Pre- and post lab questions for Spectroscopy Lab
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Lab_Manual_CHM-241_Ed-2.pdf
Bergen Community College Physical Sciences Department
General Chemistry II CHM 241
LABORATORY MANUAL 2019
Second Edition
Dr. Ara Kahyaoglu Prof. Jean Acken Associate Professor Assistant Professor
BERGEN COMMUNITY COLLEGE (BCC) is committed to providing quality material that promotes the best in inquiry-based
science education. However, conditions of actual use may vary, and the safety procedures and practices described in this resource
are intended to serve only as a guide. Additional precautionary measures may be required. BCC and the authors do not warrant or
represent that the procedures and practices in this resource meet any safety code or standard of federal, state, or local regulations.
BCC and the authors disclaim any liability for personal inquiry or demand to property arising out of or relating to the use of this
resource, to include any of the recommendations, instructions, or materials contained therein.
Bergen Community College
400 Paramus Road, Paramus, NJ 0765
Bergen Community College 2 General Chemistry II Laboratory
PREFACE
To the instructor,
This manual is available as a free download from the Bergen Community College website.
The Science Department’s aim has been to provide low cost, safe, and interesting, yet
relevant experiments.
Students are to complete the pre-laboratory exercises prior to the laboratory session. The
departmental format for the laboratory report can be found in the course syllabus.
Much effort has been made by the chemistry faculty to review this manual and make it as
error-free and accurate as possible. However, some errors will have escaped our notice.
Your help in forwarding to us any errors, inaccuracies, and/or suggestions will be greatly
appreciated. We will definitely welcome your comments and suggestions. We will make the
changes and improvements as soon as we can to make the updated manual ready for
succeeding semesters. We can be contacted at [email protected] and
Good luck and have a good semester. We look forward to hearing from you.
To the students,
This manual is available to BCC students for free download from the Bergen Community
College website. The Science Department’s aim has been to provide low cost, safe, and
interesting, yet relevant experiments that illustrate the concepts presented in the lecture
course.
Safety is everyone’s number one priority. Do not hesitate to ask your instructor if you do not
understand the procedure.
Keep in mind that your instructor expects you to be prepared for every laboratory session.
We appreciate the efforts of professors PJ Ricatto, Linda Box, Gary Porter, Brent Chapman,
Frank Ramdayal, Farah Rezae and Riwa Dandan through their suggestions and corrections.
Dr. Ara Kahyaoglu, author
Prof. Jean Acken, contributor and editor
Bergen Community College 3 General Chemistry II Laboratory
TABLE OF CONTENTS
Page
Course Schedule 4
Laboratory Safety 5
Integrity of Data Guidelines 8
Experiment 1. Heat of Fusion 9
Experiment 2. Intermolecular Forces 14
Experiment 3. Spectroscopy 23
Experiment 4. Percent Copper in Brass 31
Experiment 5. Freezing Point Depression 39
Experiment 6. Chemical Kinetics 50
Experiment 7. Le Châtelier's Principle 63
Experiment 8. Coordination Number 75
Experiment 9. Identification of a Weak Acid 83
Experiment 10. Solubility Product 94
Experiment 11. Qualitative Analysis of Cations 101
Experiment 12. Titration of Hydrogen Peroxide 109
Experiment 13. Electrochemistry 119
Experiment 14. Bicarbonate-carbonate mixture 128
Appendix A. Common Laboratory Equipment 135
Appendix B. Volumetric Glassware 136
Appendix C. Graphing 137
Appendix D. Titration 139
Appendix E. Filtration 140
Appendix F. Periodic Table 141
Bergen Community College 4 General Chemistry II Laboratory
COURSE SCHEDULE
Fifteen Week Semester Twelve Week Semester
1. Lab Safety and Exp. 1 1. Lab Safety and Exp. 1
2. Exp. 2 2. Exp. 3 and Exp. 2, Parts A&B (or Part C)
3. Exp. 3 3. Exp. 4 and Exp. 2, Part C (or Parts A&B)
4. Exp. 4 4. Exp. 5
5. Exp. 5 5. Exp. 6
6. Exp. 6 6. Exp. 7 and Exp. 8 part A
7. Exp. 7 7. Exam 1 and Exp. 8 part B
8. Exam 1 and Exp. 8 part A 8. Exp. 9 and Exp. 8 part C
9. Exp. 8 parts B and C 9. Exp. 10*
10. Exp. 9 10. Exp. 11
11. Exp. 10* 11. Exp. 12
12. Exp. 11 12. Exam 2 and Exp. 13
13. Exp. 12
14. Exp. 13 or Exam 2
15. Exam 2 or Exp. 13
*The NaOH solution standardized in Experiment 9 is used again in this experiment.
Bergen Community College 5 General Chemistry II Laboratory
BERGEN COMMUNITY COLLEGE
SAFETY REGULATIONS FOR THE CHEMISTRY LABORATORY
1. Read these safety regulations carefully and be sure you understand them. Before each
laboratory session, your instructor will discuss any safety hazards that may be associated
with that day’s experiment. Therefore, it is imperative that you come to lab on time.
2. Due to safety concerns students who arrive after the pre-lab presentation may not be allowed to perform that particular lab experiment.
3. It is strongly suggested that you obtain a hall locker from the Security Office. Only your lab manual, notebook, and calculator are allowed on the lab bench.
4. Report all accidents, no matter how minor, to your instructor at once. No one in the lab is permitted to give out bandages or medication. You must see the College Nurse.
5. Safety glasses or goggles are required and must be worn by everyone in the lab when experiments are being conducted. Contact lenses are not recommended in the chemistry
lab. Safety glasses are provided by the college, but students may purchase their own.
6. Do not perform any unauthorized experiment.
7. Do not taste anything in the laboratory. Never eat, drink or smoke in any of the labs.
8. You must tie back long hair. Do not wear open-toed shoes, shorts, fuzzy sweaters, loose sleeve shirt or any dangling jewelry. You must cover bare midriffs. You are advised to
wear a lab coat or old clothing to the lab.
9. Do not fill pipettes by mouth. Rubber bulbs or pipette pumps are provided. The instructor will demonstrate how these are to be used.
10. Exercise care when noting the odor of fumes. Use ‘wafting’ if you are directed to note an odor.
11. Do not force glass tubing or a thermometer into rubber stoppers. Lubricate with water and introduce it gradually and gently into the stopper, or insert through a cork borer.
Protect your hands with toweling when inserting without a cork borer.
12. Never point a test tube containing a reaction mixture (especially when heating) toward yourself or another person.
13. No ‘fooling around’ in the laboratory. A less than serious approach to lab work may result in an accident.
14. Before connecting or disconnecting electrical equipment, make sure that the switches are in the off position.
Safety Regulations
Bergen Community College 6 General Chemistry II Laboratory
15. Never work in the laboratory alone.
16. Make sure all apparatus is properly supported on the workbench.
17. Read the label on every bottle twice before using it in the laboratory. Many chemical names are very similar but are very different chemically.
18. Replace caps and stoppers on bottles immediately. Return spatulas to their correct place immediately after use. Do not mix them up.
19. Do not remove or relocate any chemical that has been placed in the hood. Sample it in the hood.
20. Never light a Bunsen burner with a cigarette lighter. Use the strikers that are provided.
21. Students are responsible for keeping their work area neat and orderly. All spills are to be cleaned up immediately using the spill kits located on the instructor’s desk. Solid
chemical waste should be and placed in the appropriately labeled container. Liquid
chemical waste should be poured into the appropriately labeled container. All waste
material should be left in the hood for subsequent disposal. If there is doubt about proper
disposal, ask the instructor.
22. Wash all glassware immediately after use. Place clean glassware on drying rack or in the designated bin on the counter.
23. Dispose of broken glassware in the labeled broken glassware boxes.
24. Wash your hands before leaving the laboratory.
25. You must notify your instructor of any chemical to which you are allergic.
26. If you are pregnant or planning to become pregnant this semester, you must notify your physician that you are enrolled in a chemistry lab course. You and your physician must
decide whether or not it is appropriate for you to remain in the course.
Note the location of the following safety equipment so that you can get to it quickly
in an emergency.
SAFETY EQUIPMENT LOCATION
FIRE EXTINGUISHER
SAFETY SHOWER
EYEWASH
EMERGENCY PHONE
FIRE ALARM
NEAREST EXIT
Safety Regulations
Bergen Community College 7 General Chemistry II Laboratory
SAFETY IN THE LABORATORY
True False
1. Safety glasses must be worn by everyone working in the lab. T F
2. Only major accidents in the lab need to be reported T F
3. Material Safety Data Sheets are provided in the lab T F
4. Eating and drinking are permitted in the lab T F
5. It is OK to taste a chemical as long as it smells good T F
6. Only authorized experiments are to be performed T F
7. You should wear shoes at all time in the lab T F
8. In order to save time, it is permissible to weigh hot objects T F
9. Broken glassware should be disposed of in the appropriate box T F
10. Working alone in the lab is an acceptable practice T F
A typical Chemistry Laboratory safety YouTube video link is given below: (hold Ctrl Key
and hover the mouse over the link) https://www.youtube.com/watch?v=UKovNdse5MU
Please complete sign this attached form. Remove it from the safety regulations and hand it to
your Laboratory Instructor.
