ____ (NH4)2SO4

From chemfiesta.com

BALANCINg WORd EQuATIONS

Balancing Word Equations

1. Solid lithium carbonate reacts with aqueous hydrochloric acid (HCl) to produce liquid water, carbon dioxide gas, and aqueous lithium chloride.

2. Aqueous sodium sulfide will react with aqueous magnesium iodide to produce aqueous sodium iodide and a precipitate of magnesium sulfide.

3. Liquid ethanol (CH3CH2OH or C2H5OH) will burn in oxygen gas to produce liquid water and carbon dioxide gas.

4. Hydrogen gas and oxygen gas can be burned, under certain conditions, to form liquid hydrogen peroxide (officially called dihydrogen dioxide).

5. Solid calcium chloride can react with liquid water to produce hydrogen chloride gas and aqueous calcium oxide.

6. Aqueous aluminum nitrate can react with aqueous potassium fluoride to produce aque- ous potassium nitrate and a precipitate of aluminum fluoride.

7. Solid silicon and oxygen gas can react to form a solid disiliconhexaoxide.

8. Aqueous ammonium chloride will react with aqueous calcium hydroxide to form aque- ous calcium chloride and aqueous ammonium hydroxide.

9. Aqueous iron (II) bromide can react with liquid water to produce aqueous iron (II) oxide and hydrogen bromide gas.

10. Silver metal can react with hydrochloric acid (HCl) to form a precipitate of silver chlo- ride and hydrogen gas.

TuTORIAL: CLASSIFYINg TYPES OF REACTIONS

Tutorial: Classifying Types of Reactions

Synthesis/combination A + B → AB Decomposition AB → A + B Single-replacement A + BC → AC + B Double-replacement AB + CD → AD + CB Combustion CxHy + O2 → CO2 + H2O 1. Balance these reactions:

a. Na (s) + Cl2 (g) → NaCl (s)

b. HgO (s) → Hg (l) + O2 (g)

c. AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)

d. CaO (s) + CO2 (g) → CaCO3 (s)

e. HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)

f. C2H6 (g) + O2 (g) → CO2 (g) + H2O (g)

g. HCl (aq) + CaS (aq) → H2S (g) + CaCl2 (aq)

h. Ca (s) + O2 (g) → CaO (s)

i. H2O (l) → H2 (g) + O2 (g)

j. Zn (s) + CuCl2 (aq) → ZnCl2 (aq) + Cu (s)

2. Which of the equations above describe each type of reaction? List all that apply.

a. Synthesis ___________________________

b. Single-replacement _________________

c. Decomposition _____________________

d. Double-replacement ________________

e. Combustion _________________________

STOICHIOMETRY PRACTICE

Stoichiometry Practice

1. Balance the following chemical equations.

a. Zn + HCl → ZnCl2 + H2 b. CO + Fe2O3 → Fe + CO2 c. BaCl2 + Na2SO4 → BaSO4 + NaCl d. Al + CuSO4 → Cu + Al2(SO4)3 e. Which of the reactions above use BrINClHOF for diatomic molecules?

a / b / c / d

2. Classify each of the following balanced equations as combination, decomposition, sin- gle or double displacement, combustion (may be more than one type for each).

a. N2 + 3H2 → 2 NH3 ______________________________

b. BaCl2 + K2CO3 → BaCO3 + 2 KCl ______________________________

c. H2CO3 → H2O + CO2 ______________________________

d. 2 K + Br2 → 2 KBr ______________________________

e. N2 + 2 O2 → 2 NO2 ______________________________

f. Cu + 2 AgNO3 → 2 Ag + Cu(NO3)2 ______________________________

g. Which of the reactions above include BrINClHOF for diatomic molecules?

a / b / c / d / e / f

3. Complete and balance the combustion reaction for each of the following fuels:

a. Methane (natural gas, formula CH4)

b. Isooctane (major component of gasoline, formula C8H18)

c. Ethanol (alternative fuel and alcoholic beverage component, formula C2H5OH)

STOICHIOMETRY 1

Stoichiometry 1

1. Ammonia gas reacts with oxygen gas according to the following equation:

4 NH3 + 5 O2 → 4 NO + 6 H2O a. How many moles of oxygen gas are needed to react with 23 moles of ammonia?

(29 mole)

b. How many grams of NO are produced when 25 moles of oxygen gas react with an excess of ammonia? (6.0 × 102 g)

c. If 24 grams of water are produced, how many moles of nitrogen monoxide are formed? (0.89 mole)

d. How many grams of oxygen are needed to react with 6.78 grams of ammonia? (16.0 g)

100

STOICHIOMETRY 1

2. The compound calcium carbide, CaC2, is made by reacting calcium carbonate with car- bon at high temperatures. The UNBALANCED EQUATION for the reaction is:

____ CaCO3 + ____ C → ____ CaC2 + ____ CO2 a. Balance the equation.

b. How many moles of carbon are required to produce 5.0 moles CO2? (8.3 mole)

c. How many grams of calcium carbide are produced when 4.00 moles of carbon react with an excess of calcium carbonate? (102 g)

d. How many moles of carbon dioxide are produced when 55 grams of calcium car- bonate react with an excess of carbon? (0.83 mole)

e. How many grams of carbon are needed to react with 453 grams of calcium carbon- ate? (136 g)

f. How many grams of calcium carbonate are needed to form 598 grams of calcium carbide? (934 g)

101

STOICHIOMETRY CALCuLATIONS uSINg BALANCEd EQuATIONS

Stoichiometry Calculations using Balanced Equations

A TOOL FOR STOICHIOMETRY

1. Gases, using gas laws and P, V, T data of gas “A”

2. Solids or liquids

3. Volume of pure liquids (mL)

PV = nRT n = PV/RT

g of “A” g of “B”mol “A” mol “B” MMA

V “A” L

MMB

V of gas = nRT/P

L of solution “B”

Molarity of solution “B”

Volume of solution “B”

Molarity of “B”mol “A”/L

or molarity of “A” solution

Molar ratio from balance equationdensity

4. Solutions volume “A” mL

ABBREVIATIONS uSEd

MM = Molar mass

V = volume

P = pressure

T = temperature

n = moles

NOTE: Gases may not be covered in CHEM& 140 but will be covered in CHEM& 162. Molarity will be covered later in this course.

102

STOICHIOMETRY: IRON ORE SMELTINg

Stoichiometry: Iron Ore Smelting

Iron ore is converted to iron by heating it with coal and oxygen in a process called smelting:

_____ Fe2O3 (s) + _____ C (s) + _____ O2 (g) → _____ Fe (s) + _____ CO2 (g) unbalanced equation!

1. Calculate the number of moles of Fe2O3 that would be needed to react exactly with 0.49 mol of O2. (Answer: 0.98 mol)

2. Calculate the number of moles of C that would need to react to produce 6429 g of Fe. (Answer: 115.1 mol)

3. How many grams of CO2 would be produced from smelting 2385 g of Fe2O3? (Answer: 1315 g)

4. How many kilograms of Fe can be smelted from 7.545 kg of Fe2O3? (Answer: 5.227 kg)

103

STOICHIOMETRY: REMOVINg CO2 IN SPACE

Stoichiometry: Removing CO2 in Space

The Contaminate Control Cartridge (CCC), which contains lithium hydroxide (LiOH), is used to remove carbon dioxide (CO2) from the space shuttle cabin.

The removal of CO2 is represented by the following equation:

2 LiOH (s) + CO2 (g) → Li2CO3 (s) + H2O (g) A typical space shuttle crew consists of six individuals and each CCC contains 750 g of LiOH. Assuming that each crew member expels 42.0 g of CO2 per hour on average, and that a mission is scheduled to last 18 days, how many CCCs must be carried on board the space shuttle?

PROBLEM ANd SOLuTION KEY (ONE APPROACH) The Contaminate Control Cartridge (CCC), which contains lithium hydroxide (LiOH), is used to remove carbon dioxide (CO2) from the space shuttle cabin. The removal of CO2 is repre- sented by the following equation

2 LiOH (s) + CO2 (g) → Li2CO3 (s) + H2O (g) 1. A typical space shuttle crew consists of six individuals and each CCC contains 750 g of

LiOH. Assuming that each crew member expels 42.0 g of CO2 per hour on average, and that a mission is scheduled to last 18 days, how many CCCs must be carried on board the space shuttle?

18 days · 24 hr

· 42.0 g

= 18,100 g per person day 1 hr

For a crew of six:

6(18,100 g) = 109,000 g CO2 per person

108,000 g CO2 · 1 mol CO²

· 2 mol LiOH

· 23.9 g LiOH

· 1 CCC

= 158 CCCs 44.0 g CO2 1 mol CO2 1 mol LiOH 750 g LiOH

From Nasa.gov “Carbon dioxide removal – Stoichiometry”

104

LIMITINg REACTANT BY MOLES

Limiting Reactant by Moles

1. A recipe for cookies states that 3 cups of flour, 2 eggs, and 1.5 tsp of baking soda can make 24 cookies. Let’s say that you have 5 cups of flour, 3 eggs, and 4 tsp of baking soda.

a. How many cookies can be made (what is the theoretical yield)?

b. What is the limiting reactant? How can you tell?

2. Ammonia gas reacts with oxygen gas according to the following equation:

4 NH3 + 5 O2 → 4 NO + 6 H2O a. Calculate the number of moles of oxygen gas needed to react with 7.00 moles of NH3.

b. If you had 7.00 moles of NH3 and 6.00 moles of oxygen gas, what would be the lim- iting reactant?

c. If you had 7.00 moles of NH3 and 10.0 moles of oxygen gas, what would be the lim- iting reactant?

105

LIMITINg REACTANT BY MOLES

3. Aluminum reacts with aqueous sulfuric acid to produce an aluminum sulfate aqueous solution and hydrogen gas.

2 Al (s) + 3 H2SO4 (aq) → Al2(SO4)3 (aq) + 3 H2 (g) a. What would be the limiting reactant if you had 2 moles of Al and 2 moles of sulfu-

ric acid?

b. What would be the limiting reactant if you had 3 moles of Al and 3 moles of sulfu- ric acid?

c. What would be the limiting reactant if you had 5 moles of Al and 7 moles of sulfu- ric acid?

d. What would be the limiting reactant if you had 0.50 moles of Al and 0.75 moles of sulfuric acid?

4. Ammonia can be synthesized by the following reaction.

2 NO (g) + 5 H2 (g) → 2 NH3 (g) + 2 H2O (g) a. If 10.0 moles of NO and 20.0 moles of hydrogen react, what is the theoretical yield

(in moles) of ammonia? (Answer: 8 mol)

b. If 0.500 moles of NO and 1.50 moles of hydrogen react, what is the theoretical yield (in moles) of ammonia? (Answer: 0.5 mol)

106

ACTIVITY: LIMITINg REACTANT

Activity: Limiting Reactant

PART 1: REVIEW STOICHIOMETRY There are three steps to doing gram-to-gram conversions in stoichiometry:

1. Convert grams of the starting substance to moles. (Use molar mass!)

2. Convert the moles of the starting substance to moles of the ending substance. (Use the mole ratio given in the balanced chemical equation.)

3. Convert the moles of the ending substance to grams. (Use molar mass!)

uSE SIgNIFICANT FIguRES! uSE uNITS! SHOW YOuR WORK!

REVIEW

1. Balance this chemical equation of the combustion of methane:

CH4 + O2 → CO2 + H2O

2. What is the theoretical yield or the maximum mass of water (in grams) that may be formed from the reaction of 5.00 grams of methane, CH4? Follow the three steps above and show your work below. (Answer: 11.2 g)

(NOTE: In this problem, oxygen is “in excess.” When there is an excess of one reac- tant that means you can assume all of the other reactant gets converted to product. (In other words, you don’t have to worry about it in the calculation.) We can say “methane reacts completely” or “methane is the limiting reactant” since the calculation is based on this assumption.)

Theoretical yield of water: _________ grams

3. Now suppose there is an excess of methane and only 5.00 grams of O2. (Oxygen is now the limiting reactant.) What is the theoretical yield of water in this case? Follow the three steps above and show your work below. (Answer: 2.82 g)

Theoretical yield of water: _________ grams

107

ACTIVITY: LIMITINg REACTANT

PART 2: THE LIMITINg REACTANT The Limiting Reactant is the reactant that determines the theoretical yield. The Limiting Reactant is completely consumed (or “used up”) while some of the other reactant (in excess) will be “left-over” at the end of the reaction.

1. Suppose you have a reaction with 5.00 grams of CH4 ANd 5.00 grams of O2 in one container. Unlike in the Review Questions #2 and #3, you are not told which reactant is limiting and which is in excess so you are not sure which to use in the calculation, 5.00 grams of methane OR 5.00 grams of oxygen.

In this situation, YOu have to determine which is the limiting reactant. You do this by setting up two stoichiometry problems—one for each reactant. Luckily, you already did this on the previous page. Let’s compare the amount of water formed from 5.00 grams of each reactant:

Amount (grams) of water formed from reacting 5.00 grams of methane: ________ grams

Amount (grams) of water formed from reacting 5.00 grams of oxygen: _________ grams

2. Which one “limits” the amount of water formed, the 5.00 grams of CH4 or 5.00 grams of O2? (Circle one.) The one you circle is called “the limiting reactant”.

How many grams of water is formed from this limiting reactant? _________ grams

How many grams of the limiting reactant should remain after reaction? _________ grams

3. To summarize, fill in the blanks:

When 5.00 grams of CH4 and 5.00 grams of O2 react, _________ grams of water can be formed.

In this case, _____________ is the limiting reactant and ____________ is the excess reactant.

4. Which substances should be present after the reaction? Circle all that apply:

CH4 / O2 / H2O / CO2

The theoretical yield is based on calculations assuming all of the limiting reactant is converted to product. In reality, this may not be the case. The amount of product actu- ally formed in a reaction is called the “actual yield” and is usually less than the amount expected in theory.

% yield actual yield

× 100 theoretical yield

5. If this reaction of methane and oxygen had a 75% yield, what is the actual mass of water produced?

108

THEORETICAL YIELd ANd LIMITINg REAgENTS

Theoretical Yield and Limiting Reagents

1. For the reaction 3 H2 (g) + N2 (g) → 2 NH3 (g), 3 mol H2 is reacted with 6 mol N2. a. __________ mol of NH3 is produced

b. __________ mol of H2 remains

c. __________ mol N2 remains

2. For the reaction 2 N2H4 (l ) + N2O4 (l ) → 3 N2 (g) + 4 H2O (l ), 160 g N2H4 is mixed with 160 g N2O4. (Answer: 125 g H2O)

a. __________ is the limiting reagent

b. __________ g H2O is produced

3. For the reaction Fe2O3 (s) + 3 CO (g) → 2 Fe (g) + 3 CO2, 224 g of CO is available to react with 400. g Fe2O3. (Answer: 278 g Fe, 331 g CO2)

a. __________ is the limiting reagent

b. __________ g iron is produced

c. __________ g CO2 is produced

4. For the reaction 2 C4H10 (g) + 13 O2 (g) → 8 CO2 (g) + 10 H2O (l ), 300 g of C4H10 is com- busted in 1000 g of O2. (Answer: 400 g H2O)

a. __________ is the limiting reagent

b. __________ g H2O is formed

From ThoughtCo. Chemistry www.thoughtco.com

109

PERCENT YIELd

Percent Yield

1. For the reaction of nitrogen and hydrogen to produce ammonia:

3 H2 (g) + N2 (g) → 2 NH3 (g) a. If 10.0 g of N2 were reacted, what is the theoretical yield of ammonia, in grams?

(Answer: 12.2 g)

b. If 10.0 g of N2 were reacted, and 9.1 g of ammonia were obtained, what is the % yield? (Answer: 74.6%)

2. Potassium superoxide, KO2, reacts with carbon dioxide to form potassium carbon- ate and oxygen. This reaction makes potassium superoxide useful in a self-contained breathing apparatus. Write a balanced equation before doing calculations.

a. How much oxygen could be produced from 2.50 g of KO2 and 4.50 g of carbon dioxide? (Answer: 0.845 g)

b. If only 0.500 g of oxygen is produced from 2.50 g of KO2 and 4.50 g of carbon dioxide, what is the % yield? (Answer: 59.2%)

3. After 3.8 mL of liquid ethanol (C2H5OH, density = 0.789 g/mL) is allowed to burn in the presence of 12.5 g of oxygen, 3.10 mL of water (density = 1.00 g/mL) is collected. Determine the limiting reactant, theoretical yield of H2O and percent yield for the reac- tion. Write the balanced chemical equation for the combustion of ethanol before doing calculations. (Answer: 3.52 g H2O, 88.1%)

110

REVIEW: STOICHIOMETRY

Review: Stoichiometry

1. For this generic chemical reaction: 2 A + 3 B → 3 C

How many moles of B are required to react with: (a) 6 mol A, (b) 2 mol A, (c) 7 mol A

2. a. Balance this equation: Ca + O2 → CaO

b. How many moles of product forms from 0.112 mol of oxygen and excess calcium?

3. a. Write a balanced chemical equation for the combustion of C3H8.

b. How many moles of each product forms from 4.6 mol of C3H8 and excess O2?

c. How many moles of oxygen is required?

4. For the decomposition of potassium chlorate (KClO3) to form potassium chloride and oxygen gas:

a. Write a balanced chemical equation

b. What is the theoretical yield of oxygen if 2.00 grams of potassium chlorate are used? (Answer: 0.783 g)

5. In this reaction: 2 S + 3 O2 → 2 SO3)

If 5 mol of S and 9 mol of O2 are used, how many mole of SO3 can be produced? Which is the limiting reactant, S or O2?

111

REVIEW: STOICHIOMETRY

6. What is the limiting reactant and theoretical yield if 0.483 mol of Ti and 0.911 mol Cl2 are used in this reaction: Ti + 2 Cl2 → TiCl4

7. In the reaction: 4 HCl + O2 → 2 H2O + 2 Cl2, when 63.1 g of HCl react with 17.2 g of O2, 49.3 g of Cl2 are collected. Determine the limiting reactant, theoretical yield of Cl2 and percent yield of the reaction. (Answer: 61.4 g, 80.3%)

8. If the theoretical yield of a reaction is 0.118 g and the actual yield is 0.104 g what is the percent yield?

112

MOLARITY 1

Molarity 1

There are several ways of expressing solution concentration. The most common is molar- ity, M (note that is a capital M). The unit M is defined as:

moles solute per liter of solution or M = mols solute/L solution

Note that the denominator is liters of solution, not liters of solvent. So this volume is the final volume of the solution, after the solute has been added.

Concentrations can be used like conversion factors:

Molarity = mol solute

L of solution

■ allows conversion of moles to liters or liters to moles

Mass percent = g of solute × 100

g of solution

■ allows conversion of g solution to g of solute or g solute to g solution

PRACTICE

1. Calculate the molarity of each solution

a. 33.2 g of KCl in 0.895 L of solution (Answer: 0.498 M)

b. 61.3 g of C2H6O in 3.4 L of solution (Answer: 0.39 M)

c. 38.2 mg of KI in 112 mL of solution (Answer: 2.05 × 10–3 M)

2. Determine how many grams of sucrose (C12H22O11) are contained in 1.72 L of 0.758 M sucrose solution. (Answer: 445 g)

3. The concentratoin of hydrochloric acid in the stomach is 0.100 M. What volume of stom- ach acid contains 0.500 g of hydrochloric acid? (Answer: 137 mL)

113

MOLARITY 2

Molarity 2

Answers are given in parenthesis.