I, the undersigned, have read the Divisional Safety Regulations for the Chemistry
Laboratories. I understand them and will abide by them.
Print your name: ________________________________________________________
Signature: ____________________________________________________________
Date: ____________________________________________________________
Course Name and Number:_______________________________________________
Bergen Community College 8 General Chemistry II Laboratory
INTEGRITY OF DATA GUIDELINES
One purpose of a laboratory course is to reinforce the concepts covered in the lecture
course. A second, equally important purpose, is to experience working in a chemistry lab, and
to learn about practices and procedures that are employed in such an environment. In addition
to specific laboratory procedures that will be covered in the array of experiments, there are
two universal practices in all laboratory settings- Laboratory Safety, which was discussed in
the previous pages, and Integrity of Data Guidelines.
These guidelines are used in all laboratory settings- from the traditional research
laboratory to hospitals and the physician’s office. The purpose of the guidelines is to ensure
that data is recorded in such a way that its veracity, or authenticity, cannot be questioned.
Taken as a whole, these practices protect the integrity of the data by preventing it from being
changed or recorded in error. Students are expected to follow these integrity of data guidelines
when collecting and recording data. The guidelines are as follows:
1. Data sheets must include the date and the student’s name.
2. Data is recorded in blue or black non-erasable ink; no white-out is permitted.
3. If a mistake is made while entering data, a single line is used to cross it out
and the correct entry is made nearby. (The original entry must be legible.)
4. No transcription is permitted. (Data is recorded directly into the data sheets.)
5. Data is recorded at the time it is observed.
In most laboratories today, notebooks are electronic rather than paper. Although this
renders a different set of guidelines, their purpose is the same- to ensure the authenticity of
data. Laboratory notebook software does not permit a change to be made once data has been
entered. In instances where a change is required, there is a record of the original entry. When
a measurement is recorded on a scrap of paper, that original data is scanned and becomes a
part of the notebook. These and other practices concerning electronic lab notebooks, along
with the guidelines described above regarding paper notebooks, work together to protect the
integrity of experimental data.
Bergen Community College 9 General Chemistry II Laboratory
Experiment 1
Heat of Fusion
OBJECTIVE: To determine the heat of fusion of water.
BACKGROUND:
When the solid phase of a molecular substance is converted to the liquid phase, energy,
in the form of heat, must be added in order to break the attractions between the molecules.
These intermolecular forces in a solid hold the molecules locked into position. Although the
molecules vibrate in place, they do not move relative to each other, i.e. they have no
translational movement. In contrast, the molecules in the liquid phase, although close to one
another, do have translational movement. They are constantly making and breaking
intermolecular attractions as they move about in random translational motion.
As heat is added to a molecular substance in the solid phase, the kinetic energy of the
molecules increases, resulting in greater vibrational motion, and evidenced by an increase in
temperature. This process continues until the melting point is reached, when molecules begin
to have sufficient energy to break the attractive forces holding them in position, and the
substance begins to melt. At this point, added energy results in breaking attractive forces rather
than in increased movement, and the temperature remains constant throughout the melting
process. When the entire sample has become a liquid, added heat increases the kinetic energy
and the temperature increases once again.
A similar transition occurs in converting a substance from the liquid to the gas phase.
As heat is added once the boiling point has been reached, this energy is used to break
intermolecular forces between molecules in the liquid phase. Again, during the process of
vaporization, the temperature remains constant.
These relationships can be summarized in a heating curve, as illustrated in the figure
on the following page.
The amount of heat required to convert a substance from the solid to the liquid phase
is quantified as the heat (or enthalpy) of fusion, ∆Hfus. It is a physical property, and can be
reported as heat per gram of substance or per mole of substance. The latter is often referred to
as the molar heat of fusion.
Experiment 1 Heat of Fusion
Bergen Community College 10 General Chemistry II Laboratory
In this experiment, the heat of fusion of water, in joules/gram, will be determined using
a coffee cup calorimeter where a sample of ice has melted in tap water. The amount of heat
given up by the tap water in the calorimeter as it cools (qwater) will be absorbed as heat by a
sample of ice as it melts (qfusion) and as this melted ice warms (qmelted ice). Assuming no loss of
heat to the surroundings, the sum of these must equal zero.
qwater + qfusion + qmelted ice = 0
Rearranging,
qfusion = – qwater – qmelted ice (eq. 1)
Values for both qwater and qmelted ice are obtained from the following equations, where m
represents mass, c represents the specific heat of water (4.18 J/g ºC) and ∆T represents the
change in temperature.
q = m c ∆T (eq. 2)
∆T = Tfinal – Tinitial (eq. 3)
Once the heat of fusion is determined, the experimental error can be found as follows.
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐸𝑟𝑟𝑜𝑟 = |𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒|
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 × 100
Three determinations will be made using different sample sizes. A graph of the heat,
in joules, absorbed in melting the ice (qfusion) as a function of the mass of melted ice, in grams,
will be constructed. The slope of the line represents the heat of fusion of water.
T em
p er
at u re
Heat added
melting
vaporization
Figure 1: Heating Curve
(eq. 4)
Experiment 1 Heat of Fusion
Bergen Community College 11 General Chemistry II Laboratory
REAGENTS: Ice EQUIPMENT: 150-mL beaker Tap water 100-mL graduated cylinder
coffee cup calorimeter 400-mL beaker
thermometer, or thermocouple
PROCEDURE:
1. Measure and record the mass of a 150-mL beaker. Set it aside ready to use in step 5.
2. Tare a coffee cup calorimeter. Add 100 mL tap water using a graduated cylinder.
Measure and record the mass. Place the calorimeter in a 400-mL beaker for stability.
3. Measure and record the initial temperature of the tap water in the calorimeter.
4. Add sufficient ice to fill the volume of water, and gently stir with the thermometer.
5. When the temperature reaches between 0ºC and 5ºC, record the final temperature.
Immediately pour the water into the 150-mL beaker, leaving the unmelted ice behind.
6. Measure and record the mass of the beaker and contents.
7. Repeat steps 1 – 6 using 70 mL tap water, and again using 40 mL.
Disposal: Water may be disposed of down the drain.
CALCULATIONS:
A. Perform the following calculations for each of the three determinations.
1. Determine the mass of the contents of the beaker. This is the mass of the original tap water
in the calorimeter, plus that of the melted ice.
2. Determine the mass of the melted ice by subtracting the mass of the tap water from the
mass of the beaker contents.
3. Determine the temperature change for the tap water, ∆Twater, using eq. 3.
4. Determine the temperature change for the melted ice, ∆Tmelted ice, using eq. 3. The initial
temperature for the ice is assumed to be 0 ºC.
5. Determine qwater using the mass and temperature change of the tap water and eq. 2.
6. Determine qmelted ice using the mass and temperature change of the melted ice and eq. 2.
7. Determine qfusion using eq. 1.
B. Prepare a graph in Excel* of qfusion, in joules, as a function of the mass of melted ice, in
grams. Determine the heat of fusion for water from the graph.
*See Appendix C for directions on graphing.
Bergen Community College 12 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 1: Heat of Fusion
Data: Determination 1 Determination 2 Determination 3
Initial mass of beaker ____________ ____________ ____________
Mass of tap water ____________ ____________ ____________
Initial temperature ____________ ____________ ____________
Final temperature ____________ ____________ ____________
Final mass of beaker ____________ ____________ ____________
Results:
Mass of beaker contents ____________ ____________ ____________
Mass of melted ice ____________ ____________ ____________
ΔT of tap water ____________ ____________ ____________
ΔT of melted ice ____________ ____________ ____________
Heat for water, qwater ____________ ____________ ____________
Heat for melted ice, qmelted ice ____________ ____________ ____________
Heat for fusion of ice, qfusion ____________ ____________ ____________
Heat of fusion of ice _______________________
Bergen Community College 13 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 1: Heat of Fusion
POST-LAB QUESTIONS:
1. The heat of fusion of water is 333 J/g. Determine the percent error using equation 4.
2. Determine the amount of heat required to raise the temperature of a 22.5-gram
sample of copper from 125 ºC to its melting point of 1084 ºC, and then melt the
copper. (The specific heat of copper is 24.4 J/mol ºC and its heat of fusion is
13.26 kJ/mol.)
3. If the ice had begun at a temperature lower than 0 ºC, would the calculated value of
the heat of fusion have been higher, lower, or unchanged? Briefly explain.
Bergen Community College 14 General Chemistry II Laboratory
Experiment 2
Intermolecular Forces
OBJECTIVE: To relate intermolecular forces of molecules to physical properties.
BACKGROUND:
The attractive forces between molecules and their neighbors are called intermolecular
forces. These forces are much weaker than the intramolecular forces within a substance- the
covalent (or ionic) bonds. Intermolecular forces are the attractions that need to be overcome
for a molecular solid to melt, and for a liquid to vaporize. Therefore, the strength of these
attractive forces influence a substance’s physical properties. The stronger the intermolecular
forces, the higher melting point, boiling point, heat of vaporization, and other properties. The
following table summarizes the boiling points of some molecular compounds.
Compound Formula Polarity Molecular Structure Boiling Point
Methane CH4 Nonpolar
- 161oC
Propane C3H8 Nonpolar
- 42oC
Butane C4H10 Nonpolar
10oC
Hexane C6H14 Nonpolar
70oC
Acetone C3H6O Polar
56oC
Ethanol C2H6O Polar
77oC
Water H2O Polar
100oC
Experiment 2 Intermolecular Forces
Bergen Community College 15 General Chemistry II Laboratory
There are three general types of intermolecular forces. All substances exhibit London
Dispersion Forces (LDF), and they are generally the weakest of the three types. These London
forces are due to the attractions between small, temporary dipoles that arise from the constant,
random movement of the electrons in a substance. As molar mass increases, the size of the
electron cloud increases as well. It becomes more easily distorted, and produces temporary
dipoles of greater magnitude. This causes the attractions to be stronger, requiring more energy
for both fusion and vaporization. For halogens, this results in increasing melting and boiling
points, shown by the fact that at room temperature F2 and Cl2 are gaseous, Br2 is liquid and I2
is solid. The extent to which the electron cloud can be distorted is called polarizability.
Dipole-dipole forces exist between molecules that are polar. Since the dipoles are
permanent, these attractions are generally stronger than London Dispersion Forces. This
means that a polar molecule with similar molar mass as a nonpolar molecule will have higher
melting points and boiling points. Not all molecules containing polar bonds are polar. The
polar bonds must be unevenly dispersed in the molecule in order to produce a polar molecule.
CO2 and CBr4, for example, have polar bonds but are not polar molecules.
The third type of intermolecular force is hydrogen bonding, a specific type of dipole-
dipole attraction that is stronger than other dipole-dipole attractions. Hydrogen bonds form
when a hydrogen atom is covalently bonded to a very electronegative atom. This causes its
electron to be drawn away from its nucleus. The positive hydrogen is then attracted to the very
electronegative atom in a neighboring molecule. In order to observe hydrogen bonding, the
hydrogen atom must be covalently bonded to fluorine, oxygen or nitrogen. A hydrogen atom
bonded to a carbon atom cannot create a hydrogen bond. It’s
important to note that, despite its name, a hydrogen bond is
an intermolecular force, not a bond. The figure to the right
illustrates H-bonding between water molecules. H-bonding
is important in biochemistry; the structure of a biopolymer is
largely determined by the formation of hydrogen bonds.
The relative strengths of the three types of intermolecular forces, and thus boiling
points, are generally as follows:
London Dispersion Forces < Dipole-Dipole Forces < H-Bonding
However, this is not always true. Since molar mass is also a factor, a large non-polar molecule
can have a higher boiling point than a compound that interacts with dipole-diploe forces, or
even a substance with H-bonding. For example, octane, a component of gasoline, has a boiling
point of 125oC- much higher than acetone (dipole-dipole) and H2O (H-bonding). This is due
to the polarizability of the large electron cloud.
To make comparisons of the intermolecular forces of a substance, evaporation rate can
be used instead of boiling point. Evaporation rate is the ratio of the change in temperature to
the change in time as a substance evaporates. A faster rate of evaporation translates to a lower
boiling point and, in turn, weaker intermolecular forces.
Experiment 2 Intermolecular Forces
Bergen Community College 16 General Chemistry II Laboratory
Boiling point is not the only physical property affected by the type of intermolecular
forces a substance has. Solubility is also dependent upon the polarity of a molecule. The term
“like dissolves like” suggests that polar solutes dissolve in polar solvents and nonpolar solutes
dissolve in nonpolar solvents. Therefore, polarity, and the associated intermolecular forces,
determine a substance’s solubility in water and in other solvents.
The solubility of a solid in a liquid is readily observed. When liquids mix forming a
homogeneous solution, they are said to be miscible; if they do not mix, they are immiscible.
If two liquids are miscible, there is no observable interface between the two. If the two liquids
are immiscible, two distinct layers are seen.
In this experiment, both miscibility and evaporation rates of acetone, ethanol, hexane,
and water will be determined. Salt solubility in an ethanol-water mixture will also be observed.
REAGENTS: acetone EQUIPMENT: Thermometers ethanol filter papers
hexane rubber band
distilled water tape
sodium chloride stop watch
5 test tubes containing a 50% by volume: small test tubes
water and acetone wood block
water and hexane
hexane and acetone
hexane and ethanol
ethanol and acetone
SAFETY ALERT:
- Do not pour any materials into the sink!
- Wash hands and laboratory bench after the experiment.
- Acetone: Extremely flammable liquid and vapor. Vapor may cause flash fire. Causes eye
irritation. Breathing vapors may cause drowsiness and dizziness. Causes respiratory tract
irritation. Aspiration hazard if swallowed. Can enter lungs and cause damage. Prolonged or
repeated contact may dry the skin and cause irritation.
Hexane: Extremely flammable liquid and vapor. Vapor may cause
flash fire. Breathing vapors may cause drowsiness and dizziness.
Causes eye, skin, and respiratory tract irritation. May be harmful if
absorbed through the skin. Aspiration hazard if swallowed and enters
lungs causing damage. Possible risk of impaired fertility. Long-term
exposure may cause damage to the nervous system of the extremities.
Bergen Community College 17 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 2: Intermolecular Forces
PRE-LABORATORY QUESTIONS
1. Which of the substances used in this experiment must be handled in the fume hood?
2. Identify the strongest type of intermolecular forces in acetone, ethanol, water and hexane.
(Structures listed on page 15.)
3. Predict the relative strength of the intermolecular forces in the four liquids above.
______________ < _______________ < _______________ < _______________
4. Water (MW = 18 g/mol) has higher boiling point than ethanol (MW = 46 g/ mol) and acetone
(MW = 58 g/mol). Why does water, such a small molecule, have such a high boiling point?
5. What is the meaning of the term "like dissolves like"?
Experiment 2 Intermolecular Forces
Bergen Community College 18 General Chemistry II Laboratory
PROCEDURE:
Part A: Evaporation Rate of Acetone and Hexane
MUST BE PERFORMED IN A FUME HOOD!
1. Obtain a thermometer that reads to the tenth of a Celsius degree.
2. Add a few mL acetone into a small test tube.
3. Wrap a piece of filter paper around the end of a thermometer; secure it with a rubber band.
4. Place the thermometer into the test tube. Wait approximately 10-20 seconds, until the
temperature is stabilized, and record this as the temperature at time = 0 seconds.
5. Remove the thermometer from the test tube, touching the tip of the filter paper to the tube
to leave behind any excess liquid. Tape it to a woodblock in the fume hood, so that the
thermometer lies horizontally, and the filter paper does not touch the workbench.
6. Record the temperature every 30 seconds for 5 minutes.
7. Repeat steps 1 – 6 using hexane.
Disposal: Return any remaining acetone and/or hexane to the appropriate collection
container. Leave the filter paper in the hood to dry, then dispose of in trash bin.
Figure 1: Thermometer, filter paper and rubber band in test tube with
liquid (left), and taped to wood block (right).
Experiment 2 Intermolecular Forces
Bergen Community College 19 General Chemistry II Laboratory
Part B: Evaporation Rate of Water and Ethanol
Working at the lab bench, not in the fume hood, follow steps 1 – 6 as described in part A
using ethanol, and again using water.
Disposal: Return any remaining ethanol to the appropriate collection container. Leave the
filter papers on the lab bench to dry, then dispose of in trash bin.
Part C: Solubility
1. Obtain a set of five prepared test tubes containing mixtures of the following:
water and acetone
water and hexane
hexane and acetone
hexane and ethanol
ethanol and acetone
2. Prepare a sixth test tube by adding equal amounts (about 1 mL) of ethanol and water.
3. Inspect the test tubes and record the solubility of each mixture: S = soluble, I = insoluble.
4. Add a pea-sized quantity of sodium chloride to the ethanol/water mixture. Gently shake
the tube. Make and record observations.
Disposal: Place the ethanol-water-salt mixture in the designated disposal container.
CALCULATIONS:
Parts A and B:
1. Subtract each temperature from the initial temperature in order to find ∆T.
2. Prepare (by hand) a graph* of ∆T as a function of time for each of the liquids on one piece
of graph paper.
3. Determine the relative evaporation rates for the four liquids by comparing the curves. A
faster evaporation rate generates a steeper curve.
*See Appendix C for directions on graphing.
Bergen Community College 20 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 2: Intermolecular Forces
Data: Parts A and B: Evaporation Rates
Results: Parts A and B: Evaporation Rates
Experimentally determined relative strengths of intermolecular forces:
_______________ < ________________ < ________________ < ________________
time,
seconds
Acetone Hexane Ethanol Water
T, °C ∆T, °C T, °C ∆T, °C T, °C ∆T, °C T, °C ∆T, °C
0.
30.
60.
90.
120.
150.
180.
210.
240.
270.
300.
Bergen Community College 21 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 2: Intermolecular Forces
Data: Part C: Solubility
Acetone Ethanol Hexane
Water
Hexane N/A
Ethanol N/A N/A
Observation: Solubility of NaCl in Ethanol/Water Mixture
Results: Part C: Solubility
Experimentally determined polarity of liquids: Given that water is polar, list the other liquids
as polar or nonpolar.