1. A 1.50 mole sample NaOH was dissolved in water to make a volume of 2.0 L. What is the molarity of this solution? (0.75 M)

2. A 0.300 mole sample NaOH was dissolved in water to make a volume of 150 mL. What is the molarity of this solution? (2.0 M)

3. A 2.00 g sample of NaOH was dissolved in water to make a volume of 100. mL. What is the molarity of this solution? (0.500 M)

4. How many milliliters of 0.250 M NaOH are needed to provide 0.200 moles of NaOH? (800. mL)

5. How many grams of NaOH are there in 80.0 mL of 0.400 M NaOH solution? (1.28 g)

6. How many moles of NaOH are contained in 150 mL of 3.00 M NaOH solution? (0.45 mol)

7. What is the molarity of a solution when 0.135 moles of LiOH are dissolved in water to give 50.0 mL of solution? (2.70 M)

8. How many grams of KMnO4 would you need to prepare 520 mL of 0.58 M solution? (48 g)

9. How many milliliters of 5.3 M KBr do you need to get 0.112 moles of KBr? (21 mL)

114

MOLARITY: APPLICATIONS

Molarity: Applications

1. Consider a solution of MgCl2 in water.

a. What is the solvent?

b. What is the solute?

c. Is the solution a homogeneous or a heterogeneous mixture?

d. What molecules, atoms and/or ions are present in the solution? Write their formulas.

e. If the solution had a concentration of 1.0 M and a volume of 1.0 L, what are the spe- cific amounts of the molecules, atoms, and/or ions present in the solution?

2. A sports beverage has a concentration of sodium phosphate of 0.0020 M.

a. How many moles of sodium ions are there in 1.00 moles of Na3PO4?

b. What is the molarity of sodium in this beverage?

3. A 75 mL of 0.211 M NaOH is diluted to a final volume of 125 mL.

a. How many moles of NaOH are there?

b. What is the final molarity of NaOH? (Answer: 0.127 M)

4. An aqueous glucose solution has a volume of 0.800 L and a concentration of 1.2 M.

a. How many moles of glucose are present in the solution? (Answer: 0.96 mol)

b. What volume of water should you add to dilute the concentration to 1.0 M? (Answer: 0.96 L)

115

MOLARITY: FILL-IN-THE-BLANKS TABLE

Molarity: Fill-in-the-Blanks Table

Complete the table for the following aqueous solutions.

Solute Mass Solute (g) Volume (liters) Molarity (mol/L or M)

LiBr 43.3 0.5

KCl 37.5 0.125

CuSO4 1.6 0.125

NaCl 146.0 25.0

Al(NO3)3 0.5 0.6

BaCl2 385.5 2.5

KOH 28.0 0.25

(NH4)2SO4 0.3 2.0

NaOH 80.0 1.6

NH4Cl 160.2 4.0

From ThoughtCo. Chemistry www.thoughtco.com

Answers (from top to bottom): 1 M, 4.02 L, 32 g, 0.999 M, 60 g, 0.74 L, 2.0 M, 80 g, 1.3 L, 0.75 M.

116

CONCENTRATION: MASS PERCENT (% M/M)

Concentration: Mass Percent (% m/m)

Concentration: A measure of the amount of solute in a given amount of solution. There are many different ways of expressing concentration, some are more appropriate than oth- ers depending on the purpose. Percent concentrations are frequently used with commer- cial products.

Important Concepts to Remember

■ Solution = solute + solvent ■ % is part × 100

whole

■ Masses are additive, volumes are not

Construct an equation for % (m /m ):

PRACTICE

1. What is the percent-by-mass, % (m/m), concentration of sucrose in a solution made by dissolving 7.6 g of sucrose in 83.4 g of water?

2. If 162.35 g aluminum hydroxide are dissolved in 6750 g of solution, what is the % (m/m) concentration of the solution?

3. Calculate the concentration in % (m/m) of a 450.0 g solution containing 0.0762 moles of iodine (I2). (Answer: 4.30%)

117

PERCENT CONCENTRATION PRACTICE

Percent Concentration Practice

1. Calculate the mass % for 1.00 mg of Nutrasweet added to 50.0 g of water.

2. Calculate the mass % for 6.6 g of sodium nitrate added to 6.0 L of water. (Assume the density of water is 1.0 g/mL.) (Answer: 0.11%)

3. Calculate the grams of glucose present in 400.0 g of a solution that is 8.0% (m/m) glu- cose. (Answer: 32 g)

4. How many grams of a 0.600% (m/m) penicillin solution should be given to patient who requires 1.0 g of penicillin? (Answer: 170 g)

5. A glass of milk is about 8.00 fluid ounces. How many grams of fat is in this glass of milk, assuming the milk is 2.0% (m/m) fat. (Answer: 4.5 g)

118

SOLuTION STOICHIOMETRY 1

Solution Stoichiometry 1

1. Determine the volume in liters of 0.150 M NaOH solution required to neutralize 55 mL of a 0.055 M hydrochloric acid solution. The neutralization reaction is: NaOH (aq) + HCl (aq) → H2O (l ) + NaCl (aq) (Answer: 0.020 L)

2. For this reaction: 2 HNO3 (aq) + Na2CO3 (aq) → H2O (l ) + CO2 (g) + 2 NaNO3 (aq) a. How many milliliters of 0.112 M Na2CO3 are needed to completely react with 27.2 mL

of 0.135 M HNO3? (Answer: 16.4 mL)

b. A 25.0-mL sample of HNO3 solution requires 35.7 mL of 0.108 M Na2CO3 to com- pletely react with all of the HNO3 solution. What was the concentration of the HNO3 solution? (Answer: 0.308 M)

3. What volume in milliliters of 0.225 M K3PO4 solution is needed to completely react with 134 mL of 0.0112 M NiCl2? (Answer: 4.45 mL)

2 K3PO4 (aq) + 3 NiCl2 → Ni3(PO4)2 (s) + 6 KCl (aq)

119

SOLuTION STOICHIOMETRY 1

4. A 10.0-mL sample of an unknown H3PO4 solution requires 112 mL of 0.100 M KOH to completely react with the H3PO4 according to the following reaction. What was the concentration of the unknown H3PO4 solution? (Answer: 0.373 M)

H3PO4 (aq) + 3 KOH (aq) → 3 H2O (l ) + K3PO4 (aq)

5. What is the minimum amount in liters of 6.0 M H2SO4 necessary to produce 15.0 g of hydrogen gas according to the following reaction? (Answer: 1.2 L)

2 Al (s) + 3 H2SO4 (aq) → Al2(SO4)3 (aq) + 3 H2 (g)

Answers: (1) 0.020 L (2a) 16.4 mL (2b) 0.308 M (3) 4.45 mL (4) 0.373 M (5) 1.2 L

120

SOLuTION STOICHIOMETRY 2

Solution Stoichiometry 2

Solve the following solutions stoichiometry problems. Write a balanced chemical equation for each and use dimensional analysis and significant figures.

M = mol solute/Liter solution

1. What volume of 0.496 M HCl is required to neutralize 20.0 mL of 0.809 M sodium hydrox- ide? (0.0326L HCl)

2. How many mL of 0.715 M HCl is required to neutralize 1.25 grams of sodium carbonate? (Producing carbonic acid, H2CO3) (0.0330 L HCl)

3. A 20.00-mL sample of H2SO4 requires 25.00 mL of 0.500 M NaOH to completely neutral- ize it. What is the molarity of H2SO4? (Answer: 0.313 M H2SO4)

4. A sample of NaHCO3 weighing 0.1015 g requires 47.21 mL of HCl(aq) to completely neu- tralize it. What is the molarity of HCl? (Answer: 0.02559 M)

121

ACTIVITY: ELECTROMAgNETIC RAdIATION

Activity: Electromagnetic Radiation

ELECTROMAgNETIC RAdIATION (LIgHT WAVES)

Wave 1

Wave 2

1. Frequency is the number of waves (cycles) which travel in a given amount of time. Which wave has a higher frequency, wave 1 or wave 2?

2. Wavelength is the distance between two crests (maxima) or two troughs (minima)— also known as the distance of a repeating unit. Mark this distance for each wave on the picture above. Which wave has a longer wavelength, wave 1 or wave 2?

3. The higher the frequency, the higher the energy of the wave. Which wave represents higher energy radiation, wave 1 or wave 2?

4. True or false? Higher energy light has longer wavelengths than lower energy light.

Radiation Type Wavelength (m)

Radio

103 Microwave

10–2 Infrared

10–5 Visible

0.5 × 10–6 Ultraviolet

10–8 X-Ray

10–10 Gamma Ray

10–12

Buildings Humans Butter�ies Needle Point Protozoans Molecules Atoms Atomic Nuclei

Approximate Scale of Wavelength

Frequency (Hz)

Figure 1. The Electromagnetic Radiation Spectrum

5. List these types of electromagnetic radiation in order of low energy to high energy. (Choices: gamma rays, radio waves, microwaves, ultraviolet, visible light, infrared, x-rays)

Low energy __________________ __________________ __________________

_____________ _____________ _____________ _____________ High energy

122

ACTIVITY: ELECTROMAgNETIC RAdIATION

6. The colors of visible light can be remembered by this mnemonic: ROY G. BIV. What are the colors represented by the letters of this mnemonic?

R_______ O_______ Y_______ G_______ B_______ I_______ V_______

7. If red light has a wavelength of 700 nm (nanometers) and violet light has a wavelength of 400 nm, which color of light is higher in energy? __________

Summary (circle the best answer choice): The shorter the wavelength, the lower / higher its frequency and the lower / higher its energy. In other words, wavelength and frequency are directly / inversely proportional, wavelength and energy are directly / inversely proportional, and the frequency and energy of a wave are directly / inversely proportional.

123

ENERgY, WAVELENgTH, ANd FREQuENCY

Energy, Wavelength, and Frequency

1. Sketch a diagram of a wave and label a crest, a trough, a wavelength, and amplitude.

2. Define frequency of a wave.

3. What are the units and symbol/abbreviations for each of these wave properties?

a. energy b. wavelength c. frequency

4. a. If the frequency (ν) of a wave is increased, what happens to the energy (E ) of the wave?

increases or decreases

b. Based on your answer, which relationship is correct? (∝ means “proportional to”) E ∝ ν or E ∝ (1/ν)

5. a. If the frequency (ν) of a wave is increased, what happens to the wavelength (λ) of the wave?

increases or decreases

b. Based on your answer, which relationship is correct? (c = speed of light)

ν ∝ λ or ν ∝ (1/λ) 6. These variables can be written into an equation but require constants in order for the

mathematics to work. The constants used are Planck’s constant (h = 6.626 × 10–34 js) and the speed of light (c = 2.998 × 108 m/s).

E = hν and c = λν

a. Verify that the first equation is consistent with your answer to 4b above. Verify that the units of Planck’s constant (h) must be Js.

b. Verify that the second equation is consistent with your answer to 5b above. Verify that the units of the speed of light (c) must be m/s.

7. Use the two equations above to derive a third equation, E = hc/λ. This is also a useful equation to use.

124

ENERgY, WAVELENgTH, ANd FREQuENCY

EXERCISES 1. The yellow light given off by a sodium vapor lamp used for public lighting has a wave-

length of 589 nm. What is the frequency of this radiation? (Answer: 5.09 × 1014 Hz or 1/s)

2. A radio station broadcasts at a frequency of 590. KHz. What is the wavelength of the radio waves? (Answer: 508 M)

3. A certain microwave has a wavelength of 0.032 meters. Calculate the frequency of this microwave. (Answer: 9.4 × 109 1/s

4. Microwave ovens emit microwave energy with a wavelength of 12.9 cm. What is the energy of exactly one photon of this microwave radiation? (Answer: 1.54 × 10–24 J)

5. What is the wavelength of light of a photon with an energy of 4.90 × 10–15 J. (Answer: 4.05 × 10–11 m)

6. Calculate the frequency of a gamma ray photon whose energy is 3.33 × 10–13 J. (Answer: 5.02 × 1020 1/s)

7. Calculate the energy of a photon with frequency of 8.5 × 1014 Hz. (Answer: 5.6 × 10–19 J)

8. Calculate the energy of a photon of radiation with a wavelength of 6.4 × 10–7 m. (Answer: 3.1 × 10–19 J)

125

ACTIVITY: THE BOHR MOdEL

Activity: The Bohr Model

Use a periodic table for this activity.

PART 1: BACKgROuNd INFORMATION Ernest Rutherford showed with the gold Foil Experiment that atoms contained a nucleus of positive charge. It was assumed the electrons “floated around” outside the nucleus. What prevented these negatively charged electrons from being sucked into the positively charged nucleus?

Danish physicist Niels Bohr proposed that electrons traveled on orbits (rings) that prevented them from spiraling into the nucleus. The number of electrons in each orbit can be deter- mined from the periodic table, as will be discussed below. These orbits are QuANTIZEd. This means electrons occupy orbits, and do not exist in between orbits.

Each PERIOd (row) of the periodic table shows a layer of an atom. The atom can be thought of as an onion, with layers of orbits (or shells). For example, hydrogen is in period 1, which represents the first orbit of the atom (closest to the nucleus, n = 1). Helium is also in period 1. n is the orbit number (also called an “energy level”).

The Bohr diagrams for each element are shown below:

Bohr Diagram for Hydrogen

Nucleus

Electron

n = 1

Bohr Diagram for Helium

n = 1

1. Review: Using a periodic table, how do you know how many electrons are in an atom? For example, how do you determine that helium contains two electrons, as shown above?

Elements in PERIOD two will have two layers of an atom, the first energy level being n = 1, and the second energy level being n = 2. As you can see above, two electrons fit in the first energy level, then the second level begins to fill.

126

ACTIVITY: THE BOHR MOdEL

2. Fill in the Bohr diagrams for lithium, nitrogen, and neon below by drawing in the elec- trons. HINT: (1) Determine the number of electrons each element has, then (2) fill in the first orbit, then (3) fill in the second orbit. Use the periodic table to guide you on when to switch from the first to the second “layer.” NOTE: It only matters how many electrons are in each orbit, not how they are placed within an orbit. Generally we pair the elec- trons, as shown with helium above.

n = 1

n = 2

n = 1

n = 2

n = 1

n = 2

3. Can you guess how to draw a Bohr diagram for phosphorus?

draw it in the box provided to the right.

*Don’t forget to draw the nucleus, the orbits, label the orbits (n = 1, etc.) and draw in the electrons.*

use the periodic table to guide you!

4. The electrons in the outermost energy level play a special role in chemistry. They are called valence electrons.

a. How many valence electrons are in Li, N, and Ne based on your drawings above (Q9)?

b. What group (column) are Li, N, and Ne in?

c. Is there a connection between the number of valence electrons and the group number?

d. How many valence electrons are in P, drawn in Q10?

e. Why do you think it’s similar to N? (HINT: Location in periodic table)

f. Without drawing them out, how many valence electrons do you predict are in a sodium (Na) and an argon (Ar) atom?

127

ACTIVITY: THE BOHR MOdEL

PART 2: HOW dO ATOMS EMIT (RELEASE) LIgHT? Below is a diagram of hydrogen. Notice that even though hydrogen is in the first period, there are many energy levels in the atom. Theoretically, every atom contains an infinite number of energy levels. Because hydrogen only has one electron, all the other energy lev- els are empty. It doesn’t mean they can’t be used later on. Also note that the energy levels get closer to each other the farther they are from the nucleus. Just to keep it simple, let’s only consider the first four energy levels.

When hydrogen’s electron is in the lowest possible energy level (closest to the nucleus), it is said that hydrogen is in its “ground state” configuration. If energy were supplied to this atom (in the form of heat or light), the electron can move to higher energy levels. This results in an “exited state” configuration.

Ground State Excited State

n = 4 n = 3 n = 2

n = 1

n = 4 n = 3 n = 2

n = 1

Absorbs Energy

This is called an electron transition. If given enough energy, the electron can undergo an n =1 → n = 4 transition, as shown above. 5. What are other possible transitions if the electron absorbs less energy than shown

above?

n = 1 → _____ n = 1 → _____ 6. Keeping in mind that the electron absorbed energy to “jump” from n = 1 to n = 4, do you

think the electron will absorb or release energy when dropping back down from n = 4 to n = 1? Explain your choice.

7. Only considering the four energy levels shown, what are the possible electron transi- tions as the atom releases energy?

n = 4 → _____ n = 4 → _____ n = 4 → _____ 8. Of the transitions you listed above, which one releases the MOST energy? Explain.

128

ACTIVITY: THE BOHR MOdEL

9. Each transition releases energy in the form of light. If three colors are emitted (red, green, and violet), which transition would go with which color? List them below.

10. Only considering the four energy levels shown, can you think of any other possible tran- sitions as the atom releases energy? (HINT: Does the electron have to start in n = 4?) List all the ones your group can think of.

EXERCISES

1. List the following types of electromagnetic radiation in order of increasing energy: gamma rays, radio waves, microwaves, visible light, x-rays, ultraviolet light, infrared radiation.

2. Which color of light comes from a greater energy transition, red or blue?

3. Is there a relationship between the Group number of an element and the number of valence electrons in an element? What is this relationship? Write it down:

4. Based on magnesium being in Group 2A, how many valence electrons do you predict magnesium has? Using the periodic table to guide you, draw the Bohr diagram for a magnesium atom and see if you’re correct.

5. An electron drops from n = 4 → 3, and then n = 3 → 1. Two frequencies of light are emitted. How does their combined energy compare with the energy of the single fre- quency that would be emitted if the electron dropped n = 4 → 1?

6. What is the meaning of QUANTIZED energy?

7. What might the spectrum of an atom look like if the atom’s electrons were not restricted to particular energy levels? (if the energy were not QUANTIZED)

129

ATOMIC ORBITAL ENERgY LEVELS

Atomic Orbital Energy Levels

*Energy levels are not drawn to scale.

7p 7p

6p 6p 6p

5p 5p 5p

4p 4p 4p

3p 3p 3p

2p 2p 2p

6d 6d 6d 6d 6d

5d 5d 5d 5d 5d

4d 4d 4d 4d 4d

3d 3d 3d 3d 3d

5f 5f 5f 5f 5f 5f 5f

4f 4f 4f 4f 4f 4f 4f

En er

130

ELECTRON CONFIguRATIONS 1

Electron Configurations 1

There are three common ways to represent an electron configuration.

1. The Orbital (Box) diagram

These show each orbital as a box (like the diagram on page 40). The boxes are filled with arrows (representing electrons). The electrons are paired with opposite spin (one arrow pointing up, one arrow pointing down).

NOTE: Hund’s rule says that when orbitals within the same subshell (like 2p, 3d, etc.) are available, you first fill each orbital with ONE electron first.

EXAMPLE

Carbon

1s 2s 2p

2. The Full Electron Configuration (a short cut of the orbital diagram)

This is like the exercise on the previous page.

EXAMPLE

Sodium’s full electron configuration is 1s2 2s2 2p63s1

3. The Noble gas Shorthand Configuration (a short cut of the full configuration)

Look at the Noble gas (element in group 8A) that comes BEFORE the element of interest.

EXAMPLE

The Noble gas before Na is Ne.

The Noble gas shortcut for sodium would be [Ne] 3s1

(The 1s2 2s2 2p6 is substituted for [Ne].) It’s very convenient!

131

ELECTRON CONFIguRATIONS 1

EXERCISES 1. Practice writing all three kinds of electron configurations (shown in #1–4 above) for

these elements: Al, Xe, Ba, Pb (the last three are long, but it’s an exercise!)

2. In the full electron configuration (#1) for all four elements done in Exercise 1, circle the all subshells (ex. 3s2) in the highest energy level (1, 2, 3, etc.). These are called valence electrons (they are those electrons in the outermost principal energy level) and are the electrons responsible for chemical reactions.

3. Make a table of each of the four elements (in Part 1), the number of valence electrons, and the group number of each element (Usually given as a roman numeral in that col- umn of the periodic table). Do you see a relationship between the number of valence electrons and the group #? What is the relationship?