Polar: _____________________________________________________
Nonpolar: __________________________________________________
Bergen Community College 22 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 2: Intermolecular Forces
POST-LAB QUESTIONS:
1. Explain the experimentally determined relative strengths of intermolecular forces. Did the
results match what was predicted in the Pre-Lab questions? Discuss any discrepancies.
2. Explain the polarity of the four liquids studied. Did the experimental miscibility results
match the polarities listed in the table on page 15? Discuss any discrepancies.
3. Explain why hexane and water do not mix.
4. Two layers are observed when sodium chloride is added to the ethanol-water mixture, a
phenomenon known as "salting out". Explain why the water and ethanol separate.
Bergen Community College 23 General Chemistry II Laboratory
Experiment 3
Spectroscopy
OBJECTIVE: To determine the molar concentration of Co+2 using spectroscopy.
BACKGROUND:
The field of Spectroscopy deals with the interaction between electromagnetic radiation
and matter. In the ultraviolet and visible portions of the spectrum, light is absorbed by a
substance when electrons move from an orbital of lower energy to one of higher energy. The
wavelength of the absorbed light corresponds to the difference in energy of the two orbitals.
When transition metal ions are dissolved in water, the energy differences of the d orbitals
correlate to wavelengths in the visible portion of the electromagnetic spectrum. Since only
certain visible wavelengths are absorbed, these solutions are colored. The observed color
corresponds to the wavelengths that are transmitted, passing through the sample without being
absorbed. Each aqueous transition metal ion has its own characteristic color. An Absorption
Spectrum shows the amount of absorbed light across a wide range of wavelengths. Because
the amount of light absorbed depends on concentration, visible spectroscopy is often used to
determine the concentration of aqueous transition metal ions.
A spectrophotometer is used to measure the amount of light that passes through a
sample. It consists of a light source, a diffraction grating to separate the light into the individual
wavelengths, a lens to focus the light, and a detector that measures the amount of light that has
passed through the sample. The light passing through the sample is quantified as the
transmittance, T, and is the ratio of the intensity of the light emerging from the sample, I, to
the intensity of the light entering the sample, I0. Oftentimes, the percent transmittance, %T, is
used. These relationships are shown below.
I
T = %T = T x 100
I0 I I0
The amount of light absorbed (removed) by a sample is quantified as Absorbance, A, and is
related to the Transmittance by the following relationship.
A = - log (%T/100) As absorbance is a logarithmic relationship, it has no unit.
When making absorbance measurements, the spectrophotometer must first be calibrated using
a solution that contains all of the components of the solution to be analyzed, other than the
absorbing species. Such a solution is referred to as a blank.
Experiment 3 Spectroscopy
Bergen Community College 24 General Chemistry II Laboratory
An absorbance spectrum is a
graph of absorbance as a function of
wavelength. The wavelength having the
maximum absorbance, λmax, is called the
optimal wavelength, or the analytical
wavelength. When making absorbance
measurements for the determination of
concentration, this wavelength is always
used. Because the absorbance value is
large, this increases the precision of the
measurements. Because this represents
the top of the curve rather than the
shoulder, it increases the accuracy.
The absorbance of a solution depends on the distance the light travels in the sample
and the molar concentration of the absorbing species, as well its identity. This relationship,
known as the Beer-Lambert Law, or more commonly, Beer’s Law, is A = × l × c, where is the molar absorptivity, a characteristic specific to the absorbing species, l is path length, and
c is the molar concentration. Since throughout the course of an experiment, both the path
length and the molar absorptivity remain constant, the absorbance depends solely on the
concentration of the absorbing species. Although possible to determine concentration by
comparison of an unknown’s absorbance to that of just one solution, it is much more common
to make several absorbance measurements in order to increase accuracy.
A Beer’s Law plot is a graph of absorbance
as a function of concentration. Absorbance values
for several standard solutions, those with known
concentration, are obtained at λmax and graphed.
The concentration of an unknown solution is found
using its absorbance and the equation of the line. If
the graph has been prepared by hand, the
absorbance of the unknown is located on the y-axis,
and the corresponding value on the x-axis is found.
It’s important to note that the Beer-Lambert relationship holds only for solutions that
are sufficiently dilute. If the unknown solution’s absorbance is greater than those used in the
plot, a dilution is required in order to determine concentration. Extrapolations are not possible.
In this experiment, the concentration of a cobalt(II) nitrate solution will be determined.
First, a solution will be prepared from solid cobalt(II) nitrate. Next, absorbance values of this
solution at various wavelengths will be measured, from which the analytical wavelength will
be determined. A series of standard solutions will then be prepared by dilution. Absorbance
values for these solutions will be measured at the analytical wavelength, from which a Beer’s
Law Plot can be constructed. Lastly, the absorbance of an unknown solution will be measured,
and its concentration determined from the Beer’s Law Plot.
A b
so rb
an ce
Wavelength
Absorbance Spectrum λ max
y = 6.4848x R² = 0.9997A
b so
rb an
ce
Concentration
Beer's Law Plot
Bergen Community College 25 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 3: Spectroscopy
PRE-LABORATORY QUESTIONS
1. Why are some solutions colored, and others are colorless?
2. How does a solution interact with light?
3. What is meant by the term analytical, or optimal, wavelength?
4. What is plotted on the x and y axes of an absorbance spectrum?
What is plotted on the x and y axes of a Beer’s Law plot?
5. What is the absorbance (A) of a solution with percent transmittance (%T) of 25.1?
Experiment 3 Spectroscopy
Bergen Community College 26 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Distilled water Beaker, 50-mL
0.100 M Co(NO3)2 Volumetric flask, 50-mL
(made by dissolving 29.10 g Beral pipettes
Co(NO3)2∙6H2O in water to Volumetric pipette, 25-mL yield 1.00 L solution.) Spectrophotometer
Cuvettes
Lint-free Wipes
SAFETY ALERT:
Cobalt(II) nitrate: May cause an allergic skin reaction, allergy, asthma symptoms or breathing
difficulties if inhaled. Suspected of causing genetic defects. May cause cancer by inhalation.
PROCEDURE:
Part A: Determination of the Optimal Wavelength
1. Turn the spectrophotometer on. Allow it to warm up for 10-15 minutes before use.
2. Set the spectrophotometer wavelength to 430 nm.
3. Obtain 40-50 mL of the 0.100 M Co+2 stock solution.
4. Prepare a cuvette by adding the Co+2 solution until ¾ full. Clean with a lint-free wipe.
5. Prepare a blank by adding d-H2O to a cuvette until ¾ full. Clean with a lint-free wipe.
6. Place the blank in the spectrophotometer and zero the instrument.
7. Replace the blank with the Co+2 cuvette. Read and record the absorbance.
8. Set the wavelength to 450 nm. Zero the instrument with the blank, replace with the Co+2,
read and record the absorbance.
9. Continue increasing the wavelength by 20 nm, zeroing with the blank, reading and
recording the Co+2, ending with a wavelength of 590 nm.
10. Determine the optimal wavelength, that corresponding to the greatest absorbance value.
Experiment 3 Spectroscopy
Bergen Community College 27 General Chemistry II Laboratory
Part B: Preparation of Beer’s Law Plot for Co+2
1. Set the spectrophotometer to the optimal wavelength obtained in Part A and zero the
instrument with the blank.
2. Replace the blank with the Co+2 cuvette used in Part A. This is standard solution 1. Read
and record the absorbance. (Do not re-zero the spectrophotometer.)
3. Dilute 25 mL of standard solution 1 to 50 mL using a volumetric pipet and volumetric
flask.* This is standard solution 2.
4. Condition the Co+2 cuvette with standard solution 2 and add until ¾ full. Wipe clean.
5. Place the standard solution 2 cuvette in the instrument. Read and record the absorbance.
6. Repeat steps 3 – 5, diluting solution 2 to prepare standard solution 3.
7. Repeat steps 3 – 5 again, diluting solution 3 to prepare standard solution 4.
8. Once again, repeat steps 3 – 5 diluting standard solution 4 to prepare standard solution 5.
(Do not re-zero or turn off the spectrophotometer until Part C has been completed.)
* See Appendix B for directions on using volumetric glassware.
Disposal: Pour all Co+2 solutions into the appropriate waste collection container.
Part C: Determination of Concentration of Unknown Co+2 Solution
1. Obtain an unknown Co+2 solution from the instructor.
2. Condition a cuvette; add the unknown solution until ¾ full. Wipe clean.
3. Measure and record the absorbance.
CALCULATIONS:
Part A: Prepare an absorbance spectrum in Excel**, marking the analytical wavelength.
Part B: 1. Determine the concentrations of solutions 2 - 5 using M1V1 = M2V2.
2. Prepare a Beer’s Law plot for Co+2 in Excel.**
Part C: Determine the concentration of the unknown using its absorbance value and the
equation of the line from the Beer’s Law Plot.
**See Appendix C for directions on graphing.
Bergen Community College 28 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 3: Spectroscopy DATA and Results:
PART A: Determination of Optimal Wavelength
Wavelength,
nm 430 450 470 490 510 530 550 570 590
Absorbance
(A)
Analytical wavelength: ____________ nm
PART B: Data for Beer’s Law Plot using the optimal wavelength (max).