132

ELECTRON CONFIguRATIONS 2

Electron Configurations 2

1. Write the unabbreviated electron configurations of the following elements:

a. sodium ___________________________________________________________

b. iron ______________________________________________________________

c. bromine ___________________________________________________________

d. barium ____________________________________________________________

e. neptunium _________________________________________________________

2. Write the abbreviated electron configurations of the following elements:

a. cobalt ____________________________________________________________

b. silver _____________________________________________________________

c. tellurium __________________________________________________________

d. radium ____________________________________________________________

e. lawrencium ________________________________________________________

3. Determine what elements are denoted by the following electron configurations:

a. 1s22s22p63s23p4 ____________________________________________________

b. 1s22s22p63s23p64s23d104p65s1 _________________________________________

c. [Kr] 5s24d105p3 _____________________________________________________

d. [Xe] 6s24f145d6 _____________________________________________________

e. [Rn] 7s25f11 ________________________________________________________

4. Determine which of the following electron configurations are not valid:

a. 1s22s22p63s23p64s2 4d10 4p5 _________________________________________

b. 1s22s22p63s3 3d5 __________________________________________________

c. [Rn] 7s25f8 _________________________________________________________

d. [Kr] 5s24d105p5 _____________________________________________________

e. [Xe] ______________________________________________________________

From chemfiesta.com.

133

ELECTRON CONFIguRATIONS 3

Electron Configurations 3

The periodic table mimics the electronic structure of the atom. Electrons fill sub-levels from lowest to highest energy. The row of the periodic table that an element is located in corresponds to the number of occupied energy levels in an atom of that element.

The number of electronic sub-levels possible is the same as the electronic energy level number. So, for example the first energy level, n = 1, has only the “s” sub-level, n = 2 has the s and p sublevels, n = 3 has s, p, and d sublevels, n = 4 has s, p, d, and f sublevels.

1. Write electron configurations for the following elements

a. O: _______________________________________________________________

b. P: _______________________________________________________________

c. Br: _______________________________________________________________

d. Mn: ______________________________________________________________

e. Pb: _______________________________________________________________

2. Write electron configuration for each of the following ions:

a. Br – ______________________________________________________________

b. K+ _______________________________________________________________

c. Ca2+ _____________________________________________________________

d. P3– ______________________________________________________________

e. Al3+ ______________________________________________________________

134

ORBITAL dIAgRAMS

Orbital diagrams

1. Aufbau principle: Electrons will fill the lowest energy orbitals first.

2. Pauli Exclusion Principle: Orbitals may hold no more than 2 electrons with opposing spins – symbolized by up and down arrows.

3. Hund’s rule: When occupying degenerate orbitals electrons will distribute to mini- mize electron-electron repulsion. For example, nitrogen has 5 valence electrons, the 3 electrons in the 2p sub-level will each occupy a different 2p orbital so that each 2p orbital has one electron with parallel spin. This minimizes electron-electron repulsion. Degenerate means “having equal energy.”

1s 2s 2p 2p 2p

135

REVIEW: ELECTRON CONFIguRATIONS

Review: Electron Configurations

1. How many electrons can a single orbital hold? __________

2. How many total electrons can fit in the 2p subshell of an atom? __________

3. What is wrong with the electron configuration 1s21p62s22p1?

4. Draw an orbital box diagram for Be.

5. Write the electron configuration for O.

6. To the right, draw an orbital energy diagram for Al.

7. Write the electron configuration for Ga.

8. Write the condensed electron configuration for K.

9. Write the condensed electron configuration for Cl.

The electrons in the outermost shell (n) of an atom are called the valence electrons. Electrons that are not valence electrons are called core electrons.

10. How many valence electrons are in a Cl atom? _____ How many core electrons?_____

11. How many valence electrons are in a K atom? _____ How many core electrons?_____

12. How many valence electrons are in F? How about Br? Compare your answers for F and Br to Cl (see above). What is the trend?

Argon, Ar, is in Group 8A—the noble gases. Argon atoms are very stable—they do not like to bond together to form molecules or make compounds with atoms of other elements.

13. How many valence electrons are in an Ar atom? __________

14. What would a Cl atom need to do to have the same number of valence electrons as Ar? What would a K atom need to do to have the same number of valence electrons as Ar?

136

ELECTRON CONFIguRATIONS: PuZZLE

Electron Configurations: Puzzle

A puzzle for you …

Instructions: Read each statement and identify the element described. Write the element symbol in the first column, then identify the first letter of that element symbol and place it in the second column.

description Element Symbol First Letter of Element Symbol

This is one of only two letters that does not appear anywhere on the periodic chart. It is alphabetically the second of the two.

The third element in the second row of the inner-transition elements.

Atoms of this element like to lose 3e– to become isoelectronic with neon.

This nonmetal has the most unpaired electrons in the second period.

Atoms of this element like to gain 2e– to become isoelectronic with xenon.

If an atom of this element were to lose 6e– it would have 86 electrons.

This neutral atom has an outer shell configuration of 3s2.

This neutral atom has an electronic configuration of 1s22s22p2.

This neutral atom has an electronic configuration of 1s2.

The sixth element in the first row of the inner-transition elements.

The fifth transition metal in period 4.

_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____!

Adapted from Boyd, S.L. J. Chem. Educ. 2007 84 619–621.

137

LEWIS dOT STRuCTuRES OF IONIC COMPOuNdS

Lewis dot Structures of Ionic Compounds

1. Draw the Lewis dot structure (drawing) for each of the elements in the table in the first blank column.

2. In the second column, determine whether each element will likely gain electrons (become an anion) or lose electrons (become a cation) and draw the ion it forms. For cations, write the element symbol and the ion charge. For anions, include Lewis dots in your drawing and square brackets around the structure and the ion charge as a superscript.

3. Write the Lewis dot structure for the ionic compound which forms between these two elements. Show the charges and the chemical formula as shown in example 3a.

a. Barium and Oxygen

b. Aluminum and Flourine

c. Potassium and Iodine

d. Calcium and Nitrogen

Element Lewis dot Structure

Lewis dot Structure after

Reaction (HINT: Cation,

Anion, or Neither?)

Ba2+ [ O ]2– BaO

Ba O

138

LEWIS dOT STRuCTuRES OF IONIC COMPOuNdS

Element Lewis dot Structure

Lewis dot Structure after

Reaction (HINT: Cation,

Anion, or Neither?)

139

TuTORIAL: LEWIS dOT STRuCTuRES OF COVALENT COMPOuNdS

Tutorial: Lewis dot Structures of Covalent Compounds

STEPS FOR dRAWINg A LEWIS STRuCTuRE 1. Calculate the total number of valence electrons (v.e.) for the compound.

# of v.e. = element’s group A # on the periodic table

■ For polyatomic anions, add the number of electrons equal to the charge. ■ For polyatomic cations, subtract the number of electrons equal to the charge.

2. Figure out the central atom (which atom goes in the middle?)

■ Usually it’s the one that appears only once in the formula or the one that yields the most symmetric structure (but not always) or atom that is nearest to metals. For more complicated cases, it may be underlined.

■ Sometimes there can be more than one central atom. Hydrogen and halogens like F, Cl, Br are rarely ever central atoms.

3. Make single bonds to the central atom. Add remaining electrons in pairs to surround- ing atoms so each atom gets an octet of electrons. Give pairs of electrons to the outer atoms first. Do all atoms have an octet yet?

4. If there are any remaining electrons, place them in pairs on the central atom. Check if all atoms have an octet.

5. If you run out of electrons and there are atoms without an octet, then use one or more pairs of non-bonding electrons (lone pairs) to bond to the central atom to form double or triple bonds.

EXAMPLE

NO2 –

140

LEWIS dOT STRuCTuRES 1

Lewis dot Structures 1

Draw the Lewis dot structure for each of these substances.

NF3

CS2

NH4 +

SCl2

PH3

SO4 2–

NO2 –

SiCl4

ClO3 –

CH2O

141

VSEPR / MOLECuLAR SHAPE

VSEPR / Molecular Shape

The 2-D Lewis structure of each molecule is provided (but not necessarily its 3-D shape). Use VSEPR to identify the electron geometry and molecular shape for each of the follow- ing. NOTE: Some are drawn with an incorrect shape!

Lewis Structure

OH H

Geometry:

Shape:

C O

Geometry:

Shape:

NO O

Geometry:

Shape:

C N

Geometry:

Shape:

Geometry:

Shape:

SO O

Geometry:

Shape:

O F F

Geometry:

Shape:

C S

Geometry:

Shape:

P ClCl

Geometry:

Shape:

N H H

Geometry:

Shape:

Geometry:

Shape:

C HH

Geometry:

Shape:

S O O

Geometry:

Shape:

OO O

Geometry:

Shape:

As HH

Geometry:

Shape:

142

LEWIS dOT STRuCTuRES, gEOMETRY, ANd SHAPE

Lewis dot Structures, geometry, and Shape

Formula of Compound Lewis dot Structure Electron geometry Molecular

Shape Bond Angle

Cl2 chlorine

H2O water

CH4 methane

NH3 ammonia

HF hydrogen fluoride

N2 nitrogen

143

LEWIS dOT STRuCTuRES, gEOMETRY, ANd SHAPE

Formula of Compound Lewis Structure Electron geometry Molecular

Shape Bond Angle

CO2 carbon dioxide

CO3 2–

carbonate

CH2O

formaldehyde

NH4 +

ammonium

SO2 sulfur dioxide

144

REVIEW: BONdINg

Review: Bonding

1. Write the Lewis structure of Mg.

2. Write the Lewis structure of the compound NaBr.

3. Use Lewis theory to predict the formula of the compound that forms between magne- sium and nitrogen.

4. Write the Lewis structure for CO.

5. Write the Lewis structure for H2CO.

6. Write the Lewis structure for the ClO– ion.

145

REVIEW: BONdINg

7. Write the Lewis structure for the NO2 – ion. Include resonance structures.

8. Predict the molecular geometry of ClNO (N is the central atom).

9. Predict the molecular geometry of the SO3 2– ion.

10. Predict whether the atoms will combine to form an ionic or covalent compound. Draw its Lewis structure.

a. I and I

b. Cs and Br

c. P and O

146

LABORATORY EQuIPMENT: NAMES ANd PICTuRES

Laboratory Equipment: Names and Pictures

Spatula

Wire Gauze

Rubber Cone

Rubber Policeman

Test Tube Brushes

Bunsen Burner

Crucible Holder

Test Tube Holder

Test Tubes

Ring Stand

Ring

Clay Triangle

Utility Clamp

Buret Clamp

Test Tube Rack

Pasteur Pipet and Bulb

147

LABORATORY EQuIPMENT: NAMES ANd PICTuRES

Buret

Measuring Pipet

Thermometer

Screw Cap Vial

Wash Bottle

Watch Glass

Erlenmeyer Flask

Beaker

Funnel

Crucible and Lid

Volumetric Flask

Stirring Rod

24 mL

EEGEE

STM

=0.05 mL 20114

200

450

600

850

1000 mL

250

200

150

100

mL 0

100

150

200

250

300 mL ± 5% ± 5%

400 mL

No. 1000

891 0

10 0. 1

3 2 1

10 0

90 80 70 60 50 40 30 20 10 0

10 20 30

100 mL

148

149

EXPERIMENT 1: RuBBER BANd BLAST!

Chem& 140 Workbook

EXPERIMENT 1 EXPERIMENT 1: RuBBER BANd BLAST!

We thank Cynthia Stanich from University of Washington, Seattle for permission to use this activity.

STudENT LEARNINg OBjECTIVES ■ Apply the scientific method to a simple problem. ■ Become more familiar with the vocabulary of science. ■ Investigate the meanings of accuracy and precision. ■ Get to know each other and have a bit of fun!

YOuR MISSION Your mission today is to figure out how to shoot a rubber band as far as possible by plac- ing it over the end of a finger, adjusting the angle of the finger relative to the ground, and stretching the rubber band some amount prior to releasing it. In other words, your stated purpose for this laboratory should concern the question “What factors cause a rubber band shot from the end of a finger to travel the greatest distance possible?” The follow- ing walks you through using the scientific method to answer this question. Each student should fill in their own worksheet to keep.

150

151

REPORT SHEETS | EXPERIMENT 1: RuBBER BANd BLAST!

REPORT SHEETS | Experiment 1: Rubber Band Blast! Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Turn in these Report Sheets only as your lab report. You do not need to include the previ- ous pages.

NAMES OF gROuP MEMBERS

(Please be in a group of 3–4 students!)

____________________________________________

1. What do you think is the purpose of this activity?

2. Identify All Possible Variables

With your group, brainstorm possible variables which may affect the distance the rub- ber band travels.

EXAMPLE

Paula can take big steps or little steps when running. These are independent variables because she manipulates them to vary the dependent variable, which may be speed. A con- stant variable might be the direction she runs, the clothes she wears, or something else.

3. Hypothesis

With these variables in mind, write a hypothesis about what will make a rubber band travel the farthest that you can test today.

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REPORT SHEETS | EXPERIMENT 1: RuBBER BANd BLAST!

4. Experimental Plan

Fill in the choices you make.

a. Pick one variable to be your independent (or manipulated) variable. ____________

b. Identify the dependent (or responding) variable. ___________________________

c. Will you have any constant (or control) variables? If so, what are they? _________ __________________________________________________________________

d. How will you test your independent variable?

dEFINITIONS

Accuracy is a measure of how close a measurement is to a true value. Precision is the mea- sure of how close together the measurements are.

e. Will you do any repeat trials? Will this choice affect your precision? Your accuracy?

5. Experimental Work

Show your work to your instructor before you do this step. Carry out the experiment and record your data in a table. As you began to collect data, did you find you needed to change anything in your experimental plan?

Our Data Table:

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REPORT SHEETS | EXPERIMENT 1: RuBBER BANd BLAST!

6. data Analysis and Visualization

Graph your data. Do you think you can throw any points out? Why? Remember that the independent (manipulated) variable always goes on the x-axis, while the dependent (responding) variable goes on the y-axis. Draw a best fit (or trend) line through your data points; DO NOT “connect the dots.” Ask if you do not know the difference.

Title: ________________________________________________________________________________

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REPORT SHEETS | EXPERIMENT 1: RuBBER BANd BLAST!

RuBBER BANd BLAST-OFF Once each group is satisfied that they have determined the scientifically proven method for the best way to shoot a rubber band, we will see which group can shoot their rubber band the farthest.

7. discussion

a. Do your experimental results uphold your hypothesis or contradict it?

b. If your experimental results did not fully match your hypothesis, write a new hypothesis.

c. How accurate do you feel your results are (how close to a true value)? Why?

d. How precise do you feel your data are (how close to other measurements)? Why?

e. Do your results allow you to answer the original question about what makes a rub- ber band travel farthest?

f. Are you ready to write a scientific law or a theory, or do you feel that more experi- mentation is needed?

155

EXPERIMENT 2: MEASuREMENTS

CHEM& 140 WORKBOOK

EXPERIMENT 2 EXPERIMENT 2: MEASuREMENTS

OBjECTIVES ■ Practice using scientific tools for measurement and recording proper measurements ■ Determine the number of decimal places to record based on the graduations of a mea-

suring tool

■ Practice metric conversions ■ Practice performing calculations rounding to the correct number of significant figures ■ Use different types of glassware, each with a different level of precision

INTROduCTION Measurements are essential to experimental sciences such as chemistry, physics, biology and geology. This experiment is intended to give practice in making accurate measure- ments and being aware of the level of precision possible with certain measuring tools.

Measurements in science are usually made using metric system (“SI”) units. The precision of a measurement is limited by the calibration of the measurement tool. A balance that gives readings to only the nearest 0.001 gram cannot be used to give masses to 0.00001 gram.

156

EXPERIMENT 2: MEASuREMENTS

Always include dimensional units in your measurements. They are required to perform your mathematical calculations. As such they act as mathematical objects that are at least as important as the numbers themselves.

MATERIALS ANd EQuIPMENT ■ Balance ■ Metric rulers ■ Meter stick ■ Alcohol thermometer

■ 125- or 150-mL Erlenmeyer flask ■ 10-, 25-, 50-, and 100-mL

graduated cylinders

■ Metallic solid unknown

PROCEduRE

MASS MEASuREMENTS

Your instructor will give directions on how to use the balance. Some general rules to follow when using the balances follow:

■ use the “tare” button. It will set the balance to zero grams, which is your starting point. ■ Always use a container or weighing paper for weighing chemicals. DO NOT place

chemicals directly on the balance pan!

■ Clean up any materials on or near the balance. ■ If a balance seems to be out of order, please tell your instructor. DO NOT attempt to

make adjustments to the balance yourself!

■ All measurements must be recorded using the Rules for Significant Figures. ■ In order to gain practice with the balance,

■ measure the mass of a penny. Include dimensional units. In this case the mass is measured in grams (“g”).

■ find an object in the lab that has a mass of about 10 grams.

LENgTH MEASuREMENTS

You will be using a provided ruler printed on paper and the plastic ruler in your lab bench drawer.

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EXPERIMENT 2: MEASuREMENTS

VOLuME MEASuREMENTS

There are 10-mL, 25-mL and 100-mL graduated cylin- ders available. Use the size most appropriate for your measurements. Consult your instructor if you are not sure what size to use.

1. Fill a 125- or 150-mL Erlenmeyer flask to the top mark with water. Then pour that water into gradu- ated cylinder(s) to measure and record its volume.

2. Measure the edge length of the provided cube using the plastic ruler.

TEMPERATuRE MEASuREMENTS

Temperature measurements are made using mercury thermometers, thermocouples, gas- filled thermometers, alcohol thermometers, etc. We will be using an alcohol thermometer.

Measurement errors can result from the way the thermometer is located in a liquid. We can minimize some sources of error if we observe the following practices:

■ Be sure the liquid is thoroughly mixed. ■ Position the thermometer probe away from the walls of the container. You may want

to attach a thermometer clamp to the pole of your experiment stand to keep the ther- mometer in place.

■ Allow the thermometer to be in contact with the liquid long enough for the temperature of the thermometer to reach equilibrium with the temperature of the liquid.

■ Temperatures should be measured to the precision allowed by the thermometer. If the thermometer scale reads to ±1.0 °C, only readings to the nearest degree are possible.

IdENTIFYINg TYPES OF gLASSWARE

Please ask your instructor to display a graduated pipet, a buret, and a volumetric flask, in addition to the beakers, Erlenmeyer flasks, and graduated cylinders that you can find in the shelves at your table.

Please carefully observe the scale (or graduation, increment, or division marks) on the glassware. Also, note in which direction the numbers increase.

90.0

80.0

70.0

60.0

50.0

40.0

30.0

20.0

10.0

100.0

50.0

40.0

Read to the Bottom of the Meniscus

Graduated Cylinder

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EXPERIMENT 2: MEASuREMENTS

Draw a beaker. Draw a graduated cylinder. Draw an Erlenmeyer flask.

Draw a pipet. Draw a buret. Draw a volumetric flask.

159

MEASuREMENTS PRACTICE/REFERENCE SHEET

Measurements Practice/ Reference Sheet

Option #1: Set up by your Instructor: Various types of chemistry lab glassware filled with volumes of water. Record a measurement for each and then fill in the other columns of information about that measurement and measuring device.

Option #2: On your own in lab: For each type of chemistry lab glassware that you encounter in lab, fill in the columns in the bold box section identifying the smallest increment, uncer- tainty, and decimal place to which you should record for that measuring device. Refer back to this table when making measurements in future experiments.

Your Recorded Measurement

# of Certain digits

# of uncertain

digits

Total # of Sig Figs

Measuring device

Smallest Increment

uncertainty (+/–)

Always Record to This

decimal Place

Decimal places:

hundreds tens ones decimal point tenths hundredths thousandths

100 10 1 . 1/10 1/100 1/100

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MEASuREMENTS PRACTICE/REFERENCE SHEET

161

REPORT SHEETS | EXPERIMENT 2: MEASuREMENTS

REPORT SHEETS | Experiment 2: Measurements Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Turn in these Report Sheets only as your lab report. You do not need to include the previ- ous pages.

You will be graded based on correct number of significant figures reported for each mea- surement and calculated value, as well as correct unit abbreviations.

MASS Record the mass of a penny. _________________

What is an object in the lab that has a mass of approximately 10 g? _________________

LENgTH Record the length and width in centimeters of the provided piece of post-it note paper using the provided paper ruler. Include proper unit abbreviations in your measurements.

Length: _________________ Width: _________________

Calculate and report the area of this provided piece of post-it note paper determined using this ruler. Include proper unit abbreviations. Show your work.

Area: _________________

Record the length and width in centimeters of the provided piece of post-it note paper using the plastic ruler in your lab bench drawer.

Length: _________________ Width: _________________

Calculate and report the area of this provided piece of post-it note paper determined using this ruler. Include proper unit abbreviations. Show your work.

Area: _________________

What is an object that has a length of about 10 cm? _________________

What is an object that has a length of about 1 m? _________________

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REPORT SHEETS | EXPERIMENT 2: MEASuREMENTS

VOLuME

VOLuME OF WATER IN AN ERLENMEYER FLASK

Fill a 125- or 150-mL Erlenmeyer flask to the top mark with water. Measure this volume of water using a combination of appropriate graduated cylinder(s). Record all volume read- ings and calculate the total volume. All volumes must be reported to the appropriate num- ber of significant figures and with correct unit abbreviations.

Volume readings (in mL): _________________ _________________

Total volume (using sig fig rules for addition): _________________

VOLuME OF METALLIC SOLId uNKNOWN

(Use the same metallic solid for the density measurement.)

using the plastic ruler, measure the edge length of the provided cube in centimeters. Report it to the appropriate number of significant figures and with correct unit abbreviation.

Edge length of cube: _________________

Calculate the volume of the provided cube in cubic centimeters. Report it to the appropri- ate number of significant figures and with correct unit abbreviation.

Volume of cube: _________________

Show your work here.

Calculate the volume of the provided cube in cubic millimeters. Report it to the appropriate number of significant figures and with correct unit abbreviation.

Volume of cube: _________________

Show your work here.

TEMPERATuRE Observe the level of graduation of a thermometer and sketch a small portion of it below. Record the temperature of the tap water in the lab. Report it to the proper number of signif- icant figures and with correct unit abbreviation.

Temperature of tap water: _____________

Search the internet to find the freezing point of water in °C: _____________

Search the internet to find the boiling point of water in °C: _____________

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LAB EXERCISE EXPERIMENT 2: MEASuREMENTS

LAB EXERCISE Experiment 2: Measurements Name: __________________________________________________ Section: __________

1. What are the base units (and symbols) of measurements of the following in metric system?

a. Mass _______________________________

b. Length ______________________________

c. Temperature ________________________

d. Volume _____________________________

2. Write the full names of each of the following units:

a. mL __________________________________

b. ng __________________________________

c. km _________________________________

d. °C __________________________________

3. Write symbols for each of the following units:

a. Liter ________________________________

b. micrometer _________________________

c. kilogram ____________________________

d. milligram ___________________________

4. Convert 5.00 cm into mm. Use the proper number of significant figures.

5. Which is better to use for measuring volume: a 100 mL beaker or a 100 mL graduated cylinder? Explain.

6. In your own words, explain how you use the graduations on a piece of glassware to determine how many decimal places to record a measurement using that glassware.

164

LAB EXERCISEEXPERIMENT 2: MEASuREMENTS

165

EXPERIMENT 3: dENSITY

CHEM& 140 WORKBOOK

EXPERIMENT 3 EXPERIMENT 3: dENSITY

OBjECTIVES ■ Practice reporting measurements to the correct number of significant figures ■ Practice using significant figures in calcualtions ■ Determine the density of a liquid. ■ Determine the density of a solid by water displacement method. ■ Distinguish between intensive and extensive properties ■ Use empirical data to identify an unknown sample

INTROduCTION The density of a substance is an intensive physical property. This means that its value does not depend on the amount of the substance being measured. It can be used to help identify a material. The formula for calculating the density D is

D = mass in grams

volume in milliliters

or, more generally

D = mass

volume

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EXPERIMENT 3: dENSITY

Intensive properties do not depend on the amount of sample present. Some examples of intensive properties are luster, color and boiling point. Extensive properties are dependent on the amount of sample, some examples include mass and volume. Density is an example of an intensive property.

To determine the density of a substance you must determine the mass of the substance, a value that can be obtained from weighing the object on a balance.

You must also determine the volume of the substance. The volume of a substance can be determined in different ways.

FOR OBjECTS/SuBSTANCES WITH A dEFINEd SHAPE

The first, most straight forward method is to measure the volume in a graduated cylinder. This method is especially useful for determining the volume of liquids but does not work for accurately determining the volume of solids.

If a solid has measurable dimensions, then the volume can be determined using geom- etry. For a rectangular solid volume is determined by the equation: where values for the length (L ), width (W ), and height (H ) are multiplied against each other.

volume of a rectangle = L × W × H

In the case of a cubic solid the length, width, and height at all the same value and the equa- tion simplifies to

volume of a cube = L 3

FOR OBjECTS/SuBSTANCES WITH AN AMORPHOuS SHAPE

Determining the volume of an amorphous (undefined shape) solid can be more difficult than the other two methods. Direct measurement in a graduated cylin- der does not account for the air that takes up some of the space and there are no measurable dimensions to apply to a geometric formula.

For an amorphous solid the volume can be determined by the water displacement method. Fill a graduated cylinder with water. Record the initial volume. Place the substance in the water and the water level rises. Record the final volume. The difference is due to the volume of the substance added.

Figure 1. Displacement of water using an amorphous solid

167

EXPERIMENT 3: dENSITY

dENSITY OF A SOLId BLOCK

1. Obtain a solid metal block from your instructor. It should have a number written on it. Make sure to record this number on your report sheet as “block # ____.”

2. Weigh the solid to the correct amount of significant figures.

3. Using a ruler, measure each side of the block in units of cm. Record using the correct number of significant figures.

4. Calculate the volume in cubic centimeters (cm3).

5. Calculate the density (in grams per cubic centimeter).

6. Determine the identity of the metal using the table of densities.

7. Check with your instructor on what the identity of the metal is.

8. Based on the correct answer, calculate the % error:

% error = ⎛ ⎜ ⎝

experimental value – true value ⎞ ⎟ ⎠

× 100% true value

dENSITY OF WATER

1. Weigh a clean, dry 10-mL graduated cylinder. Record its mass.

2. Add tap water to the graduated cylinder, bringing the water level to between 8 and 9 mL. Record the volume of the water to the correct number of significant figures.

3. Weigh the graduated cylinder with the water. Calculate the mass of the water by sub- tracting the mass of the cylinder from the mass of the cylinder plus the water.

4. Use a thermometer to measure the temperature of the water to the nearest °C.

5. Calculate the density from the mass of water and the measured volume using the for- mula for density.

dENSITY OF AN AMORPHOuS SOLId

1. Obtain a vial containing metal shots. Record the number that is written on the vial on your report sheet.

2. Carefully empty the contents of the entire vial onto a tared piece of weigh paper. Record the mass of the metal shot.

3. Fill a 10 mL graduated cylinder with water so it is about a quarter full. Record the start- ing volume using the correct number of significant figures.

168

EXPERIMENT 3: dENSITY

4. Carefully slide, without splashing, the unknown solid objects into the graduated cylinder so it becomes totally submerged. Observe that the water level rises. It should be below 10 mL in volume. If it is not, start with less water initially.

5. Read and record the final water level (final vol- ume, V final).

6. Calculate the volume of the solid object by sub- tracting the initial volume from the final vol- ume. V solid = ∆V water = V final – V initial

7. Calculate the density of the solid object using the formula for density and identify which metal it is likely to be.

8. Obtain the true identity of the solid from your instructor and calculate % error.

9. Dry the unknown solid and return it in the vial provided.

density of Water as a Function of Temperature1 density of Various Metals

Temperature (°C) Density of Water (g/mL) Material Density at 25°C (g/mL)

20 0.9982 Aluminum 2.70

21 0.9980 Zinc 7.14

22 0.9978 Iron 7.87

23 0.9975 Brass 8.55

24 0.9973 Copper 8.94

25 0.9970 Lead 11.3

Figure 2. Graduated cylinder before and after addition of unkown solid

Water

Solid

Final Level

Initial Level

1 From the Handbook of Chemistry nand Physics, 86th edition, 2005–2006.

169

REPORT SHEETS | EXPERIMENT 3: dENSITY

REPORT SHEETS | Experiment 3: density Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Turn in these Report Sheets only as your lab report. You do not need to include the previ- ous pages.

You will be graded based on correct number of significant figures reported for each mea- surement and calculated value, as well as correct unit abbreviations.

dENSITY dETERMINATION FOR THE METAL BLOCK BLOCK # _________

Report all measurements and calculated values to the appropriate number of significant figures and with correct unit abbreviations.

1. Mass of the metal block: ____________

2. Length, width and depth of the metal block: __________ __________ __________ cm

3. Show the calculation for the volume of the metal block (include units). Use sig figs.

4. Show the calculation for the density of your metal block (include units). Use sig figs. This will be the “experimental value.”

5. Based on the table of densities, your block is made out of ____________________________

6. The correct identity of your block is ____________ which has a density of ___________. (This will be the “true value.”)

7. Calculate the % error using this equation:

% error = ⎛ ⎜ ⎝

experimental value – true value ⎞ ⎟ ⎠

× 100% true value

170

REPORT SHEETS | EXPERIMENT 3: dENSITY

dENSITY dETERMINATION FOR WATER Report all measurements and calculated values to the appropriate number of significant figures and with correct unit abbreviations.

Mass of Empty 10-mL Graduated Cylinder

Mass of Empty 10-mL Graduated Cylinder Plus Water

Calculated Mass of water

Volume of Water

Calculated Density of Water

Temperature of the Water to the Nearest Degree C

1. Fill in the table for the mass of the cylinder empty and with water. Show your mass cal- culation here and report it as “mass of water” in the table above:

2. Show your density calculation here and report the density in the table above. This is your “experimental value.”

3. Use the table in this lab to find the density of water at the temperature you measured and report it here (include units): ______________. This is your “true value” for the density.

4. Calculate your percent error for the density of water at this temperature. Show your work here and use sig figs.

% error = ⎛ ⎜ ⎝

experimental value – true value ⎞ ⎟ ⎠

× 100% true value

Your % error: _________

171

REPORT SHEETS | EXPERIMENT 3: dENSITY

dENSITY dETERMINATION OF A METALLIC SOLId VIAL # _______

Report all measurements and calculated values to the appropriate number of significant figures and include correct unit abbreviations.

Mass of solid object

Initial volume of Water in _____-mL Cylinder (Record the size of the cylinder you used.)

Final Volume of Water and Solid Object

Calculated Volume of Solid Object

Calculated Density of Solid Object

1. Fill in the mass, initial volume and final volume into the table above. Show your volume calculation and report the volume of the solid in the table above. Use sig figs.

2. Show your density calculation and report it in the table above. Use sig figs.

3. Based on the table of densities, your solid is made out of ____________________

4. The correct identity of your solid is ____________ which has a density of ___________. (This will be the “true value.”)

5. Calculate your percent error for the density of the solid object. Show your work here:

% error = ⎛ ⎜ ⎝

experimental value – true value ⎞ ⎟ ⎠

× 100% true value

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LAB EXERCISE | EXPERIMENT 3: dENSITY

LAB EXERCISE | Experiment 3: density Name: __________________________________________________ Section: __________

1. a. Write the equation used to calculate density.

b. If the density and volume of a substance are given, write an equation to calculate the mass of this substance.

c. If the density and mass of a substance are given, write an equation to calculate the volume of this substance.

2. Read the procedure. In your own words, summarize how you will determine the density of an amorphous solid.

3. What are the commonly used units of density (in the metric system) for each of the following:

a. Solids: _______________ b. Liquids or solutions __________

c. Gases _______________

4. a. Describe in your own words the meaning of the term physical property.

b. Is density a physical property or a chemical property? ___________________________

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LAB EXERCISE | EXPERIMENT 3: dENSITY

5. Refer to the Introduction section for this experiment.

a. Describe in your own words the difference between the terms extensive property and intensive property.

b. Is the denstiy an extensive property or an intensive property? ________________

6. Search the internet to find:

a. Density of ice at 0°C: _________

b. Density of water at 20°C: _________

c. Density of steam at 100°C: _________

7. What is unusual about density values of water in its different states?

8. Why do density values change with a change in temperature? (HINT: What part of the equation for density might change with temperature? What do atoms do as heat is added to a sample?

175

EXPERIMENT 4: CHEMICAL CHANgES

Chem& 140 Workbook

EXPERIMENT 4 EXPERIMENT 4: CHEMICAL CHANgES

From CHEM&121, Introduction to Chemistry (Inorganic) by Karyn Mlodnosky, Jennie Mayer, Gina Fiorini at http://opencourselibrary.org. Licensed under CC BY 3.0, https://creativecommons.org/licenses/by/3.0/.

INTROduCTION We can characterize chemical substances by their properties. Some of these are chemi- cal properties (i.e., how they react with other substances), and some of these are physical properties. Most chemical changes are accompanied by changes in physical properties. Since many physical properties can be observed by an experimenter’s senses, changes in physical properties are very often used to detect and to provide evidence that chemical changes have taken place.

Simple examples of evidence of chemical change include: a temperature change away from room temperature, changes in phase (a gas, a liquid, or a solid), change in color, solu- bility or precipitation (forming a new solid), how clear a solution is, and anything new and different or unexpected.

OBjECTIVES In this experiment you will

■ Observe physical changes as evidence that chemical changes have taken place. ■ Use vocabulary terms relating to physical and chemical changes. ■ Reinforce the practices of chemical safety and waste disposal.

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EXPERIMENT 4: CHEMICAL CHANgES

To perform this laboratory you do not need any prior chemical experience. You will need to be familiar with our laboratory’s safety procedures and follow them carefully. You will also need to read labels with care and to follow instructions exactly.

During this experiment you will need to make observations. Some changes are obvious, some are subtle; some observations are relevant to what you are studying, and some are not. To make a good observation you need to make a comparison between an initial and a final condition in order to note a change, or better yet, make a comparison between two samples, one of which differs only by the variable you are interested in. The best observa- tions are those that can readily be verified by another experimenter.

Read the section in your textbook that discusses physical and chemical changes. Read the sections in the laboratory manual that discuss safety and waste disposal practices. Then, using your book’s index/glossary, a dictionary, or another reference, complete the Prelab assignment for this laboratory and be prepared to have it checked before beginning work.

HAZARdS The acids and bases used in this experiment can damage eyes and harm skin. Safety gog- gles must be worn by everyone if any chemical work, including cleanup, is in progress in the laboratory. Wash your hands before leaving lab. You will be rotating to different experi- mental stations: do not carry out any procedure until you have read through the entire pro- cedure for that section!

PROCEduRE

gENERAL INSTRuCTIONS

■ Work with your lab partner to perform the experiments at each of the six stations. Each station will require between two and ten minutes. You may complete the experiments in any order you wish.

■ Make sure the station and test tubes are clean for the next group. Since the test tubes will be filled with mixtures containing water, it is not necessary to use dry test tubes.

■ Fill in your observations on the first page of the report. As you proceed through the sta- tions, read (but do not fill in) the questions on each reaction that are on the second page of the report; complete these questions after you have completed all of the experiments.

177

EXPERIMENT 4: CHEMICAL CHANgES

REACTION STATIONS

1. Calcium Chloride Solution with Sodium Carbonate Solution

a. Place about 20 drops of 5% sodium carbonate (Na2CO3) solution in a test tube.

b. Add about 5 to 10 drops of 5% calcium chloride (CaCl2) solution.

c. Record your observations.

d. Pour the solution into the waste container and rinse the test tube with water. You do not need to dry the test tube. Leave the station clean for the next group.

2. Formation of a Potassium Nitrate Solution

a. Put about a teaspoon of postassium nitrate (KNO3, a solid) in a test tube.

b. Touch the bottom of the test tube to the inside of your arm to get a sense of its temperature.

c. Add deionized water to the test tube (about 1/3 full) and stir with a glass stirring rod.

d. Touch the bottom of the test tube to the inside of your arm again. If you don’t feel a difference, add more potassium nitrate. Record your observations.

e. When finished, place the contents of the test tube in the waste container. Rinse the test tube with water. You do not need to dry the test tube. Leave the station clean for the next group.

3. Limestone with Acidic Water

a. Place a limestone chip (use smooth white chips for best results) on a watch glass and add about 5 drops of deionized water. Did anything happen?

b. Add 5 drops of 3 M hydrochloric acid (HCl). Did anything happen? What is different?

c. When finished observing, rinse the limestone chips with water and dry for reuse by others. Rinse the watch glass with water into the waste container. You do not need to dry the watch glass. Leave the station clean for the next group.

4. Vegetable dye with Acidic and Basic Water

a. Obtain 3 test tubes and fill each one with half a pipet full (0.5–1 mL) of red cabbage juice.

b. To the first test tube, add an equal amount of deionized water. To the second test tube, add an equal amount of 0.1 M HCl (hydrochloric acid). To the third test tube, add an equal amount of 0.1 M NaHCO3 (baking soda).

c. Note any color changes and record them on your data sheet.

d. Pour the liquids into a waste container. Thoroughly rinse the test tubes with water. Leave the station clean for the next group.

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EXPERIMENT 4: CHEMICAL CHANgES

5. Bleaching of Paper

a. Obtain three test tubes and label them 1, 2, 3. To the first one, add approximately 5 mL of deionized water. To the second one, add approx. 5 mL of chlorox bleach. To the third one, add approx. 5 mL of regular chlorine bleach.

b. Using scissors, cut three strips of brown/colored paper or bright colored cloth (enough to fit inside the test tube with some of it submerged in the liquid and some of the paper above the liquid). Place one strip into each test tube.

c. Let it sit for 1 min at room temperature. Record your observations.

d. Heat the above test tubes in the hot water bath for 5 minutes. Record observations.

e. Remove the paper from the test tubes. Place the used paper in the garbage. Discard your bleach solutions in the waste container. Rinse out the test tubes with plenty of water, and leave them for the next group (they do not need to be dry). Leave the station clean for the next group.

6. Aluminum with Copper Ions

a. Obtain a test tube and add about 2–3 mL of 5% (m/v) CuCl2•2H2O (copper (II) chlo- ride dihydrate) solution.

b. Record your observations. (Include color, temperature, states of matter, etc.)

c. To this test tube, add an approximately 1-inch piece of aluminum wire. Touch the test tube to the inside of your arm. Wait for a minute and record your observations. (Include color, temperature, states of matter, etc.)

d. When finished, place the aluminum wire and copper chloride solution in a labeled waste container. Rinse the test tube with water (it does not need to be dry). Make sure the station is clean for the next group. Leave the station clean for the next group.

179

REPORT SHEETS | EXPERIMENT 4: CHEMICAL CHANgES

REPORT SHEETS | Experiment 4: Chemical Changes Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Turn in these Report Sheets only as your lab report. You do not need to include the previ- ous pages.

dATA For each reaction, write down your observation(s) before and after the reaction that you believe provide evidence that a chemical change took place. Be clear and concise; use only enough detail to communicate your findings.

Reactions Observations

Before After

1. Calcium Chloride + Sodium Carbonate

2. Formation of Potassium Nitrate Solution

3. Limestone + Acidic Water

4. Vegetable Dye + Acid/Base Solutions

5. Bleaching Paper

6. Al Wire + Copper (II) Chloride Solution

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POST-LAB QuESTIONS 1. Reaction of Calcium Chloride Solution with Sodium Carbonate Solution: What term

from the pre-laboratory assignment (besides “chemical change”) describes this result? Explain.