Solution # 1 2 3 4 5
Concentration, M
Absorbance (A)
PART C: Determination of Concentration of Unknown Co+2 Solution.
Unknown # Absorbance (A) Concentration (M)
Show calculations for Parts B and C on a separate sheet.
Bergen Community College 29 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 3: Spectroscopy
POST-LABORATORY QUESTIONS
1. A solution of Co2+ ion contains some small impurities of Cu2+ and Ni2+ ions. The absorption
spectrum of each ion’s 0.10 M solution is given in the figure below. What wavelength should
be used in order to determine the concentration of the Co2+ ions in the unknown solution?
2. a. If the cuvettes used for the standard solutions are rinsed with d-H2O, but not
conditioned or dried, will the calculated concentration of the unknown be too
high, too low, or unaffected? Briefly explain.
b. If the flasks used to prepare the standard solutions are rinsed with d-H2O, but not
conditioned or dried, will the calculated concentration of the unknown be too high,
too low, or unaffected? Briefly explain.
Bergen Community College 30 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 3: Spectroscopy
POST-LABORATORY QUESTIONS (continued)
3. What are the common calibration errors in the Beer's Law plot?
4. Curcumin (MW = 369 grams/mol) is one the active compounds of the turmeric and has anti-
inflammatory and antibacterial properties. It is a yellow substance, insoluble in water but
soluble in alcohol. A chemist extracts curcumin by dissolving 0.500 grams of turmeric in 500.0
mL of alcohol and determines the absorbance of the solution to be equal to 1.400 absorbance
unit. The path length of the cuvette is 1.0 cm and the molar absorptivity of curcumin is 55,
000 L × mol-1 × cm-1. Determine the mass percent of curcumin in turmeric.
Bergen Community College 31 General Chemistry II Laboratory
Experiment 4
Percent Copper in Brass
OBJECTIVE: To determine the percentage by mass of copper in a brass sample.
BACKGROUND:
Brass is an alloy of copper and zinc. The mass percentage of copper in brass can be
determined by a two-step methodology. In the first step, the brass is reacted with concentrated
nitric acid. This produces aqueous copper(II) nitrate and aqueous zinc nitrate. Copper reacts
readily with oxidizing agents. In this reaction, concentrated nitric acid oxidizes copper metal
to the copper(II) ion. As products of this reaction, copper(II) nitrate is a water-soluble salt that
produces a blue solution, and nitrogen monoxide, NO, oxidizes in presence of the air to a
dense, toxic, red-brown gas, NO2. The zinc from the brass forms aqueous zinc nitrate, which
is colorless. The balanced chemical equations are:
3 Cu(s) + 8 HNO3(aq) → 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(l)
4 Zn(s) + 10 HNO3(aq) → 4 Zn(NO3)2(aq) + NH4NO3(aq) + 3 H2O(l)
Aqueous copper(II) ion absorbs light in the visible region, while aqueous zinc ion does
not. In step two, a visible spectrophotometer, or colorimeter, is used to analyze the color
intensity of the copper(II) nitrate solution that forms.
A solution of copper(II) nitrate appears blue because when white light passes through
the solution, the dissolved copper(II) ions absorb the red components of the light, while
transmitting the blue portions of the light. Most transition metal salts are hydrated and colored.
This is due to unfilled d-orbitals. Copper ions are colored; however, zinc ions usually produce
colorless solutions, due to filled 3d-orbitals. Most hydrated metal salts differ in color. The
following table illustrates the number of moles of crystal water and the color of the salts.
SALT Moles
Water COLOR FORMULA
Molar
Mass
g/mol
Analytical
Wavelength
nm
Zinc nitrate 6 Colorless Zn(NO3)2∙6H2O 297.49 Does not absorb
visible light
Copper(II)
nitrate 3 Blue Cu(NO3)2∙3H2O 241.54 635
Experiment 4 Percent Copper in Brass
Bergen Community College 32 General Chemistry II Laboratory
The absorbance of a colored solution is measured as A = ∙b ∙c, where is the molar absorptivity (M/cm), b is path length (cuvette diameter, 1 cm) and c is concentration (mol/L).
The absorbance of incident light by a colored solution is depicted below.
The molar absorptivity constant for an absorbing species can be determined using
Beer’s Law* by measuring the absorbance of different solutions with known concentrations.
A plot of absorbance as a function of concentration gives a curve with the slope representing
the molar absorptivity constant. The blank used to zero the spectrophotometer contains dilute
nitric acid solution. The working (analytical) wavelength has been determined using a 0.10 M
Cu(NO3)2 solution, and finding the highest absorbance reading at wavelengths between 400
nm and 700 nm. As a copper solution absorbs red light (635 nm), the analytical wavelength,
λmax, is 635 nm.
After constructing the calibration curve (Beer’s Law Plot*), the absorbance of the
unknown solution is measured. Its concentration is obtained by reading the corresponding
value from the x-axis. If the concentration of the unknown exceeds that of the standard
solutions used in the calibration, the unknown is further diluted in order to be able to measure
its concentration. Calculations are then adjusted accordingly.
Once the concentration of the solution prepared from the oxidation of the copper in the
brass sample is obtained, the number of moles of copper and, thus, the mass, can be obtained.
The percentage copper in the brass sample is then determined as follows:
% 𝐶𝑢 = 𝑚𝑎𝑠𝑠 𝐶𝑢
𝑚𝑎𝑠𝑠 𝑏𝑟𝑎𝑠𝑠 𝑠𝑎𝑚𝑝𝑙𝑒 × 100 (equation 1)
In this experiment, the percent composition of copper in a brass (copper-zinc) sample
will be determined. First, a 0.30 M copper(II) nitrate stock solution, as well as a series of
standard solutions, will be prepared. Next, their absorbances will be measured, and a Beer’s
Law plot* will be constructed. The brass will then be reacted with concentrated nitric acid,
producing aqueous copper(II) and zinc ions. Finally, the absorbance of the unknown solution
produced in the reaction of the brass will be measured, and its concentration will be found
from the Beer’s Law plot.* Thereafter, the percent composition of the brass sample will be
determined.
*See Experiment 3 for information on Spectroscopy and Beer’s Law.
Bergen Community College 33 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 4: Percent Copper in Brass
PRE-LABORATORY QUESTIONS
1. How many grams of zinc nitrate zinc hexahydrate are needed to prepare 250.0 mL of
0.0500M Zn(NO3)2∙6H2O? (Zn = 65.39 g/mol; N = 14.01 g/mol; O = 16.00 g/mol; H= 1.008 g/mol)
2. How you would prepare 50.00 mL of 0.0400M Zn(NO3)2 from the solution in Question 1?
3. The following data were obtained for CuSO4∙5H2O solutions at a wavelength of 635 nm.
Concentration (M) Absorbance
0.0000 0.000
0.0100 0.120
0.0200 0.250
0.0400 0.460
0.0800 0.906
a. Using Excel,* plot a graph of absorbance as a function of concentration using the
above data. Include the graph with the pre-laboratory questions.
b. From the graph, determine the molarity of a copper sulfate pentahydrate solution
whose absorbance is 0.150 at 635 nm.
*See Appendix C for directions on graphing.
Experiment 4 Percent Copper in Brass
Bergen Community College 34 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Solid Cu(NO3)2 ∙ 3H2O Beakers, 100-mL and 150-mL Distilled water Volumetric flask, 100-mL
Brass sample Beral pipettes
6.0 M HNO3 Spectrophotometer
dilute nitric acid Cuvettes
Lint-free Wipes
Hot plate
Graduated cylinders, 10-mL and 25-mL
SAFETY ALERT
Safety Precautions
6 M nitric acid is very corrosive to all body tissue and toxic by ingestion and inhalation.
Nitric acid is also a strong oxidizing agent. Copper(II) nitrate is an oxidizing agent, and
moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves, and
a chemical-resistant apron. Please consult Material Safety Data Sheets for further safety,
handling, and disposal procedures.
Disposal
Consult the Catalog/Reference Manual for general guidelines and specific procedures
governing the disposal of laboratory wastes. The excess nitric acid solution may be
disposed of according to Flinn Suggested Disposal Methods #24b. The copper nitrate
solution may be flushed down the drain with excess water according to Flinn Suggested
Disposal #26b. www.flinnsci.com
PROCEDURE
Preparation of the Stock Solution
1. Turn the spectrophotometer on. Allow it to warm up for 10-15 minutes before use.
2. Tare a 100-mL beaker. Add 7.245 g Cu(NO3)2 ∙ 5H2O.
3. Add 50 mL distilled water and stir to dissolve. Transfer to a 100-mL volumetric flask.
4. Add 10 mL d-H2O to the beaker, swirl to rinse the beaker, and add to the flask.
5. Fill the flask to the mark with d-H2O. Cap the flask, and invert several times to mix. This
Makes 100 mL 0.300 M Cu+2 and is standard solution 1.
Experiment 4 Percent Copper in Brass
Bergen Community College 35 General Chemistry II Laboratory
Preparation of the Calibration Curve
1. Prepare a Cu+2 cuvette from solution 1 and a dilute nitric acid cuvette (which serves as
the blank) by filling ¾ full. Clean with a lint-free wipe.
2. Set the spectrophotometer to 635 nm, and zero the instrument with the blank.
3. Replace the blank with the Cu+2 cuvette. Read and record the absorbance. 4. Using a graduated cylinder to measure, add 40.0 mL of the stock solution to a 50-mL
volumetric flask. Add d-H2O to the mark and invert to mix. This is standard solution 2.
5. Condition and fill a cuvette ¾ full with standard solution 2. Wipe clean. Place in the
spectrophotometer. Read and record the absorbance.
6. Pour the standard 2 solution in the appropriate waste collection container, and rinse the
flask with d-H2O. Repeat steps 4 and 5 to prepare five additional standard solutions
according to the amounts given in the data table. Measure and record their absorbances.
Disposal: Pour all Cu+2 solutions into the appropriate waste collection container.
Preparation of the Sample
PERFORM THIS PART OF THE EXPERIMENT IN THE FUME HOOD!
1. Tare a 150-mL beaker and add 1 gram of the brass sample. Record the mass.
2. Set a hot plate to 60-70°C. While heating, add 25 mL of 6.0 M HNO3 to the beaker to
dissolve the brass.
3. When the brass is completely dissolved, stop heating and allow the sample to cool.
4. Slowly poor this solution to a 100-mL volumetric flask and fill to the mark with d-H2O.
5. Take a sample form this solution and measure and record the absorbance.
Disposal: Pour all Cu+2 solutions into the appropriate waste collection container.
CALCULATIONS
1. Determine the concentrations of solutions 2 - 7 using M1V1 = M2V2.
2. Prepare a Beer’s Law plot for Cu+2 in Excel. See Appendix C for directions on graphing.
3. Determine the concentration of the unknown using its absorbance value and the graph.
4. From the solution’s volume and concentration, determine the number of moles of Cu+2.
5. From the molar mass of copper, 63.54 g/mol, determine the number of grams of copper.
6. Determine the percentage copper in the sample using equation 1.
Bergen Community College 36 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 4: Percent Copper in Brass DATA:
Preparation of the Calibration Curve
Standard
Solution
Volume
0.3 M Cu(NO3)2
solution, mL
Cu2+ concentration,
mol/L
Absorbance
at 635 nm
1 50.0
2 40.0
3 25.0
4 10.00
5 5.00
6 2.00
7 1.00
Preparation of the Sample
Mass of brass sample ____________
Absorbance of Unknown Solution ____________
Results:
Cu+2 concentration ____________
Number of moles of copper ____________
Mass of copper ____________
Percent copper in brass sample ____________
Show calculations on a separate sheet.
Bergen Community College 37 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 4: Percent Copper in Brass
POST-LABORATORY QUESTIONS
1. A student weighs 1.050 grams of brass sample and dissolves it in 6 M nitric acid solution. After following the same experimental procedures, the absorbance of the
sample is recorded at 0.150. Using the graph obtained in the experiment. determine the
mass percent of the copper in the sample.
2. Explain why the Zn2+ ions do not absorb in the visible region of the electromagnetic spectrum. (Hint: Look at the electronic structure of the Zn+2.)
Bergen Community College 38 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 4: Percent Copper in Brass
POST-LABORATORY QUESTIONS (continued)
3. Duralumin is an alloy of aluminum (95%) and copper (5%). Could the technique used in this laboratory experiment to determine the percent copper also be used to
determine the duralumin in a sample? Briefly explain. (Hint: Check the color of
Aluminum nitrate)
4. UV-VIS spectroscopy was used to measure the spectra of two solutions, A and B. Solution A was a pink color, while solution B was a green color. The student recorded
the absorbance of each solution in the visible region of the electromagnetic spectrum
and the result is shown below.
If equal volumes of Solutions A and B are mixed, what wavelength should be used to
measure the concentration of Solution B in the mixture? Briefly explain.
Bergen Community College 39 General Chemistry II Laboratory
Experiment 5
Freezing Point Depression
OBJECTIVE: To determine the molar mass of an unknown organic compound through freezing point depression.
BACKGROUND:
Colligative properties are of solution properties that depend on the concentration of
solute particles, but not their identity. A pure substance has a set boiling point, freezing point
and vapor pressure under specified physical conditions. However, when solute is added to the
liquid, its presence alters the vapor pressure, in turn changing the boiling and freezing points
as well. Compared to the pure solvent, the solution has a lower vapor pressure, a lower freezing
point and a higher boiling point. The following graph illustrates these relationships.
Figure 1: Vapor Pressure as a function of Temperature
The change in the freezing point, that is the difference between the freezing point of
the solution and that of the pure solvent, ∆Tf, is quantified by the following:
∆Tf = Kf × i × m Equation 1
where Kf is the molal freezing point depression constant, i is the van’t Hoff factor, and m is
the molality. The freezing point depression constant, Kf, is unique to every solvent. Its value
for water is -1.86 oC/m. The van't Hoff factor is the number of dissociated particles per formula
of dissolved solute. For molecular solutes, such as C6H12O6, i = 1. For dilute solutions of ionic
solutes, i is the number of ions in the formula, e.g. 2 for NaCl and 3 for Na2SO4. Due to ion-
ion interactions, however, this value lessens as concentration increases. Concentration is
measured as molality, moles solute per kilogram solvent. Unlike molarity, molality is not
vulnerable to changes in temperature.
solution
1 atm
pure
solvent
Vapor
Pressure
Temperature solution boiling point
pure solvent boiling point
solution freezing point
pure solvent freezing point
Experiment 5 Freezing Point Depression
Bergen Community College 40 General Chemistry II Laboratory
One important application of colligative properties is in the determination of molar
mass of an unknown substance. A solution is prepared with the unknown dissolved in a solvent
whose freezing point and freezing point constant are known. Measuring the change in freezing
point, the molality of the solution can be found. With the masses of solute and solvent used in
preparing the solution, molar mass is then obtained.
In order to determine its freezing point, the temperature of a liquid is observed during
the cooling and freezing process, and a cooling curve, a graph of temperature as a function of
time, is then constructed.
Below is a cooling curve for a pure liquid. Note that during the freezing process, the
temperature remains constant. Once frozen, the temperature begins to drop as the liquid cools.
The freezing point is determined by extrapolation of the plateau to time = 0, and reading the
temperature from the graph. In rare instances, supercooling may occur, in which the solution
reaches a very low temperature and then rises back up to its freezing point, as depicted in the
graph. This is due to rearrangement of particles, which is an exothermic process. Hereafter,
the temperature rises and reaches its plateau.
Figure 2: Cooling Curve of a Pure Substance
The cooling curve of a solution is shown on the following page. Because the solute is
not incorporated into the solvent crystal lattice as it forms, the solute concentration during the
freezing process increases. This, in turn, causes the freezing point to be lowered, and no
plateau is observed. To determine the freezing point from the graph, the intersection of the
line representing the cooling at the onset, and the line representing the temperature during the
freezing process, is found. The freezing point is determined by extrapolation of this point to
time = 0, and reading the temperature from the graph. As with the cooling curve of a pure
solvent, supercooling may occur.
Experiment 5 Freezing Point Depression
Bergen Community College 41 General Chemistry II Laboratory
Figure 3: Cooling Curve of a Solution
In this experiment, the molar mass of an organic (molecular) solute will determined. First, the
freezing point of pure water will be found by collecting time-temperature data and constructing a
cooling curve. Then, the solute will be dissolved in water, and the freezing point of the solution will
be obtained in a similar manner. From the change in freezing point, the molar mass will be calculated.
REAGENTS: EQUIPMENT:
deionized water large test tubes
table salt thermometers or thermocouples
crushed ice 600 mL beaker
unknown organic compound two 100mL beakers
(glucose, sucrose or urea) magnetic stir bar
magnetic stirrer
rubber stopper
Safety Alert
Wear eye goggles at all times
Wash hands following every experiment
Make sure that lab workbench and glassware used are cleaned after completing experiment
Do not dispose of any chemical without asking your instructor.
Be sure to handle all chemicals, solutions, materials, etc. with proper care
Freezing Point
Time
T em
p er
at u re
Bergen Community College 42 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 5: Freezing Point Depression
PRE-LABORATORY QUESTIONS
1. Pure benzene, C6H6, freezes at 5.45 oC. A solution containing 7.24 g of C2Cl4H2 in 115
g of benzene freezes at 3.55oC. Based on this data, what is the molal freezing-point
constant of benzene, Kf? Molar masses: C = 12.00; H = 1.01; Cl = 35.54
2. Determine the freezing point of an 2.5 m solution of C12H22O11 in water. Kf = -1.86 oC.
3. Given 0.01 m aqueous solutions of each of the following, arrange the solutions in order
of increasing freezing point. Kf = -1.86 oC.
NaCl CaCl2 K3PO4 C6H12O6
Experiment 5 Freezing Point Depression
Bergen Community College 43 General Chemistry II Laboratory
PROCEDURE
Freezing Point of Pure Water 1. Prepare an ice/salt water bath in a 600 mL beaker by adding 100 g of salt, 500 mL of
ice, and 100 mL of tap water. Place a large stirring bar in the beaker and turn on the
magnetic stirrer, in order to mix well. 2. Add 10 grams of deionized water to a large test tube.