2. Formation of a Potassium Nitrate Solution: It could be correctly argued that this is a physical change or that this is a chemical change. a. What did you observe that gives evidence that it is a chemical change?

b. For what practical purposes do you think that this reaction can be used?

3. Reaction of Limestone with Acidic Water:

a. Were the results for the deionized water similar or different than for the acid?

b. What do you think acid rain would do to a limestone statue or gravestone?

4. Reaction of a Vegetable dye with Acidic and Basic Solutions: What was the purpose of adding deionized water to one of test tubes containing cabbage juice?

5. Bleaching of Paper:

a. What did you observe that indicates that this is a chemical change and not just a physical change?

b. Did the reaction occur instantaneously? What do you think would happen if you waited longer?

6. Aluminum with Copper Ions:

a. Based on your observation and chemicals used, what substance do you think deposited on the aluminum wire? What observations explain your choice?

b. What practical application is there for what you observed?

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PRE-LAB | EXPERIMENT 4: CHEMICAL CHANgES

PRE-LAB | Experiment 4: Chemical Changes Name: __________________________________________________ Section: __________

Complete the following questions BEFORE class. Refer to your textbook or the web if needed.

1. Read through the experimental procedure. Most of the experiments require use of clean test tubes, but why is it not necessary to use dry test tubes?

2. What should you do with any waste materials from this laboratory? If you take excess chemicals that you don’t use, what should you do with it?

3. Use your textbook’s glossary and index to define the following terms:

a. physical property

b. chemical change

c. reactants

d. products

e. precipitate

f. solution

4. List three types of evidence that a chemical change has taken place. (Refer to Section 6.1 in your Timberlake textbook.)

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PRE-LAB | EXPERIMENT 4: CHEMICAL CHANgES

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CHEM& 140 WORKBOOK

EXPERIMENT 5 EXPERIMENT 5: EMPIRICAL FORMuLA

OBjECTIVE Determine the empirical formula of magnesium chloride.

BACKgROuNd The word equation for the reaction we will examine is

Solid magnesium reacts with hydrochloric acid (an aqueous solution of hydrogen chlo- ride) to produce hydrogen gas and an aqueous solution of magnesium chloride.

Mg (s) + HCl (aq) → H2 (g) + ? Hydrochloric acid will be added to a measured amount of magnesium. The amount of hydro- chloric acid solution used is in excess to ensure that all of the magnesium reacts. The reac- tion is quite vigorous and a substantial amount of heat will be released. Reactions that release heat are called exothermic. When the reaction is complete, the hydrogen will have escaped into the room. The beaker will contain magnesium chloride dissolved in water and excess hydrochloric acid. In order to determine the mass of magnesium chloride, the water and excess hydrochloric acid must be removed. This can be accomplished by heating the sample to dryness. The remaining powder is the magnesium chloride.

184

HAZARd HCl is corrosive. Avoid contact with skin and clothing. Avoid inhalation. Perform all heating steps in the hood. goggles must be worn at all times.

PROCEduRE 1. Weigh a clean dry and empty 250 ml beaker and record its mass.

2. Weigh approximately 1.00 gram of magnesium in this beaker and record mass.

3. While working in the hood, slowly add approximately 50 mL of 3 M HCl (hydrochloric acid solution) to the magnesium. Record observations as the reaction proceeds. You may carefully touch the beaker. Do not inhale any of the vapors.

4. After the metal has all reacted, heat the reaction mixture using a hot plate in the hood to boil off the water and drive off the excess HCl. Avoid splattering of the hot liquid when the volume is low by heating gently.

5. When the material in the beaker seems dry, remove the beaker from hot plate and allow it to cool to room temperature. Weigh the beaker and the magnesium chloride. Record the mass.

6. After weighing, reheat the beaker and solid for 5 minutes on the hot place in the hood. If you have water droplets condensing towards the top of the beaker, use a cool flame of a Bunsen burner to remove them. Allow the beaker to cool and reweigh. Record the mass. Use the lowest mass reading for calculations. Do not discard the dried magnesium chloride until you complete your calculations and get these checked by your instructor.

7. Scrape off the solid magnesium chloride in beaker and place it in a labeled waste con- tainer. You may find it helpful to squirt a little water into the beaker during scraping.

8. Please rinse the beaker with plenty of water in the sink before putting it away.

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REPORT SHEETS | EXPERIMENT 5: EMPIRICAL FORMuLA

REPORT SHEETS | Experiment 5: Empirical Formula Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Show all calculations. Record all data and report all calculated results to the appropriate number of significant figures. Include units.

Mass of beaker and magnesium _________

Mass of empty beaker _________

Mass of magnesium _________

Moles of magnesium _________

Mass of beaker and magnesium chloride after the first heating _________

after the second heating _________

use the mass after the second heating to complete your calculations.

Mass of magnesium chloride _________

Mass of chloride _________

Moles of chloride _________

Calculate the ratio of moles of the Mg to the moles of Cl ____ mole Mg : ____ mole Cl (Do not round to nearest whole number yet.)

Empirical formula _________ (The metal is written first. Round your ratio to the nearest whole number.)

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REPORT SHEETS | EXPERIMENT 5: EMPIRICAL FORMuLA

POST-LAB QuESTIONS 1. Why do you think the solid sample of magnesium chloride was re-heated?

2. What changes occurred during the entire experiment? Record your observations here.

3. Write a balanced chemical equation for the reaction that occurs as solid magnesium reacts with hydrochloric acid solution to produce hydrogen gas and an aqueous solu- tion of magnesium chloride. Include physical states of matter. (If your experiment did not result in the correct formula for magnesium chloride, do not use your formula in this equation, instead use the correct formula.)

187

LAB EXERCISE | EXPERIMENT 5: EMPIRICAL FORMuLA

LAB EXERCISE | Experiment 5: Empirical Formula Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Show your work. Use appropriate number of significant figures and units.

1. A compound containing nitrogen and oxygen is decomposed (separated into nitro- gen and oxygen) in the laboratory and produced 1.78 g of nitrogen and 4.05 g of oxy- gen. Calculate the empirical and molecular formula of this compound by following the steps below.

a. Calculate the moles of nitrogen formed during the decomposition:

b. Calculate the moles of oxygen formed during the decomposition:

c. Calculate the simplest ratio of moles of nitrogen and oxygen:

d. Write the empirical formula of this compound containing nitrogen and oxygen.

e. Calculate the empirical formula mass of this compound.

f. The molar mass of this compound is 92.0 g/mol. Calculate the ratio of the molar mass to the empirical formula mass. Based on this ratio, what is the molecular formula?

g. Identify this compound as an ionic or molecular compound. Circle your answer.

h. Write the systematic name of this compound.

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LAB EXERCISE | EXPERIMENT 5: EMPIRICAL FORMuLA

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CHEM& 140 WORKBOOK

EXPERIMENT 6 EXPERIMENT 6: CHEMICAL REACTIONS

OBjECTIVES Become more familiar with different types of chemical reactions by

1. Conducting several reactions

2. Discussing your observations and recording the results

3. Writing correct formulas and balanced chemical equations

INTROduCTION Chemical reactions take place all around us and even inside us. Chemical reactions can be relatively simple, like the combination of oxygen and hydrogen to form water. They can also be complex, like the synthesis of large protein molecules from thousands of much sim- pler molecules.

Chemical reactions are changes in the way atoms and molecules combine with each other to form different kinds of matter. Because it is not generally possible to actually see the atoms and molecules in a chemical reaction directly, we need other means to identify when and how a chemical reaction has taken place. Fortunately, many chemical reactions pro- duce observable changes that are easily detected. Color changes, the formation of gases, the formation of solids, heat absorption or emission, and light emission can all be evidence

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EXPERIMENT 6: CHEMICAL REACTIONS

of a chemical reaction. However, they are not definitive evidence that a chemical reaction has taken place. Only chemical analysis can show that the substances we started with ini- tially have been changed into other substances.

In a chemical reaction, the substances present before the chemical changes are called reactants, and the substances present after the change are called products. Chemical reactions convert reactants into products.

We represent a chemical reaction by a balanced chemical equation, where reactants are stated on the left and products on the right. Instead of the “=” we use in a mathematical equation, chemical equations use “→”to indicate the conversion of reactants into products. In writing the chemical formula of different substances (reactants and products) we also include the states of matter of the reactants and products using the following convention:

(g ) – gas (l ) – liquid (s ) – solid (aq ) – aqueous

PROCEduRE Caution: You must wear safety goggles at all times during this experiment.

Follow all directions carefully!

■ You will be rotating to different experimental stations. ■ Use fume hoods for reactions 6, 7, and 8.

For each of the following reactions answer the following questions:

■ How did you know a chemical reaction took place? ■ Formula of reactant(s), refer to your text book on writing correct formulas. ■ Formula of product(s) ■ Write the balanced equation for the reaction (including all physical states). ■ Type of chemical reaction

REACTION 1

Place 10 drops of 0.2 M silver nitrate into a small test tube. Add 10 drops of 0.2 M sodium chloride to the same test tube.

Equation: Aqueous silver nitrate reacts with aqueous sodium chloride to produce solid silver chloride (also called precipitate) and aqueous sodium nitrate.

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EXPERIMENT 6: CHEMICAL REACTIONS

REACTION 2

Place a small piece of zinc into a small test tube. Add enough 3.0 M hydrochloric acid (HCl) to just cover the piece of zinc.

Equation: Solid zinc reacts with aqueous hydrochloric acid to produce aqueous zinc chloride and hydrogen gas.

REACTION 3

Place a small piece of magnesium ribbon (about 10 mm long) into a small test tube.

Add enough 0.1 M copper (II) sulfate to just cover the ribbon. Set aside for 15 minutes. Examine the contents of the test tube.

Equation: Aqueous copper (II) sulfate reacts with magnesium to produce aqueous mag- nesium sulfate and solid copper.

REACTION 4

Place 10 drops of 0.2 M barium chloride into a small test tube. Add 10 drops of 0.2 M sodium sulfate to it.

Equation: Aqueous barium chloride reacts with aqueous sodium sulfate to make solid barium sulfate and aqueous sodium chloride.

REACTION 5

Place 20 drops of hydrogen peroxide, H2O2, into a small test tube. Add a small amount of manganese (IV) oxide to the tube (just enough to cover the bottom of the test tube).

NOTE: Manganese (IV) oxide is a catalyst. It is neither a reactant nor a product: It speeds up the rate of the reaction. A catalyst is indicated in a chemical reaction by writing the for- mula of the catalyst over the arrow.

Equation: Aqueous hydrogen peroxide, H2O2, produces liquid water and oxygen gas.

REACTION 6 (uSE FuME HOOd)

Place a small amount of sulfur into a deflagration spoon. Heat over a Bunsen burner until the sulfur begins to burn.

Equation: Solid sulfur (S8) reacts with oxygen gas to yield sulfur dioxide gas.

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EXPERIMENT 6: CHEMICAL REACTIONS

REACTION 7 (uSE FuME HOOd)

Place a small amount (small scoop) of hydrated copper (II) sulfate into a small test tube. Heat the hydrate over a Bunsen burner until the chemical reaction is complete (the reac- tion should be obvious).

NOTE: A hydrate is a compound that contains water. The formula of a hydrate is writ- ten a little differently than formulas for other compounds. The formula for the hydrate of copper (II) sulfate is written CuSO4 · 5 H2O. This means there are five water molecules bound to one CuSO4 formula unit. (The correct name for this compound is copper (II) sul- fate pentahydrate).

Equation: When solid copper (II) sulfate pentahydrate is heated, solid anhydrous cop- per (II) sulfate and water vapor are produced. (Anhydrous means “no water.”)

REACTION 8 (uSE FuME HOOd)

Use the matches/lighter to ignite the liquid methanol, CH3OH, in the lamp by lighting the wick. Extinguish the flame by putting the cap over the wick.

Equation: Liquid methanol, CH3OH, reacts with oxygen to produce carbon dioxide gas and water vapor.

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REPORT SHEETS | EXPERIMENT 6: CHEMICAL REACTIONS

REPORT SHEETS | Experiment 6: Chemical Reactions Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

For each reaction, write observations, type of reaction, and a balanced chemical equation that includes physical states.

REACTION #1

Observations:

Type of reaction:

Equation:

REACTION #2

Observations:

Type of reaction:

Equation:

REACTION #3

Observations:

Type of reaction:

Equation:

REACTION #4

Observations:

Type of reaction:

Equation:

REACTION #5

Observations:

Type of reaction:

Equation:

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REPORT SHEETS | EXPERIMENT 6: CHEMICAL REACTIONS

REACTION #6

Observations:

Type of reaction:

Equation:

REACTION #7

Observations:

Type of reaction:

Equation:

REACTION #8

Observations:

Type of reaction:

Equation:

POST-LAB QuESTIONS 1. Summarize the major types of visual evidence you observed indicating that a chemical

reaction has occurred in your experiment reactions 1 to 8?

2. Why can the formation of a precipitate be evidence of a chemical reaction?

3. When a hydrate compound is heated, does the water get broken down into its compo- nent elements, hydrogen and oxygen?

4. When any hydrocarbon compound (CxHy) is combusted, what products are formed?

195

LAB EXERCISE | EXPERIMENT 6: CHEMICAL REACTIONS

LAB EXERCISE | Experiment 6: Chemical Reactions Name: __________________________________________________ Section: __________

1. List at least three possible types of evidence of a chemical reaction.

2. What is a reactant in a chemical reaction?

3. What is a product in a chemical reaction?

4. Give the four symbols for physical states of reactants and products and explain what each means.

5. Why must chemical equations be balanced?

6. Which of the following is a chemical reaction? If it is a chemical reaction, give your reasons.

Liquid water freezes to form solid ice _________

Copper turn greens on exposure to air _________

When sodium bicarbonate is combined with hydrochloric acid, bubbling is observed _________

Butane evaporating out of a butane lighter _________

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LAB EXERCISE | EXPERIMENT 6: CHEMICAL REACTIONS

7. Classify each of the following reactions as a synthesis, decomposition, single-dis- placement, double-displacement, or combustion.

2 Al (s) + 2 H3PO4 (aq) → 2 AlPO4 + 3 H2 (g) ____________________________ 2 K (s) + Br2 (l ) → 2 KBr (s) ____________________________ CuSO4 (aq) + 2 KOH (aq) → Cu(OH)2 (s) + K2SO4 (aq) ____________________________ CuCl2 (aq) → Cu (s) + Cl2 (g) ____________________________

8. Balance the following equations:

______Sb (s) + ______Cl2 (g) → ______SbCl3 (s)

______C2H6 (g) + ______O2 (g) → ______CO2 (g) + ______H2O (g)

______Fe (s) + ______HCl (aq) → ______FeCl3 (aq) + ______H2 (g)

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CHEM& 140 WORKBOOK

EXPERIMENT 7 EXPERIMENT 7: SYNTHESIS OF ALuM

OBjECTIVES ■ Apply the principles of synthetic chemistry, the process by which one pure substance

is transformed into another.

■ Practice stoichiometric calculations including theoretical and percent yield.

INTROduCTION In this experiment we will perform a synthesis—taking one or more compounds to form another. During the synthesis we will transform the aluminum in aluminum foil into a sub- stance known as alum, which has the chemical formula KAl(SO4)2 · 12 H2O. (This chemical formula is described on the next page.)

Alum is a versatile substance used in pickling, leather tanning, dyeing fabrics, treatment to fireproof textiles, and for water purification in municipal water treatment plants. Alum is also an ingredient in baking powder and in aftershave products. In medicine, it has appli- cations as an astringent (constriction agent) and a styptic (to coagulate the blood and stop the bleeding).

Aluminum foil is made of pure aluminum (Al).

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EXPERIMENT 7: SYNTHESIS OF ALuM

STEP 1

The first step in this experiment is to react the aluminum foil with potassium hydroxide (KOH) solution. The KOH helps to dissolve the thin oxide coating on the aluminum that pro- tects it from further oxidation. This is an oxidation-reduction reaction, in which the alu- minum is oxidized and the water is reduced. Transfer of electrons from the aluminum to the water produces hydrogen gas and hydroxide ions. The hydroxide ion then combines with the dissolved aluminum cation to form a complex ion Al(OH)4

–, the tetrahydroxyalu- minate ion.

The balanced chemical equation for this oxidation-reduction reaction is

2 Al (s) + 2 KOH (aq) + 6 H2O(l ) → 2 KAl(OH)4 (aq) + 3 H2 (g)

STEP 2

The second step of the procedure is to convert the KAl(OH)4 to alum by the addition of sul- furic acid (H2SO4). Addition of a slight excess of sulfuric acid to the products neutralizes the excess KOH and 4 OH- groups on the Al(OH)4

– ion. This is an acid-base neutralization reaction in which acid and base combine to form water. The solution then contains Al3+, K+ and SO4

2– ions in the proportions in which they are found in alum. The formula of crystal- line alum is KAl(SO4)2 ·12 H2O. There are 12 water molecules associated with each formula unit of the salt.

The alum has a limited solubility in water, so it precipitates from the solution. The balanced chemical reaction that occurs in this step is

KAl(OH)4 (aq) + 2 H2SO4 (aq) + 8 H2O (l ) → KAl(SO4)2 · 12 H2O (s)

STEP 3

Next, the solution containing the dissolved alum is cooled on ice. The solubility of alum decreases as the temperature is lowered. Cooling the alum solution increases the amount of solid alum that can be recovered.

To get pure alum, it is necessary to filter out the crystals and wash them with ethanol to remove the excess H2SO4. The ethanol will evaporate quickly and does not dissolve the alum. The crystals of alum can be dried at room temperature.

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EXPERIMENT 7: SYNTHESIS OF ALuM

SAFETY ■ You must wear safety goggles at all times during this experiment. ■ Gloves recommended for strong acids & bases. ■ 1.5 M KOH (caution: caustic/strong base) ■ 3 M H2SO4 (caution: corrosive/strong acid)

PROCEduRE 1. Cut a piece of aluminum foil 10 × 13 cm (this yields between 0.5–0.6 g Al).

2. Cut the aluminum into small squares, about 0.5 × 0.5 cm.

3. Weigh an empty 150-mL beaker and record its mass to the nearest milligram.

4. Add the 0.5–0.6 g of small pieces of aluminum (from Steps 1 & 2) and record the mass of the beaker containing the aluminum to the nearest milligram.

5. Measure 25 mL of 1.5 M KOH solution in a graduated cylinder.

Work under the fume hoods!

1. Add the 25 mL of 1.5 M KOH to the 150 mL beaker that contains the small pieces of aluminum.

2. Place the beaker on a hot plate set at low heat. Warm the solution: do not allow it to boil. The aluminum should react and be consumed in about 20 minutes.

3. Set up a glass funnel1 using a ring stand and iron ring to support the funnel.

4. Place the stem of the funnel into a second, clean 150-mL beaker.

5. Fold a large circular filter paper into quarters and place it into the funnel2. Pour a lit- tle water onto the filter paper so that it will cling to the inner surface of the funnel and stay in place.

6. Remove the beaker from the hot plate. Be sure that all aluminum has reacted.

7. Obtain 15 mL of 3 M H2SO4 and carefully and rapidly pour all of the H2SO4 into the bea- ker, stirring steadily to avoid formation of insoluble Al(OH)3 precipitate. If white crys- tals are visible, return the beaker to the hot plate and stir until all of the material is dis- solved. You may need to add up to 5 mL more of the 3 M H2SO4 if the white crystals do not disappear adequately.

1 This is a glass gravity funnel not a vacuum Büchner funnel.

2 If you haven’t installed filter paper into a funnel before, please consult the instructor.

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EXPERIMENT 7: SYNTHESIS OF ALuM

8. Pour the hot solution into the previously prepared glass filter funnel. Caution: Do not overfill the filter funnel and do not touch the filter paper with the stirring rod. The fil- trate (the liquid that has passed through the filter) should be clear and colorless. Save the filtrate for next step. Discard any residual crystals on the filter paper in a labeled waste container and the filter paper in the trash.