3. Place one temperature probe (or thermometer) into the test tube, and one into the
ice/salt water bath. Do not remove them until the experiment is finished.
4. Place the test tube with the water into the bath. Make sure that the water level in the
test tube is fully submerged in the ice/salt water bath, as shown below.
Figure 4. Apparatus to determine the freezing point of water or
freezing point depression of water-soluble compound.
5. Record the temperature of the water in the test tube every 30 seconds until the
temperature has not changed for at least three minutes.
Experiment 5 Freezing Point Depression
Bergen Community College 44 General Chemistry II Laboratory
Freezing Point of Solution with Unknown
1. Prepare an ice/salt water bath in a 600 mL beaker by adding 100 g of salt, 500 mL of
ice, and 100 mL of tap water. Place a large stirring bar in the beaker and turn on the
magnetic stirrer, in order to mix well.
2. Obtain a sample of an unknown nonelectrolyte from the instructor and record the code.
Grind the sample using a mortar and pestle.
3. Tare a large test tube, using a beaker to keep it upright. Add 2 g of the unknown and
record the mass. Tare the test tube again. Add 10 g d-H2O water; record the mass.
4. Mix the contents of the test tube well. Make sure the solid is completely dissolved.
5. Place one temperature probe (or thermometer) into the test tube, and one into the
ice/salt water bath. Do not remove them until the experiment is finished.
6. Place the test tube with the solution containing the unknown into the bath. Make sure
that the whole solution is fully submerged in the ice/salt water bath, as shown on the
previous page.
7. Record the temperature of the water in the test tube every 30 seconds, until
completely solidified.
Disposal: Unknown solutions may be washed down the drain.
CALCULATIONS
1. Prepare a graph in Excel* of temperature as a function of time for both the pure water
and the solution with the unknown.
2. Determine the freezing point of pure water as described on page 41. Use a ruler and a
pencil to extend the line of the plateau to the y-axis to read the freezing point.
3. Determine the freezing point of the solution containing the unknown as described on
page 42. Use a ruler and a pencil to draw best-fit lines through the initial cooling
portion of the graph and the portion for the freezing process. Where these lines
intersect, draw a straight line to the y-axis and read the freezing point.
4. Determine the change in freezing point by the difference in the temperatures.
5. Determine the molality of the solution using Equation 1. Kf is -1.86 oC/m and i = 1.
6. Determine the number of moles of the unknown from the molality and the mass of the
solvent, water.
7. Determine the molar mass of the unknown by dividing the number of grams by the
number of moles.
*See Appendix C for directions on graphing.
Bergen Community College 45 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 5: Freezing Point Depression
DATA: Freezing Point of Pure Water
Time, minutes Temperature, °C Time, minutes Temperature, °C
Bergen Community College 46 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
DATA: Freezing Point of Solution with Unknown
Unknown Code __________ Mass unknown __________ Mass water __________
Time, minutes Temperature, °C Time, minutes Temperature, °C
Bergen Community College 47 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 5: Freezing Point Depression
Results:
Freezing point of pure water: _______________ (from the Excel Graph)
Freezing point of the solution: ______________ (from the Excel Graph)
Freezing point depression (ΔT): ______________
Molality of solution: _________________
Moles of unknown: __________________
Molar mass of unknown: _________________
Show calculations below or on a separate sheet.
Bergen Community College 48 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 5: Freezing Point Depression
POST-LABORATORY QUESTIONS
1. A solution is 2.0 m NaCl. What is the freezing point of the solution? Kf = -1.86
2a. A solution contains 5.0 g of KCl in 200. g water. What is its freezing point? Kf = -1.86
2b. What mass of sucrose is needed to obtain the same freezing point as the above solution?
Bergen Community College 49 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 5: Freezing Point Depression
POST-LABORATORY QUESTIONS (continued)
3. An organic compound contains 74.07% C, 8.64% H and 17.28% N. When 1.921
grams of this compound is dissolved in 48.92 grams of water, the freezing point of
the solution becomes - 0.45oC. molar masses: C = 12.00; H = 1.00; N = 14.00.
a. Determine the empirical formula of the compound.
b. Calculate the molality of the solution. Kf = -1.86 oC.
c. Calculate the molar mass of the solute.
d. Determine the molecular formula of the organic compound.
Bergen Community College 50 General Chemistry II Laboratory
EXPERIMENT 6
Chemical Kinetics
OBJECTIVE: To determine the rate law, rate constant, activation energy, and collision frequency factor for the reaction of iodide with iron(III) ions.
BACKGROUND:
The field of chemical kinetics involves the study of reaction rates. The rate of a reaction
is defined as the change in concentration of reactants or products per unit time. Concentration
is measured in moles per liter, and the time unit varies depending on the speed of the reaction.
Commonly, the time unit is seconds; in this case, the rate is measured in mol/L∙s.
As a reaction proceeds, the concentrations of the reactants decrease, while the
concentrations of the products increase. Because of this, when the rate is measured using the
concentration of a reactant, whose concentration change is negative, the negative of this
change is used to give a positive value for the rate. Since the balancing coefficients may differ
from one reactant or product to another, the changes in concentration may not be the same
when the rate is measured using different species. In order to standardize reaction rates, it is
convention to use the inverse of the balancing coefficient. Thusly, the rates of disappearance
of the reactants and formation, or appearance, of the products are related as follows.
For the hypothetical reaction
a A + b B c C + d D
the rates are given as
− 1
𝑎
∆[A]
∆𝑡 = −
1
𝑏
∆[B]
∆𝑡 = +
1
𝑐
∆[C]
∆𝑡 = +
1
𝑑
∆[D]
∆𝑡
For example, in the gas-phase formation of ammonia, N2 + 3 H2 2 NH3, the relationship
of the rates with respect to changes in concentration of hydrogen, nitrogen, and ammonia are:
− ∆[N2]
∆𝑡 = −
1
3
∆[H2]
∆𝑡 = +
1
2
∆[NH3]
∆𝑡
Thus, if the rate of disappearance of N2 is 0.012 mol/L∙s, the rate of appearance of NH3 is
twice the rate of disappearance of N2, and Δ[NH3]/Δt = 2 × 0.012 mol/L∙s = 0.024 mol/L∙s.
timeinchange
ionconcentratinchange Rate =
Experiment 6 Chemical Kinetics
Bergen Community College 51 General Chemistry II Laboratory
There are many factors that influence the rate of a reaction. The physical state of the
reactants plays a large role. Ionic reactions are almost instantaneous, whereas reactions that
require bond breaking are slow. The presence of a catalyst and the temperature also affect
reaction rate. The latter is investigated in the laboratory exercise following this. The reaction
rate is dependent on concentration as well. A higher concentration of reactants causes more
total collisions to take place in a given amount of time, increasing the frequency of collisions.
For simple reactions that occur in a single step, the rate of the reaction is dependent on the
concentrations of each of the reactants. This relationship is summarized in the rate law, or rate
equation. For the reaction A + B → C occurring in a single step, the rate law is
∆[C]
∆𝑡 = 𝑘[A][B] where k represents the rate constant
A larger value for the rate constant corresponds to a faster reaction. It is important to note that
the rate constant is dependent on temperature.
Other than the simplest types of chemical reactions, however, reactions generally occur
in a series of elementary steps, known as a reaction mechanism. Oftentimes, a product is
formed in the initial step that is then consumed in a later step. These species do not appear in
the overall reaction, and are referred to as intermediates. In the example given below, Q is the
catalyst and QA, Y, and Z are intermediates.
Step 1 A + Q QA + Z (fast)
Step 2 B + QA Y + D (slow)
Step 3 Z + Y Q + C (fast)
The overall reaction: A + B C + D
The energy progress diagram of a multi-step
reaction is given in the figure to the right. The
slowest step in the mechanism has the largest
activation energy. For a multi-step reaction, the
reaction rate depends only on the concentrations
of the reactant species in the slowest step in the
mechanism, called the rate-determining step.
For the multi-step reaction aA + bB cC + dD the rate equation is
Rate = k [A]x [B]y where x is the order with respect to A, y is the order
with respect to B, and k again is the rate constant
Experiment 6 Chemical Kinetics
Bergen Community College 52 General Chemistry II Laboratory
Because the mechanism for a reaction is not evident, and therefor the rate-determining
step is not known, the values of x and y must be determined experimentally. In fact, the kinetics
of a reaction is often investigated specifically for the purpose of identifying a plausible
reaction mechanism.
There are two methods to determine the rate of a reaction: the method of initial rates
and the integral method. In the latter method, the concentration of a reactant or product is
measured over time. From this data, the dependence on concentration, i.e. the rate law, can be
determined. In the method of initial rates, several determinations of rate are performed with
varying initial concentrations of one reactant, while keeping the concentration of the other
reactants the same. Initial concentration is denoted using a subscript of zero, e.g. [A]0.
Comparison of these initial rates allows the order for this reactant to be determined. If the
reaction is zero order with respect to this reactant, changing its concentration does not change
the rate. If it is first order, a change in concentration will cause the same change in the rate. If
it is second order, the rate changes by the square of the change in concentration. The same
process is then done for the other reactants, such that the order for each reactant is determined.