9. Allow the 150-mL beaker containing the filtrate to cool.

10. Take a 600-mL beaker and fill it halfway with crushed ice. Add water to this to create a 0 °C ice bath.

11. Place the 150-mL beaker, containing the filtrate, in the ice bath for about 20 minutes, allowing it to cool. Stir the filtrate mixture frequently as the crystals of alum form.

12. Insert a Büchner funnel into a vacuum filter flask and attach the flask’s nozzle to a vac- uum hose. Attach the other end of the hose to the vacuum nozzle in the hood.

13. Select a 7.5 cm filter paper. Using a pencil, label the filter paper with your name and record its mass. Place the filter paper evenly around the inside of the funnel so it can catch all of the alum you’ve produced.

14. Moisten the filter paper and turn on the vacuum to the Büchner funnel.

15. Remove the 150-mL beaker from the ice bath. Stir the contents and transfer all of the crystals to the Büchner funnel. You may use a plastic spatula to make sure you transfer all of the crystals produced.

16. Add two 10-mL portion of ethanol to wash the excess H2SO4 out of the crystals.

17. Continue to draw air through the Büchner filter for 15 minutes to dry the alum crystals.

18. Using the forceps (tweezers), remove the filter paper containing the alum crystals from the Büchner funnel and allow alum crystals to air dry till next week.

19. Weigh 10 g of sodium bicarbonate. Add it to the remaining solution in the filter flask. Swirl the flask until the bubbling ceases. Add more NaHCO3 till you do not observe any more fizzing. This is a neutralization reaction of the sulfuric acid H2SO4 by sodium bicarbonate NaHCO3:

H2SO4 (aq) + 2 NaHCO3 (s) → Na2SO4 (aq) + 2 CO2 (g) + 2 H2O (l ) 20. Pour the neutralized solution (filtrate) in the filtration flask down the sink1 and rinse

the flask.

21. After the crystals have dried, weigh the filter paper with alum crystals. Record the mass and calculate the mass of alum, KAl(SO4)2 · 12 H2O (s), crystals obtained. This is the actual yield of alum.

22. Calculate the theoretical yield of alum based on the mass of aluminum foil used as alu- minum is the limiting reactant in this reaction. Calculate the percent yield of alum after you recorded the mass of dried alum crystals.

1 The neutralized solution will not damage the pipes.

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REPORT SHEETS | EXPERIMENT 7: SYNTHESIS OF ALuM

REPORT SHEETS | Experiment 7: Synthesis of Alum Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

data

Mass of Empty 150-mL Beaker

Mass of Empty 150-mL Beaker + Aluminum

Mass of Aluminum Metal Used (calculated)

Mass of Filter Paper

Mass of Filter Paper + Alum Crystals

Mass of Alum Crystals Obtained (calculated)

Show your work and express your answer in correct number of significant figures based on your data. Use appropriate units.

1. Calculate the number of moles of aluminum that you used. __________________ Show your calculations.

2. Calculate the molar mass of alum KAl(SO4)2 · 12 H2O and round it to two decimal places. Include units. (HINT: The dot means each crystal of KAl(SO4)2 contains 12 H2O mole- cules. How does this affect its molar mass?) __________________ Show your calculations.

3. Calculate the maximum number of grams of alum that you could produce (theoretical yield), assuming that aluminum is the limiting reactant. __________________

4. Refer to your data table. Which piece of data is your actual yield of alum? Rewrite your actual yield here. __________________

5. Calculate the percentage yield of alum. __________________ Show your calculations.

% yield actual yield

× 100 theoretical yield

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LAB EXERCISE | EXPERIMENT 7: SYNTHESIS OF ALuM

LAB EXERCISE | Experiment 7: Synthesis of Alum Name: __________________________________________________ Section: __________

Show your work and report your answer in correct number of significant figures.

1. For the reaction, ammonia reacts with oxygen to produce nitrogen monoxide and water, answer the following questions:

_______NH3 (g) + _______O2 (g) → _______NO (g) + _______H2O (g) a. Balance the equation to whole number ratios.

b. How many molecules of water can be produced from 60. molecules of O2?

c. How many moles of O2 would be required to react with 0.330 moles of NH3?

d. If 2.00 g of NH3 are used, how many moles of NO can be produced?

e. If 1.11 moles of NH3 are consumed in a reaction, how many grams of H2O are produced?

f. How many grams of O2 would be required to react with 2.00 g of NH3?

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LAB EXERCISE | EXPERIMENT 7: SYNTHESIS OF ALuM

2. For the reaction of nitrogen gas with hydrogen gas to produce ammonia gas, NH3, answer the following questions:

a. Write a balanced chemical equation for this reaction, include physical states of matter.

b. Calculate the theoretical yield of ammonia, in grams, when 10.00 g of N2 is reacted.

c. Calculate the percent yield of this reaction when 10.00 g of N2 is reacted and 9.10 g of ammonia is produced.

HINT: See the previous part in which you calculated the theoretical yield of ammo- nia based on 10.00 g of N2 reacting. The actual yield is the actual amount produced in the reaction, not the theoretical (expected or calculated) yield or the starting reactant amount.

% yield actual yield

× 100 theoretical yield

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CHEM& 140 WORKBOOK

EXPERIMENT 8 EXPERIMENT 8: STOICHIOMETRIC ANALYSIS

OF AN ANTACId

From CHEM&121, Introduction to Chemistry (Inorganic) by Karyn Mlodnosky, Jennie Mayer, Gina Fiorini at http://opencourselibrary.org. Licensed

under CC BY 3.0, https://creativecommons.org/licenses/by/3.0/.

INTROduCTION In this lab, you will use the concept of stoichiometry to solve two sequential problems. First, you will try to determine the products of a certain reaction (below), choosing between three possibilities. Then, you’ll use your results of this first part to determine the amount of sodium bicarbonate in a common household substance.

CH3COOH (aq) + NaHCO3 (s) → CO2 (g) + ??? You’ve probably seen this reaction in elementary school—add a few drops of red food col- oring, and you have the classic “volcano reaction.” Or, you can perform it easily in your kitchen by mixing vinegar (dilute acetic acid) and baking soda (sodium bicarbonate). The most noticeable sign of the reaction is vigorous bubbling, a result of very rapid carbon diox- ide generation.

Gaseous carbon dioxide is one of the products, as you can see with your own eyes. (You can prove the gas to be carbon dioxide by collecting it in a flask, and inserting a burning match into the flask. The flame will be immediately extinguished.) Aside from carbon diox- ide, what else is produced by the reaction?

206

A chemist approaching this problem would most likely form some hypotheses about the other products, and then design experiments to evaluate which hypothesis is best sup- ported by experimental evidence. For this experiment, we’ll supply three possible reac- tions, shown below. Notice that they are all balanced equations.

A. CH3COOH (aq) + NaHCO3 (s) → 2 CO2 (g) + CH2O (aq) + Na+ (aq) + 3 H+ (aq) B. CH3COOH (aq) + NaHCO3 (s) → CO2 (g) + H2O (l ) + CH3COO–Na+ (aq) C. CH3COOH (aq) + 2 NaHCO3 (s) → CO2 (g) + Na2CO3 (aq) + H2O (l ) + 2 CH2O (aq)

Your job is to determine which of these three possibilities is correct, using some simple lab- oratory measurements. Focus on the differences between the three proposals: Reactions A and C produce formaldehyde (CH2O), but Reaction B doesn’t. The products of Reaction A are acidic (H+ is produced); those of B and C are basic (CH3COO

–Na+ and Na2CO3 are pro- duced). These things could be tested, but an even simpler method would be take advan- tage of the different amounts of carbon dioxide produced, relative to sodium bicarbonate:

In Reaction A, 1 mole NaHCO3 produces 2 moles CO2

In Reaction B, 1 mole NaHCO3 produces 1 mole CO2

In Reaction C, 2 mole NaHCO3 produces 1 mole CO2

You will measure the ratio of moles NaHCO3 used to moles of CO2 produced, and if it is approximately 1:2, you may conclude that Reaction A is correct; if the ratio is around 1:1, you can bet that Reaction B is correct, and if it’s about 2:1, you should choose Reaction C.

Keep in mind that your results may not give you exact whole number mole-to-mole ratios because of basic experimental errors. Your results may be off by as much as 20% for this experiment, but you will still be able to choose between the three reactions (A, B, or C) with a fair amount of confidence if you work carefully and collect good data.

Determining the moles of NaHCO3 is easy: Use the measured mass you scoop out of the container to use. (The other reactant, acetic acid will be used in excess, so its exact amount will have no relationship to the amount of carbon dioxide generated.)

Determining the moles of CO2 is less straightforward; it’s not so simple to collect and mea- sure the mass of a gaseous substance, as you can imagine. In each of the three reactions above, carbon dioxide is the only gas, and all other reactants and product are liquids, sol- ids, or aqueous. As the reaction occurs, carbon dioxide will bubble out of the reaction solu- tion and escape into the laboratory.

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EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

Therefore, the mass of your reaction mixture after the reaction will be lighter due to the loss of carbon dioxide, and a simple subtraction tells you how much carbon dioxide was produced:

Mass of CO2 = ⎛ ⎜ ⎝

mass of reaction mixture before reaction

⎞ ⎟ ⎠

– ⎛ ⎜ ⎝

mass of reaction mixture after reaction

⎞ ⎟ ⎠

One small complication is that some of the CO2 produced will remain dissolved in the reac- tion mixture because carbon dioxide is somewhat soluble in water. This means that the mass you calculate by subtraction in the above equation is somewhat too low—i.e., you have not accounted for the carbon dioxide that goes into the water. You will account for this with a correction factor in your calculations.

Once you have chosen the correct reaction between acetic acid and sodium bicarbonate, you can use it to measure the amount of sodium bicarbonate in Alka-Seltzer tablets using a similar methodology. In this case, the mass of sodium bicarbonate will be an unknown. You can measure the amount of CO2 produced as you did before, and use the mole-to-mole ratio of the chosen reaction to calculate the number of moles and the mass of sodium bicarbon- ate which reacted. Finally, you will determine your experimental error by comparing your experimentally determined mass of sodium bicarbonate present with what the manufac- turer reported on the package of Alka-Seltzer.

OBjECTIVES In this experiment, you will

■ Determine the stoichiometry of a reaction experimentally. ■ Weigh by difference a reaction mixture before and after the reaction in order to find the

mass of a gas produced.

■ Practice molar mass and mole ratio calculations. ■ Calculate a percent error and determine how an inaccuracy in a specific measurement

affects the outcome.

HAZARdS Even though the reagents in this lab are fairly safe, please wear safety goggles and dispose of waste in the labeled waste container.

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EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

PROCEduRE

PART A: REACTION STOICHIOMETRY

Pour a small amount of sodium bicarbonate into a small beaker to use as your personal sup- ply for the experiment. You can always get more if needed. Clean up any sodium bicarbon- ate spills with paper towels and water, and dispose of excess down the drain with water.

Run the reactions in a clean and dry 250-mL beaker, with a large watch glass as a cover. Tare a watch glass, measure about 2 g of sodium bicarbonate (NaHCO3) and record the mass in the data table. Measure 50 mL of acetic acid into a graduated cylinder and pour this into a 250-mL beaker.Stack the watch glass on top of the beaker (with the sodium bicarbon- ate still in the watch glass). Re-zero a balance and carefully measure the mass of the whole stack altogether: beaker with acetic acid, plus watch glass and sodium bicarbonate on top. Record the mass in the data table.

Without spilling any of the sodium bicarbonate or reaction mixture, return to your work area, and initiate the reaction by carefully transferring all of the sodium bicarbonate from the watch glass into the acetic acid. (Note that the solution may overflow if you dump the sodium bicarbonate too quickly.) Put the watch glass back on top of the beaker and swirl the reaction mixture gently. After a few minutes, when the reaction is complete and no more bubbling is observed, weigh and record the total mass of the beaker, reaction mix- ture, and empty watch glass.

Pour the reaction mixture down the sink with plenty of water. Rinse the glassware thor- oughly with water (no soap required) and dry completely. Never put a wet beaker onto a balance!

Repeat for at least one more trial. Do all the calculations for Part A. get approval from your instructor before you proceed.

PART B: ANALYSIS OF AN ANTACId

Using your conclusion about which of the three reactions are correct from Part A, you can now measure the amount of sodium bicarbonate contained in Alka-Seltzer tablets.

Obtain one pack of two tablets from the lab staff. This will allow you to conduct two trials of the reaction.

Use the same procedure as you used in Part A, replacing the sodium bicarbonate with the Alka-Seltzer tablet each time. Give the reaction at least 5 minutes to take place fully, and don’t worry if there are small amounts of white solid in the reaction mixture after the reac- tion is complete.

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REPORT SHEETS | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

REPORT SHEETS | Experiment 8: Stoichiometric Analysis of an Antacid Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

PART A: REACTION STOICHIOMETRY

Trial 1 Trial 2 Average

data

Mass of NaHCO3 —

Mass of Beaker + Acetic Acid + NaHCO3 + Watch Glass before Reaction

—

Mass of Beaker + Reaction Mixture + Watch Glass

after Reaction —

Calculated Results (see below for guidance)

Mass of CO2 Gas Released —

Moles of CO2 Gas Released —

Moles of CO2, Corrected for Amount Dissolved*

(see below for note)

Moles of NaHCO3 Used

NOTE: *Calculate by adding 0.0040 moles to the “Moles of CO2 Gas Released.” This correc- tion accounts for the amount of CO2 that dissolves in 50 mL of aqueous solution.

Calculate (and show your work for at least Trial #1) for the following:

1. Calculate the mass of CO2 released and then convert it to moles of CO2.

2. Calculate the number of moles of CO2, corrected, by following the special note.

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REPORT SHEETS | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

3. Using the mass of NaHCO3, calculate the number of moles of NaHCO3.

4. Compare the calculated number of average moles of CO2 and NaHCO3 in the table above by calculating a ratio. Which of the three possible reactions, A, B or C (refer to Introduction section for the options), is best supported by these results? Write out the complete balanced chemical equation you chose, and explain your reasoning.

PART B: ANALYSIS OF AN ANTACId

Tablet 1 Tablet 2

data

Mass of Beaker + Acetic Acid + Tablet + Watch Glass before Reaction

Mass of Beaker + Reaction Mixture + Watch Glass after Reaction

Calculated Results

Mass of CO2 Gas Released

Moles of CO2 Gas Released

Moles of CO2, Corrected for Amount Dissolved* (see note below)

Moles of NaHCO3** (see note below)

Mass of NaHCO3

Average Mass of NaHCO3 for 2 Tablets

*Calculate by adding 0.0040 moles to the “Moles of CO2 Gas Released.”

**Use the balanced chemical equation from Part A, Question 4, for the mole-to-mole ratio to convert moles of CO2 (corrected) to moles of NaHCO3.

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REPORT SHEETS | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

Calculate the following (and show your work for at least Trial #1):

1. Calculate the mass of CO2 gas released by the tablet and convert it to moles of CO2.

2. Calculate the number of moles of CO2, corrected, released by the tablet by following the special note*.

3. Using the mole ratio of CO2 to NaHCO3 in the balanced chemical equation that you determined in Part A, calculate the moles of NaHCO3 present in the tablet.

4. Convert the moles of NaHCO3 in the tablet to mass of NaHCO3 in the tablet.

5. Calculate for the average mass of NaHCO3 for the two tablets.

6. The mass of sodium bicarbonate in each tablet of Alka-Seltzer is reported as 1916 mg on the package (which we will call the “actual value”).

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REPORT SHEETS | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

Using your experimental value for the average mass of NaHCO3, calculate the percent error. Show your work.

% error = (actual value – experimental value)

× 100 % actual value

7. In Part B, suppose the tablet was mostly dissolved and had mostly reacted when some of your solution splashed out of the beaker.

a. How would this affect the perceived mass of CO2? It would be artificially high or artificially low? Explain.

b. Would your final calculated mass of sodium bicarbonate in the tablet be artificially high or artificially low as a result of this splashing? Explain.

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LAB EXERCISE | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

LAB EXERCISE | Experiment 8: Stoichiometric Analysis of an Antacid Name: __________________________________________________ Section: __________

1. When acetic acid and sodium bicarbonate react, why does total mass decrease?

2. Suppose you are running a trial of the reaction in Part A. You use 2.80 g of NaHCO3, and determine that 2.83 g of carbon dioxide are produced. Using these data and a periodic table, show calculations to:

a. Convert the mass of NaHCO3 to moles using the molar mass of NaHCO3.

b. Convert the mass of CO2 to moles using the molar mass of CO2.

c. Find the simplest ratio between the number of moles of NaHCO3 and CO2. (HINT: divide both numbers by the smaller of the two to get a ratio.)

The mole ratio of NaHCO3: CO2 is ____:____.

d. Compare this ratio to the ratio of the coefficients for NaHCO3 and CO2 in the three balanced chemical equations (given on page 1). Which of the three possible reac- tions discussed is consistent with these results?

Write the equation here.

NOTE: The result in part C is not necessarily the correct answer. It is a hypothetical situation for practicing the calculations required in this lab. Do not use this actual result in the lab!

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LAB EXERCISE | EXPERIMENT 8: STOICHIOMETRIC ANALYSIS OF AN ANTACId

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CHEM& 140 WORKBOOK

EXPERIMENT 9 EXPERIMENT 9: TITRATION OF VINEgAR

OBjECTIVES ■ Determine the concentration of acetic acid in vinegar (mass percent and molarity) ■ Understand how to perform a titration

INTROduCTION Acid base neutralization is a reaction between an acid and a base to produce a salt and water. An example of this type of reaction is the reaction of the strong acid HCl with the strong base NaOH to produce the salt sodium chloride and water:

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) The same result occurs when a strong acid reacts with a weak base, or when a weak acid reacts with a strong base. In each case the acid and base react to form a salt and water; neutralizing the acid and base. The neutralization reaction provides us with one method for determining the amount of either the acid or the base in a solution.

The general method employed for determining the amount of an unknown substance in a solution by using another solution with a known concentration and the reaction stoichi- ometry from the balanced chemical reaction is known as titration. Titration is a procedure for determining the amount of substance A by adding a carefully measured volume of a

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EXPERIMENT 9: TITRATION OF VINEgAR

solution B with a known concentration until the reaction between them is just complete. Titrations are not used exclusively with acids and bases but anytime you need to determine the concentration of a species in a solution.

In this lab, we will titrate a vinegar solution with a solution of sodium hydroxide, NaOH. The active ingredient in vinegar is the weak acid, acetic acid, CH3COOH or HC2H3O2. The structure is shown here (lone pairs of electrons are not shown), the hydrogen that is attached to oxygen is the acidic hydrogen, meaning it will be the only hydrogen involved is this acid-base reaction. The abbreviation for acetic acid is HAc.

The sodium hydroxide will neutralize the acetic acid, HAc, in the vinegar. The Ac actually represents the acetate anion, CH3COO

– or C2H3O2 –.

HAc (aq) + NaOH (aq) → NaAc (aq) + H2O (l ) The point at which all the acetic acid has reacted with the sodium hydroxide is called the equivalence point. Once the neutralization reaction is complete, any additional base added will produce a basic solution and the pH will greatly increase. We can determine when the titration is complete by employing an acid-base indicator that will change color in a basic solution.

Phenolphthalein is an acid-base indicator that changes color with changes in pH. This particular indicator is colorless in acidic solutions but turns pink in basic solutions. Phenolphthalein changes color at pH = 8.2. If one extra drop of base is added after the acid has been neutralized, the pH will jump significantly as the solution becomes basic and it will change to a pink color.

Colorless

O O

Pink

+ 2 OH– + 2 H2O

As soon as very faint pink color persists for 30 seconds you have reached the endpoint. In this case the equivalence point is very close to the endpoint, therefore the endpoint can be used to approximate the equivalence point.

At the equivalence point, the number of moles acid will have exactly reacted with the num- ber of moles of base according to the mole ratio in the balanced equation.

Number of moles of OH– = Number of moles of H+

H O

HO CC

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EXPERIMENT 9: TITRATION OF VINEgAR

We can determine the molarity of the acid by titrat- ing a known volume of the acid solution with a known concentration and volume of a base.

The concentration of the known solution is deter- mined through a process called standardization. For this laboratory the standardization of NaOH has already been performed for you.

INFORMATION ABOuT STANdARdIZATION (FOR YOuR FuTuRE INFORMATION)

Hydroxides are very hydroscopic (absorb water from the atmosphere) making them difficult to accurately weigh out on a balance. Hydroxide solu- tions can absorb carbon dioxide from the atmo- sphere, which can lead to changes to the concen- tration of the solution via acid-base reactions. And, although simple measuring is a good approximation of concentration, it is not precise enough for use to determine the concentration of a solution for use in a titration. Instead, the concentration of the known solution (in this case the NaOH solu- tion) can be determined through reaction with an appropriate primary standard.

LABORATORY TECHNIQuE FOR BuRETS Burets are used to deliver a recorded amount of liquid or solution to another container. A buret is marked in milliliters like a graduated cylinder, but buret markings show 0 mL at the top, and the numbers increase as you go down the buret. The stopcock controls the liquid flow. It is open when parallel to the length of the buret and closed when perpendicular to the length of the buret.

■ Washing and rinsing the buret: To clean a buret, wash its interior with soap and tap water. Next, rinse the buret with 5–10 mL portions of DI water. With the buret over the sink and the stopcock open, pour the water into the buret and let it drain out the tip. Use a beaker to pour solutions into the buret— most breakage occurs during washing, and burets do NOT fit under the faucet.

■ Conditioning the buret: After the buret is well-drained, close the stopcock and add about 5 mL of the titrant (the solution to be used into the buret). Tilt the buret sideways and roll the barrel to com- pletely rinse the inner walls of the buret. Drain the solution through the buret tip to insure the tip is also conditioned. Repeat this step at least twice to be sure all interior surfaces are rinsed with titrant.

46 47 48 49 50

42 43 44 45

38 39 40 41

34 35 36 37

30 31 32 33

26 27 28 29

22 23 24 25

18 19 20 21

14 15 16 17

10 11 12 13

7 8 9

3 4 5

2 1

Figure 1. Equipment set-up for titration

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EXPERIMENT 9: TITRATION OF VINEgAR

■ Filling the buret: Close the stopcock. Use a clean funnel to fill the buret with titrant just above the “0” mark. Place a container under the buret tip, and open the stopcock briefly to fill the buret tip with solution, leaving no air bubbles, and to get the level of meniscus to fall within the markings of the buret. If the tip does not fill with solution when the stopcock is in the open position, there may be an air bubble in the stopcock. Consult your instructor. NOTE: The initial level of titrant need not be exactly at 0.00 mL as the initial level of liquid will be recorded and subtracted from the final volume to determine the volume delivered after a titration is complete.

■ Reading the buret: Always remove the funnel used to fill the buret before taking any measurements. Record the volume of titrant by noting the bottom of the meniscus. On the buret shown below, numbers marked for every 1 mL, and the ten lines between each number represent every 0.1 mL. Thus, the level of titrant in the buret can be esti- mated to one more decimal place than the markings or to the nearest 0.01 mL. Thus, in the figure to the right, the meniscus is about halfway between 25.0 and 25.1 mL, so the level of titrant can be recorded as 25.04 mL, 25.05 mL, or 25.06 mL depending on whether the bottom of the meniscus appears to be just above, just at, or just below halfway, respectively.

■ Cleaning the buret: Afterwards, empty the buret, disposing of the titrant according to the waste disposal instructions for each experiments. Wash the buret with soap and tap water, then rinse with several portions of tap water, allowing some tap water to run through the tip. Do a final rinse with small portions of DI water, allowing the DI water to run through the tip, then return the buret to the stockroom.

SAFETY PRECAuTIONS ! CAuTION: Harmful Chemicals

NaOH is caustic. Handle with care. In case of contact with skin, rinse the area with large amounts of water and notify your instructor. Wear goggles at all times in the chemistry lab- oratory. You may wear gloves for extra protection.

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EXPERIMENT 9: TITRATION OF VINEgAR

PROCEduRE 1. Record the molarity of NaOH and obtain about 50 mL of this NaOH solution. (Make sure

your goggles are on!) You will use a small amount of NaOH (about 15 mL) for the next step, and the rest for the titration.

2. Prepare a buret for titration as described under, “Laboratory Techniques for Burets” on the previous page. Fill the buret with the sodium hydroxide solution and follow the usual procedures for eliminating air bubbles and setting the initial level.

3. Use a graduated pipet to measure out 2.00 mL of vinegar. Record the actual volume used. Transfer the vinegar to a clean Erlenmeyer flask (the 250-mL size should be large enough). Add about 25 mL of distilled water. (Does adding water change the num- ber of moles of acid in your sample? No, it just dilutes it to lower concentration. You recorded the initial volume of vinegar used and you will use that volume to calculate the concentration.)

4. Add 2–4 drops of phenolphthalein indicator to the flask.

5. Record your starting buret reading. This is your initial volume reading, V initial. (Have your lab partner verify your measurement—use the proper number of significant fig- ures! A buret has an uncertainty of ± 0.01 mL.)

6. Using the buret, add NaOH dropwise to the vinegar sample with constant swirling of the solution (analyte flask). The endpoint may be easier to see if you have a white sheet of paper under the flask. Stop when a last drop (or partial drop) of hydroxide solution results in a faint pink color that persists for at least 30 seconds.

7. Record your ending buret reading. This is your final volume reading, V final. Calculate the volume of NaOH delivered (added to flask) by finding the difference as V final – V initial.

8. Pour the contents of the flask into the chemical waste container. Rinse the flask with tap water, then a couple of times with distilled water. Use the volume of NaOH deliv- ered for the first trial to estimate if you have enough NaOH in your buret for two more trials; add more NaOH to the buret if necessary. Repeat the titration twice more with two new samples of vinegar by repeating Steps 3 through 7. The volumes delivered for each of the trials in a titration need to be within ± 0.20 mL.

CHEMICAL WASTE Everything should go into the appropriate labeled waste container in the fume hood.

MATERIALS ANd EQuIPMENT Ceramic stand, buret, buret clamp, 250-mL Erlenmeyer flask, 5- or 10 mL-graduated pipet, green pipet pump, 25-mL graduated cylinder, a large beaker to hold waste and rinse solu- tions, two small beakers to hold NaOH and vinegar.

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EXPERIMENT 9: TITRATION OF VINEgAR

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REPORT SHEETS | EXPERIMENT 9: TITRATION OF VINEgAR

REPORT SHEETS | Experiment 9: Titration of Vinegar Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

dATA NOTE: The volumes delivered for each of the trials in a titration should be within ± 0.20 mL.

Record your data using appropriate number of significant figures and units.

Table 1. Acid-Base Titration

Trial 1 Trial 2 Trial 3

Volume of Vinegar (mL)

Molarity of NaOH

Initial Volume (mL) of NaOH in Buret

Final Volume (mL) of NaOH in Buret

Volume of NaOH Delivered (mL)

CALCuLATIONS Show all of your work using proper significant figures and units to calculate the results.

(For all of the calculations that follow, show a sample calculation for at least one trial.)

1. Calculate the volume of NaOH (in liters) used to neutralize the acetic acid.

2. Using the molarity and volume of NaOH solution (in liters) used, calculate the moles of NaOH used to neutralize the acetic acid.

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REPORT SHEETS | EXPERIMENT 9: TITRATION OF VINEgAR

3. Using the balanced chemical equation for the reaction of NaOH with CH3COOH, calcu- late the moles of acetic acid in vinegar.

4. Using the moles of CH3COOH present and initial volume of vinegar pipetted, calculate the molarity of acetic acid in vinegar.

5. Using the molarity of acetic acid in vinegar for each trial, calculate the average molar- ity of acetic acid in vinegar.

6. Using the moles of acetic acid and molar mass of acetic acid, calculate the mass of acetic acid used in each titration.

7. Using the initial volume of vinegar pipetted and the known density of vinegar 1.005 g/mL, calculate the mass of vinegar solution.

8. Using the mass of acetic acid and the mass of vinegar solution, calculate the (m/m) % of acetic acid in vinegar.

9. Calculate the average (m/m) % of acetic acid in vinegar.

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REPORT SHEETS | EXPERIMENT 9: TITRATION OF VINEgAR

Tabulate your results here:

Table 2. Acid-Base Titration

Trial 1 Trial 2 Trial 3 Average

Volume of 0.10 M NaOH Delivered (L) —

Moles of 0.10 M NaOH Used —

Moles of Acetic Acid in Vinegar —

Molarity of Acetic Acid

(m/m) % of Acetic Acid

POST-LAB QuESTIONS 1. Most commercial vinegar solutions are 5.0 % (m/m) acetic acid.

a. Calculate the percent error for your average % (m/m) acetic acid compared to a commercial solution of 5.0 % (m/m) acetic acid.

b. For this lab, percent errors greater than 10% indicate that the experiment was not very accurate. How accurate were your results? Include your calculated percent error to support your answer.

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REPORT SHEETS | EXPERIMENT 9: TITRATION OF VINEgAR

225

LAB EXERCISE | EXPERIMENT 9: TITRATION OF VINEgAR

LAB EXERCISE | Experiment 9: Titration of Vinegar Name: __________________________________________________ Section: __________

1. What is the formula for calculating density?

2. What is the formula for calculating the % (m/m) of a solution?

3. What is the formula for calculating the molarity of a solution?

4. If you were given 250.0 mL of 0.500 M NaCl, which has a density of 1.003 g/mL, what would be the % (m/m) of the NaCl solution. Show your work for credit. Report your answer to an appropriate number of significant figures and include units.

a. Using volume and molarity of NaCl, calculate the moles of NaCl present.

b. Convert moles of NaCl to grams of NaCl.

c. Calculate the mass of the 250.0 mL of solution using its density 1.003 g/mL.

d. Calculate the % (m/m) of the NaCl solution.

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LAB EXERCISE | EXPERIMENT 9: TITRATION OF VINEgAR

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CHEM& 140 WORKBOOK

EXPERIMENT 10 EXPERIMENT 10: SPECTROSCOPY

OBjECTIVES ■ To identify the unknown element by comparing its flame color to that of known elements. ■ To learn about light and the relationship between energy, frequency, and wavelength. ■ Observe continuous vs. line spectra resulting from emission from gas discharge tubes

and identify the element in fluorescent light bulbs.

INTROduCTION Light is a form of energy called electromagnetic radiation. Light travels at a speed of 2.998 × 108 m/s, a constant value in a vacuum, and it is denoted by the lowercase letter “c.” A chart of the electromagnetic spectrum is shown below.

Radiowave Microwave

Longer Wavelength Lower Frequency Lower Energy

Shorter Wavelength Higher Frequency

Higher Energy

Infrared Visible Ultraviolet X-Ray Gamma Ray

Figure 1. Types of electromagnetic radiation

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EXPERIMENT 10: SPECTROSCOPY

Which type of electromagnetic radiation is more dangerous to humans, radiowaves, or x-rays? _________________

Which type of electromagnetic radiation has the highest energy? _________________

Which type of electromagnetic radiation is associated with the sun? _________________

What is the relationship of the frequency of a wave to the energy of the electromagnetic radiation? _________________

One model used to describe electromagnetic radiation is called the wave model. This model uses several variables to describe this radiation: wavelength, frequency and energy.

Wavelength is denoted by the symbol lambda (λ) and is the distance between the crests (peaks) or troughs or one complete wave cycle of a wave. Wavelength can be measured in any length unit. For example, radio waves are often measured in meters. Gamma rays are often measured in nanometers.

Another unit for wavelength is the micrometer (10–6 meters, or called microns). It’s often used in biology since it is the approximate size of cells and bacteria. You will encounter microns when using most microscopes.

The second variable used to describe waves is called frequency (ν), which is the number of cycles which pass per unit time (see Figure 2 above for “cycle”). The unit of frequency is the hertz (Hz). Once per second is 1 hertz, twice per second is 2 hertz, etc. One thousand hertz is expressed as a kilohertz (kHz). One million hertz is a megahertz (MHz).

Frequency is denoted by the symbol nu (ν). Frequency and wavelength of light are related to the speed of light by this equation:

c = λν

Shorter ↓ wavelength = higher ↑ frequency. This is an inverse relationship. (See the chart above to confirm this.)

The energy carried by light is related to both wavelength and frequency. X-rays and ultra- violet rays are high energy radiation and can be damaging to your cells. On the other hand, radio waves are lower in energy and you need not use “blocking” agents to protect yourself!

Crest

Trough

Cycle

Figure 2. Terminology of a wave.

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EXPERIMENT 10: SPECTROSCOPY

The equation that relates energy and frequency is:

E = hν

where h is Planck’s constant with a value of 6.626 × 10–34 J ∙ s

Higher ↑ frequency = Higher ↑ energy. This is a direct relationship. (See the chart above to confirm this.)

A second model is often used to describe the energy carried by light. This model describes light as having particle characteristics. These “particles” are bundles of energy called pho- tons. Each photon carries a small bundle of energy. A single photon associated with ultra- violet light has greater energy than a single photon associated with radio waves. Using this model, we can think of light as a stream of photons.

Let’s compare radiowaves (a) with ultraviolet radiation (b):

Which type of radiation is higher energy?

Looking at radiowaves, when you tune your radio you are receiving radio waves with fre- quencies between 530 and 1700 kHz for AM and 87.5 to 108 MHz for FM. (This corresponds to the “numbers” on your dial.) A radio wave cor- responding to 1700 kilohertz on your AM radio has a wavelength of approximately 200 meters, approximately twice the length of a foot- ball field. These waves pass through us with- out harm.

Sunscreens protect your skin from uV radiation by absorbing wavelengths between 290 and 400 nanometers. This is a short wavelength compared to radio waves. The frequency of UV light compared to radio waves is extremely high. UV light have frequencies in excess of 700 trillion (7 × 1014) hertz. Thus, photons of UV light have a higher energy than photons of radiowaves.

Both radio waves and ultraviolet light are invisible to the human eye. Colors are due to vis- ible light, which have wavelengths between radiowaves and UV, with a range between 400–700 nanometers. Each color has a corresponding wavelength. For example, red light has a wavelength in the 700 nm range, while violet light has a wavelength near 400 nano- meters. The visible spectrum (the “rainbow”) can be remembered with a mnemonic ROY G BIV (Red, Orange, Yellow, Green, Blue, Indigo, V iolet). An interesting bit of trivia—the British use the following mnemonic—R ichard O f York Gave Battle In Vain.

(a)

(b)

Figure 3. Comparing two waves corresponding to different types of electromagnetic radiation.

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EXPERIMENT 10: SPECTROSCOPY

ORIgIN OF SPECTRAL LINES The origin of the spectral lines can be explained using the planetary model of the atom, in which the nucleus is in the center and surrounded by orbits (n) on which electrons travel (NOTE: this is not completely accurate—this is an oversimplified model). Atoms are usu- ally in a low energy state, which means their electrons are in the lowest energy orbits such as level 1 (n = 1). This is called the ground state for an atom. If energy is added to an atom, the electrons can absorb energy and move to a higher energy state, called an excited state (in Figure 4a, the electron jumps from level 1 to the second level, n = 2). This energy can be added to atoms by an electric discharge or by heat, an incoming photon. This added energy is emitted when the atom gives off a photon and the excited electron returns to the ground state (see Figure 4b). Figure 5 shows electron transitions focusing on energy levels of electrons.

(b)(a)

n = 1

n = 2

Photon

n = 1

n = 2

Photon

Figure 4. (a) An atom in its ground state, with the electron (green dot) moving from level 1 to level 2 by absorbing energy from an incoming photon. (b) An atom in an excited state emitting a photon as its electron moves to a lower energy level. Only two levels (n = 1, n = 2) are shown for simplicity.

n = 3

n = 2 n = 2 ∆E for

n = 3 → 1

n = 1

n = 3

n = 2

n = 1 e–

n = 3 e–

n = 1 e–

Energy is added. Electron is excited.

Electron relaxes. Light is emitted.

Figure 5. Electron transitions in an atom from the perspective of energy levels.

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EXPERIMENT 10: SPECTROSCOPY

The light emitted by atoms have a definite wavelength and if visible, a color that depends on the amount of energy originally absorbed. Each excited electron will emit one photon of light as the electron moves to a lower energy state. Billions of excitations and emissions are possible because a typical sample is made up of billions of atoms. There are also many possible energy jumps for each electron. Each of these “jumps,” electron transitions (∆E), matches a different color of light. These colors constitute the emission spectrum of the excited atoms. A large jump corresponds to a color associated with high energy photons. A small jump corresponds to a color associated with a low energy photon.

Each element has a unique set of spectral lines. Since each element has a unique set of spectral lines, spectral studies can help us identify an element. The line spectrum of ele- ments is much like a “fingerprint” for the element.

SAFETY PRECAuTIONS Use caution with the Bunsen burners and wooden splints. Wear goggles and tie hair/loose clothing back to prevent fires. Do not touch the connections when the power supply is plugged into an outlet. A serious electric shock can result. Also, these bulbs get very hot. Allow them to cool before touching. Turn off the power supply when you are not using it. Place chemical waste in the labeled waste container.

PROCEduRE

PART 1: FLAME TESTS (ALL OF THE SAMPLES ARE FINELY gROuNd POWdERS)

1. Light a Bunsen burner. Adjust the air vent to obtain a blue flame.

2. Obtain a vial containing a known sample (note the element symbol on the vial) and a piece of weigh paper. Using a spatula, take a scoop of solid from the vial onto the weigh paper. Dip a wooden splint it into one of the beakers of water to moisten it. Gently roll the splint onto the weigh paper with the solid so the tip of the splint is covered with sample.

3. Place the splint into the clear blue flame and observe/record the color of the flame. Dip the splint into the other beaker of water to douse the flame.

4. Repeat this procedure for each known.

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EXPERIMENT 10: SPECTROSCOPY

PART 2: IdENTIFICATION OF AN uNKNOWN SOLId

Repeat the above procedure for two unknowns. Check your results with your instructor before you proceed.

PART 3: CONTINuOuS ANd LINE SPECTRA

Using a spectroscope, examine the spectrum given off by an incandescent light bulb and a fluorescent light bulb. Sketch the spectrum for each labeling your diagram with colors and wavelengths.

PART 4: OBSERVINg LINE SPECTRA WITH THE SPECTROSCOPE

1. Your instructor has set up several gas discharge tubes in power supplies. Use your spectroscope to examine the spectrum for each of the tubes. (See Hazards note.)

2. Sketch the results on your data sheet, noting the color and wavelength of the observed lines. Look at tubes containing He, Ne, Hg and H2 (and others).

3. Use your spectroscope to re-examine the fluorescent light bulbs in the lab room.

233

REPORT SHEETS | EXPERIMENT 10: SPECTROSCOPY

REPORT SHEETS | Experiment 10: Spectroscopy Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

PART 1: FLAME TESTS FOR KNOWN ELEMENTS Write the color of the flame next to each element.

Lithium ___________________________ Calcium __________________________________

Sodium ___________________________ Strontium ________________________________

Copper ___________________________ Potassium ________________________________

Iron ______________________________ Barium ___________________________________

PART 2: IdENTIFICATION OF SOLId uNKNOWNS Solid unknown # __________ Color of flame ___________ Element ___________________

Solid unknown # __________ Color of flame ___________ Element ___________________

PART 3: CONTINuOuS ANd LINE SPECTRA (Draw what you see through the spectroscope).

Incandescent light bulb:

Fluorescent light bulb:

234

REPORT SHEETS | EXPERIMENT 10: SPECTROSCOPY

PART 4: gAS-dISCHARgE TuBES Draw what you see through the spectroscope when you examine each of the following gas discharge tubes. Include colors and wavelengths. Also, describe the color of the glowing discharge tube as seen by naked eye (without the spectroscope).

Helium

Color of the glowing discharge tube: ____________________

Helium spectrum

Neon

Color of the glowing discharge tube: ____________________

Neon spectrum

Mercury

Color of the glowing discharge tube: ____________________

Mercury spectrum

Hydrogen

Color of the glowing discharge tube: ____________________

Hydrogen spectrum

235

REPORT SHEETS | EXPERIMENT 10: SPECTROSCOPY

POST-LAB QuESTIONS 1. What are some differences in the emission spectra of the different elements?

2. In which discharge tubes are electrons undergoing the largest and smallest energy changes (from what we can tell by only observing gas tubes with the naked eye)?

3. a. Which discharge tube emits visible light with the highest frequency, as seen with the naked eye?

b. Which discharge tube emits visible light with the longest wavelength, as seen with the naked eye?

4. What are some differences in emission between the incandescent and fluorescent light bulb?

5. Identify the light sources that produce a line spectrum. Identify those that produce a continuous spectrum.

236

REPORT SHEETS | EXPERIMENT 10: SPECTROSCOPY

6. Which element in the glowing discharge tubes is most likely to be the element in the flu- orescent lamp? Explain your reasoning.

7. To answer a–d, consider the energy level diagram to the right.

a. Draw arrows on the diagram showing all of the transitions an electron could make that would result in the emission of a photon.

b. How many emission lines could we observe from this atom?

c. Which transition would emit light with the highest frequency (shortest wavelength)?

From n = ______ to n = ______

d. Which transition would emit light with the lowest frequency (longest wavelength)?

From n = ______ to n = ______

4 3

237

LAB EXERCISE | EXPERIMENT 10: SPECTROSCOPY

LAB EXERCISE | Experiment 10: Spectroscopy Name: __________________________________________________ Section: __________

Complete the following questions BEFORE class. Refer to your textbook or the web if needed.

1. a. List the colors in the visible spectrum, from highest to the lowest frequency.

b. Which color of visible light has the lowest energy? ________________________________

The longest wavelength? _________________________________________________________

The lowest frequency? ___________________________________________________________

2. The following diagrams of waves are labeled λ1, λ2, and λ3. Fill in each blank with the best choice.

λ3

λ2λ1

a. Which wave has the longest wavelength? _____ The shortest wavelength? _____

b. Which wave has the greatest frequency? _____ The smallest frequency? _____

c. Which wave would be a good model for UV light? _____ Radio waves? _____ Visible light? _____

d. Which wave has photons with the greatest energy? _____ The smallest energy? _____

3. What is the difference between a continuous spectrum versus a line spectrum?

238

LAB EXERCISE | EXPERIMENT 10: SPECTROSCOPY

239

EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

CHEM& 140 WORKBOOK

EXPERIMENT 11 EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

PART A The first set of molecules we will examine contain only two atoms. These are considered to have a linear molecular shape.

PART B 1. The molecules in this part contain more than 2 atoms. Draw the Lewis structure for

each of them.

2. Make a model of the molecule. Use the short gray plastic connecting rods for single bonds and lone pairs of electrons.

■ Use the longer flexible connecting rods for double and triple bonds (2 for a double bond; 3 for a triple bond).

■ Use the following colored spheres: ■ C, Si = black ■ N, P, As = blue

■ H = white ■ O, S = red

■ Cl, Br, I = purple or orange

■ If the “red” sphere is the central atom, you won’t be able to see the lone pairs of electrons; use a black sphere in these cases.

240

EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

3. The electron geometry is determined by the number of electron groups on the central atom. (HINT: Look at the arrangement of the connecting rods.)

■ 4 groups result in a tetrahedral arrangement of the electron groups. ■ 3 groups result in a trigonal planar arrangement of the electron groups. ■ 2 groups result in a linear arrangement of the electron groups.

4. Determine the molecular shape of the molecule. While the lone pairs of electrons influ- ence the molecular shape, the arrangement of the atoms is used to describe the molec- ular shape. Do not use any abbreviations in your answers!

POSSIBLE SHAPES FOR MOLECuLES WITH 4 ELECTRON GROUPS AROuNd THE CENTRAL ATOM (BONd ANgLES NEAR 109.5°)

Molecular Shape 3-d Sketch and Lewis Structure

Tetrahedral

A four sided figure. Each side is an equilateral triangle. H

H H

Trigonal pyramidal

A four sided figure. Three sides are isosceles triangles. They are sitting

on an equilateral triangular base.

H H

H N

H H

H N

Bent

A nonlinear arrangement of atoms. H H O

H O

POSSIBLE SHAPES FOR MOLECuLES WITH 3 ELECTRON GROUPS AROuNd THE CENTRAL ATOM (BONd ANgLES NEAR 120°)

Molecular Shape 3-d Sketch and Lewis Structure

Trigonal Planar

A flat molecule. The atoms bonded to the central atom form an equilateral triangle. The central atom sits in the center of the triangle. F F

B F FBF

Bent

A nonlinear arrangement of atoms. F O NF

241

REPORT SHEETS | EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

REPORT SHEETS | Experiment 11: Molecular Model Building Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

PART A

Species/ Total v.e.

draw the Lewis dot Structure Name of Molecular Shape

Name of Element or Compound

H2 2 v.e.

Br2

HBr

CONCLuSIONS

If only two atoms are bonded, the molecular shape will always be _______________________.

242

REPORT SHEETS | EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

PART II Sp

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243

REPORT SHEETS | EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

Sp ec

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This page contains polyatomic ions. They have gained or lost valence electrons to become charged.

*Nitrite (NO2 –) and carbonate (CO3

2–) have resonance structures—different but equiva- lent Lewis structures.

244

REPORT SHEETS | EXPERIMENT 11: MOLECuLAR MOdEL BuILdINg

245

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

CHEM& 140 WORKBOOK

EXPERIMENT 12 EXPERIMENT 12: SEPARATION OF A

TERNARY MIXTuRE

SuggESTEd PRE-LAB ACTIVITY This lab has many parts to it. It is advised to sketch out the procedure to this lab as a pre- lab activity to see how the components of this mixture are separated. A flow chart or a schematic diagram showing the steps in picture format may be helpful.

OBjECTIVES ■ Understand the differences between physical and chemical separations ■ Identify pure substances and mixtures ■ Distinguish homogenous mixtures from heterogeneous mixtures

INTROduCTION Mixtures are physical combinations of two or more different substances where each of the substances are mixed but do not combine with each other chemically. For example, in muddy water (a mixture of dirt and water) both the dirt and the water maintain their chemical identity. Mixtures can be classified as either homogeneous or heterogeneous. Homogeneous mixtures have the same composition throughout the mixture and appear uniform. Homogeneous mixtures are often solutions, a situation where one substance (the solute) is dissolved in another (the solvent). Dissolving is a physical process and does

246

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

not change the substance chemically. Vinegar and household ammonia are examples of homogeneous mixtures. In contrast, heterogeneous mixtures have varying composition throughout the mixture and do not appear uniform. Chocolate chip cookie dough, seashells on a beach and chicken noodle soup are all examples of heterogeneous mixtures.

Chemists often need to separate a mixture into its pure components. Chemical and physi- cal properties can be used to separate some mixtures. Solubility is an example of a physi- cal property used to separate mixtures. When water is added to a mixture of salt and sand, the salt dissolves but the sand does not. The mixture can then be separated by filtration. The dissolved salt passes through the filter paper with the water (the liquid or soluble mix- ture that passes through the filter paper is called filtrate) but the insoluble sand remains behind on the filter paper (the solid or insoluble substance left on filter paper is called res- idue). Neither the salt nor the sand has been chemically changed; the separation is a phys- ical process.

In this laboratory, you will be using both physical and chemical properties to separate the components of a ternary (three substances) mixture. You will also determine the mass per- centage of each substance in the mixture.

When distilled water is added to a mixture of NaCl (sodium chloride), CaCO3 (calcium car- bonate) and SiO2 (silicon dioxide = sand), the water dissolves the NaCl (a physical change), leaving the insoluble CaCO3 and SiO2 behind. Filtration separates the solids (residue) from the NaCl aqueous solution (filtrate).

Beaker, NaCl, CaCO3, SiO2 → Beaker, NaCl, CaCO3, SiO2 Water is removed from the NaCl aqueous solution (filtrate) by evaporation. The mass lost after the filtrate is dried is the mass of NaCl in the mixture. From the residue (containing CaCO3 and SiO2) left on the filter paper, CaCO3 is removed by adding HCl (hydrochloric acid) through a chemical change).

Evaporating dish, CaCO3, SiO2 → Evaporating dish, SiO2 The HCl reacts with the CaCO3 to produce water, carbon dioxide (a gas) and CaCl2 (which is dissolved in the water). The CaCO3 is chemically converted into two new substances CO2 and CaCl2 according to this chemical equation

CaCO3 (s) + 2 HCl (aq) → CO2 (g) + CaCl2 (aq) + H2O (l ) The CaCl2 solution is removed by decantation, which means removing the aqueous layer by carefully pouring it off. The mass lost after the residue has been dried is the mass of the CaCO3 that was in the mixture before its chemical transformation. Since the sand does not dissolve in water or react with the acid, its mass can be determined by weighing the mate- rial left in the evaporating dish. The mass of the sand can then be calculated by subtracting

water

acid

1 The chemical reaction states that 1 unit of solid CaCO3 combines with 2 units of HCl acid in water solution to produce 1 unit of CO2 gas, 1 unit of CaCl2 in solution and 1 unit of water. Note that this representation conserves the number and kind of atoms found on either side of this chemical equation.

247

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

the mass of the evaporating dish from the mass of the evaporating dish and sand combined. Because the unknown contained only CaCO3, SiO2 and NaCl, the mass of CaCO3 can be determined by subtracting the isolated masses of both the sand and NaCl from the mass of the total unknown.

The mass percent of each component is determined by dividing each of the isolated com- ponents (SiO2, NaCl and CaCO3) by the mass of the initial mixture and multiplying by 100%:

% component = mass of the component

× 100% mass of the mixture

PROCEduRE 1. Obtain a 150-mL beaker.

2. Obtain an unknown mixture and accurately weigh (±0.001 g) out between 2.5 g and 3.5 g into the beaker. Record the actual mass used in your data table.

3. Add 10 mL distilled water to the mixture and stir. This should dissolve the NaCl but leave the CaCO3 and SiO2 intact.

4. Prepare a vacuum flask with a Büchner fun- nel (see diagram). Obtain a 7.5-cm circular fil- ter paper that fits inside. Connect the flask to the vacuum line using a rubber hose. Clamp the flask to a ring stand so it does not fall over.

5. Carefully filter the mixture by slowly pouring through the filter paper with the vacuum. Be careful not to lose any of the filtrate (NaCl (aq), the salt water solution).

6. Rinse the original beaker with approx. 5 mL of water and pour over the filter paper (repeat two times). This should ensure that all of the remaining residue (CaCO3 and SiO2) is on the filter paper and that all of the NaCl is in the liquid portion (filtrate). Retain the residue for use in Step 11.

7. Weigh a clean, dry 250-mL beaker and a clean, dry watch glass and record the actual masses on your Report Sheet.

8. Pour the filtrate (the liquid in your suction flask) into the beaker. Be careful not to lose any solution. Use a small amount of water from your distilled water bottle to rinse the flask contents into the beaker.

1000 mL

850

600

450

200

Büchner Funnel

Filter Paper

Vacuum Flask

To Vacuum Source

248

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

9. Place the beaker containing the solution of NaCl on a hot plate and gently boil to remove the water. Once the volume has been reduced, you will want to place a watch glass over the beaker to prevent splattering (convex side down). Once NaCl is fully dried, wait for the beaker and watch glass to cool down to room temperature and then record their mass.

10. Weigh (±0.001 g) a clean, dry evaporating dish and record the mass on your data sheet.

11. Using forceps (tweezers), gently transfer the filter paper with residue(CaCO3 and SiO2) from step 6 to the evaporating dish. Be careful not to lose any residue.

12. Obtain approximately 15 mL of 2.0 M HCl in a graduated cylinder. Be careful: HCl is a strong acid. Add HCl dropwise to the residue and filter paper until you no longer observe the evolution of gas bubbles. The HCl will react with the CaCO3 but not the SiO2. You will probably not need to use all of the HCl. (Excess goes in the waste bea- ker.) When the reaction is completed, you will have CaCl2 solution and SiO2 in the evap- orating dish.

13. Use a 250-mL beaker as a waste beaker. Carefully decant the solution in to your waste beaker, making sure not to lose any of the solid SiO2.

14. Using forceps (tweezers), pick up the filter paper and gently rinse it with distilled water from the wash bottle retaining the liquid in the evaporating dish.

15. Discard the filter paper in the appropriate waste container.

16. decant (carefully pour off) the rinse water from the SiO2 in your waste beaker.

17. Rinse the SiO2 a third time and decant the solution into the waste beaker making sure not to lose any of the SiO2.

18. Place the evaporating dish onto a clay triangle and/or iron ring clamped on your ring stand and gently heat the dish carefully over a cool (low) flame for 10 minutes until the SiO2 is dry. If the SiO2 does not appear dry, keep heating. It is very important to use a cool flame to ensure that the SiO2 does not spatter.

19. Allow the evaporating dish to cool, then weigh the evapo- rating dish and SiO2.

20. Heat the evaporating dish and SiO2 again for two minutes, let it cool, and reweigh it. If the masses differ by more than 0.002 g, heat the dish for another two minutes followed by cooling, and weighing. Remember, the last weighing is the number you will use in your calculations.

249

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

WASTE dISPOSAL Place all waste for this experiment in a labeled waste container.

SAFETY You must wear safety goggles at all times during this experiment.

It is recommended you wear gloves when handling acids: 2.0 M HCl can cause chemi- cal burns. If you spill any on you, wash it immediately with plenty of water and notify your instructor.

Several heating steps are involved. Handle the hot glassware and evaporating dishes with tongs or oven mitts. Make sure to let the hot beaker and hot evaporating dish cool to room temperature before weighing.

MATERIALS ■ Beakers (150 mL, 250 mL) ■ Watch glass ■ Ring stand ■ Büchner funnel ■ Filtration flask ■ Vacuum hose ■ Filter paper (7.5 cm)

■ Clamps ■ Evaporating dish ■ Hot plate ■ Tweezers ■ Pasteur pipets with rubber bulb ■ Bunsen burner and striker ■ Gloves

250

EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

251

REPORT SHEETS | EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

REPORT SHEETS | Experiment 12: Separation of a Ternary Mixture Name: __________________________________________________ Section: __________

Lab Partner: ________________________________________________________________

Show calculations in the provided space. Always include units and use the proper num- ber of sig figs!

1. Mass of unknown mixture ______________

2. Mass of a 150-mL beaker and watch glass ______________

3. Mass of dried NaCl + beaker + watch glass ______________

4. Mass of dried NaCl ______________

5. Percent sodium chloride in unknown mixture ______________

6. Mass of clean dry evaporating dish ______________

7. Mass of evaporating dish plus dried SiO2 (after HCl reaction—after each heating)

1st mass: ____________ 2nd mass:___________ 3rd mass (see Step 20): __________

8. Mass of SiO2 ______________

9. Percent silicon dioxide in unknown mixture ______________

10. Mass of CaCO3 ______________

(Remember: you must subtract the mass of the SiO2 and NaCl from the mass of the mixture)

11. Percent calcium carbonate in unknown mixture ______________

12. Fill in the table below using the data that you have obtained

% NaCl % SiO2 %CaCO3 Total

(Make sure total % adds up to 100%, or nearly so.)

252

REPORT SHEETS | EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

POST-LAB QuESTIONS 1. When water was added to the unknown in Step 3 did you observe a physical or a chem-

ical change? Explain.

2. When decanting and rinsing with water in Steps 13–17, a student accidentally poured off some of the fine particles of SiO2. How would the student’s percents of NaCl, CaCO3, and SiO2 be affected (i.e., too high, too low, or unchanged)? For each substance, circle your answer and explain.

% NaCl: increase / decrease / unchanged

% CaCO3: increase / decrease / unchanged

% SiO2: increase / decrease / unchanged

3. In an experiment, a student weighed out 2.510 grams of unknown but recorded the mass as 5.210 grams in the student’s laboratory notebook. How would the percents of NaCl, CaCO3, and SiO2 be affected (i.e., too high, too low, or unchanged)? For each sub- stance, circle your answer and explain.

% NaCl: increase / decrease / unchanged

% CaCO3: increase / decrease / unchanged

% SiO2: increase / decrease / unchanged.

253

REPORT SHEETS | EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

4. The mass of an evaporating dish with a mixture of sand, NaCl, and CaCO3 is 26.817 g. The mass of the evaporating dish is 22.437 g. Water is added to the mixture, dissolv- ing the salt, and then the liquid is poured off leaving a solid. After the solid is dried, the mass of the remaining solid mixture and the dish is 25.332 g. When 50 mL of 1.1 M HCl is added to the mixture, it fizzes, chemically changing CaCO3 to CaCl2 and CO2 gas. After this residue is dried, the mass of the residue and the dish is 24.007 g. Calculate the % of sand, % of NaCl, and the % of CaCO3 in the mixture. Show all calculations and report your answer to an appropriate number of sig. figs.

254

255

LAB EXERCISE EXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

LAB EXERCISE Experiment 12: Separation of a Ternary Mixture Name: __________________________________________________ Section: __________

1. Identify the following changes as chemical or physical:

a. The steak dinner you ate is being digested and used for energy. ________

b. Iron rusts when left in a moist environment. ________

c. Dry ice (solid carbon dioxide) sublimes to form carbon dioxide gas. ________

d. A student drops a 150-mL beaker that breaks into 17 pieces. ________

e. A battery terminal corrodes when battery acid is spilled on the terminal. ________

2. Name the physical methods that are used in this experiment to separate the mixture. Identify the procedure steps where each is used. (Example: dissolving in Step 3.)

3. Identify the step of this experiment that is a chemical method used to separate the mix- ture. Write the chemical equation for the reaction that represents this chemical change.

256

LAB EXERCISEEXPERIMENT 12: SEPARATION OF A TERNARY MIXTuRE

257

NOTES

Notes

258

NOTES

Notes

259

NOTES

Notes

260

NOTES

Notes

261

NOTES

Notes

262

NOTES

Notes

263

NOTES

Notes

264

POLYATOMIC IONS

ammonium NH4 + sulfite SO3

2–

hydronium H3O + hydrogen sulfate (bisulfate) HSO4

–

hydroxide OH– hydrogen sulfite (bisulfite) HSO3 –

carbonate CO3 2– perchlorate ClO4

–

hydrogen carbonate (bicarbonate) HCO3 – chlorate ClO3

–

nitrate NO3 – chlorite ClO2

–

nitrite NO2 – hypochlorite ClO–

phosphate PO4 3– cyanide CN–

hydrogen phosphate HPO4 2– permanganate MnO4

–

dihydrogen phosphate H2PO4 – thiosulfate S2O3

2–

chromate CrO4 2– oxalate C2O4

2–

dichromate Cr2O7 2– borate BO3

3–

acetate C2H3O2 – or

CH3COO – citrate

C6H5O7 3– or

C3H5O(COO –)3

sulfate SO4 2–

NuMBERS ANd FORMuLAS TO KNOW

Avogadro’s Number: 6.022 × 1023

mass percent (% m/m) =

mass of solute × 100

mass of solution

volume percent (% v/v) = volume of solute

× 100 volume of solution

mass volume percent (% m/v) =

mass of solute (g) × 100

volume of solution (mL)

molarity (M) = moles of solute

liters of solution

T K = T°C + 273

T°F = (1.8 × T°C) + 32

C 1V 1 = C 2V 2