Reaction rates do not only
depend on the concentration of the
reactants, but also depend on
temperature. Molecules have a
higher average kinetic energy at a
higher temperature. Thus, a larger
portion of molecules possess the
minimum energy required to begin
the bond-breaking process at a
higher temperature than at a lower
one. This is shown in the energy
distribution figure to the right.
The minimum energy that
the reactant molecules must
possess in order to overcome the
energy barrier for a reaction to
occur is called the activation
energy. In the energy diagram for
an exothermic reaction to the left,
the activation energy can be seen as
the difference between the energy
of the reactants and that of the
transition state. The relationship
between temperature and activation
energy is given by the Arrhenius
equation, found on the next page.
R el
at iv
e N
u m
b er
o f
M o le
cu le
s
Energy
Required
Energy
Lower Temperature
Higher Temperature
Experiment 6 Chemical Kinetics
Bergen Community College 53 General Chemistry II Laboratory
Arrhenius Equation: k = A e-Ea/(RT) (eq 1)
k is the rate constant, which depends only on temperature
A is the Arrhenius constant, or the collision frequency factor
Ea is the activation energy of the reaction R is the Universal gas constant (8.31 Jmol-1K-1)
T is the Kelvin temperature
Taking the logarithm and rearranging, ln k = ln A e-Ea/(RT)
ln k = ln A + ln e-Ea/(RT)
ln k = ln A - Ea/(RT)
ln k = - Ea/(RT) + ln A
This gives the equation in the slope-
intercept form of a line. Thus, the activation
energy and collision frequency factor can be
obtained by plotting ln k as a function of
reciprocal kelvin temperature, as shown to
the left. The slope is equal to -Ea/R, and the
y-intercept is equal to ln A. The activation
energy can also be calculated algebraically
if the rate constant is known for two
temperatures.
In this experiment, the method of initial rates will be used to determine rate law and
rate constant, and an Arrhenius plot will be constructed to determine activation energy and
collision frequency factor. The reaction that will be studied is that of iron(III) ion with iodide
ion in acidic solution. The iron(III) is reduced to iron(II), and the iodide is oxidized to the
triiodide ion. Acidic conditions are needed to keep the iron(III) ions in the 3+ oxidation state.
2Fe3+(aq) + 3I-(aq) 2Fe2+(aq) + I3 -(aq)
The change in concentration of triiodide will be measured by adding small quantities of both
thiosulfate ions and starch. As the triiodide ions are produced, they are immediately consumed
by thiosulfate.
I3 -(aq) + 2S2O3
2-(aq) 3I-(aq) + S4O6 2-(aq)
When there are no thiosulfate ions remaining, the triiodide ions will combine with the starch
to give a characteristic intense blue color. The time between the mixing of reactants and the
blue color will be measured so that the rate can be determined. Since there are two thiosulfate
ions required to consume one triiodide ion, the change in concentration of triiodide will be
equal to half the change in concentration of thiosulfate. Expression for the reaction rate and
rate constant are given below.
Reaction rate = Δ[I3 -]/Δt = k [Fe+3]x[I-]y and k = reaction rate/[Fe+3]x [I-]y
Arrhenius Plot: ln k vs inverse T
Bergen Community College 54 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
PRE-LABORATORY QUESTIONS
1. Consider the following reaction, which is thought to occur in a single step.
OH ˉ + CH3Br → CH3OH + Brˉ
What is the rate law?
2. Consider the following mechanism:
Step 1 Slow NO2(g) + NO2(g) → NO(g) + NO3(g)
Step 2 Fast NO3(g) + CO(g) → NO2(g) + CO2(g)
a. What is the overall reaction?
b. Are there any intermediates?
c. What is the rate law for this mechanism?
3. What are the units of k for the rate law: rate = k[A][B]2, when concentration is in mol/L?
Experiment 6 Chemical Kinetics
Bergen Community College 55 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
0.050 M KI Three 10-mL graduated cylinders
0.015 M FeCl3 Ten 50-mL Erlenmeyer flasks
0.010 M HNO3 Beral pipettes
0.0040 M Na2S2O3, freshly prepared Timers
2.0% starch solution, freshly prepared Hot water baths
distilled water
Safety Alert
Eye Goggles must be worn at all times, and hands must be washed at the conclusion of the
experiment. Dispose of all the chemicals in the designated container. If unsure, ask the how
or where to dispose of the waste products. Do not dispose of any solutions in the sink.
PROCEDURE
1. Obtain four 10-mL graduated cylinders and label one each as Fe+3, H+, S2O3 -2, and I-.
2. Obtain two 50-mL Erlenmeyer flasks and label them A and B.
3. Measure the volumes of each reactant into the appropriate clean, dry flasks, as specified in
the table on the following page for experiment 1. Swirl to mix.
4. Transfer the contents of flask B into flask A and immediately begin the timer. Pour the
contents back and forth between the flasks several times to mix. Continue to swirl the flask.
Record the time at which the contents turn blue.
5. Measure and record the temperature of the contents of the reaction flask.
6. Repeat steps 3 and 4 using clean, dry flasks for a second trial.
7. Follow the same procedure (steps 3 and 4) for experiments 2 and 3, performing two trials
each, using the reactant volumes specified in the table. (The temperature is not necessary for
these experiments.)
8. Obtain a water bath and adjust the temperature to 35oC for experiment 4.
9. Measure the volumes of each reactant into the appropriate clean, dry flasks, as specified in
the table on the following page for experiment 1, 4, and 5. Swirl to mix.
Experiment 6 Chemical Kinetics
Bergen Community College 56 General Chemistry II Laboratory
10. Immerse both flasks into the water bath. Hold the flasks in place by adding an O-ring to
the neck of each. Place a thermometer in flask A; when the temperature of the solution reaches
the set temperature, transfer the contents of flask B into flask A and immediately begin the
timer. With the flasks still immersed in water, swirl the solution to mix. Maintain the
temperature constant as much as possible. Record the time at which the contents turn blue.
11. Measure and record the temperature of the contents of the reaction flask.
12. Repeat steps 9 through 11 using a 50°C water bath for experiment 5.
Dispose of all solutions in the appropriate waste container.
Solution Volumes:
CALCULATIONS
The order of the reaction for Fe+3 and I- is given by the values of x and y in the rate law:
rate = Δ[I3 -]/Δt = k [Fe+3]x[I-]y
Determination of reaction order, x and y, and rate constant (Experiments 1, 2 & 3):
1. Determine the initial concentrations, [Fe+3]0, [I -]0, and [S2O3
-2]0, using M1V1=M2V2.
2. Determine the concentration changes, Δ[S2O3 -2] and Δ[I3
-]:
Δ[S2O3 -2]: Final concentration (zero) – initial concentration
Δ[I3 -]: From the balancing coefficients, ∆[ I3
-]= -½ ∆[S2O3 -2]
3. Determine the average time to color, Δt, in seconds, for each experiment.
4. Determine the reaction rates for each experiment: r = Δ[I3 -]/Δt
Experiment
Flask A Flask B
0.015 M
FeCl3
0.010 M
HNO3
2% starch
solution 0.050 M KI
0.0040 M
Na2S2O3
1, 4, 5 10.00 mL 5.00 mL 1.0 mL
(20 drops) 5.00 mL 5.00 mL
2 5.00 mL 10.00 mL 1.0 mL
(20 drops) 5.00 mL 5.00 mL
3 5.00 mL 5.00 mL 1.0 mL
(20 drops) 10.00 mL 5.00 mL
Experiment 6 Chemical Kinetics
Bergen Community College 57 General Chemistry II Laboratory
=
5. Determine the ratio of rates for the corresponding experiments.
6. To determine the order of the reaction for iron(III) ion, x:
The ratio of rates is given by
r1 k [Fe +3]x[I-]y
r2 k [Fe +3]x[I-]y
Since k and [I-] do not change in these experiments, and [Fe+3] in experiment 2 is one
half that experiment 1, this is simplifies as
r1/r2 = 2 x Taking the logarithm, log (r1/r2) = log 2
x = x log 2
Thus, the order, x, is given by the following:
𝑥 = log 𝑟1/𝑟2
0.301
Similarly, 𝑦 = log 𝑟3/𝑟2
0.301
7. Round the values of x and y to the nearest integer.
8. Determine the rate constant, k, for each of the three experiments:
k = reaction rate/[Fe+3]x [I-]y
9. Determine the average for the rate constant.
10. Write the Rate Law for the reaction, including values for k, x, and y.
Determination of Activation Energy and Collision Frequency Factor (Experiments 1, 4 & 5):
1. Determine the initial concentrations and concentration changes as in steps 1 and 2 above.
The values are the same for all three experiments.
2. Determine the time change, Δt, in seconds, for each experiment. (Use the average for #1.)
3. Determine the reaction rate for each experiment: r = Δ[I3 -]/Δt.
4. Determine the rate constant, k, using the values of x, and y determined above.
k = reaction rate/[Fe+3]x [I-]y
5. Determine the natural logarithm of the rate constant.
6. Record the Celsius temperature (from data), convert to Kelvin, and find the reciprocal.
7. Prepare an Arrhenius plot in Excel of ln k as a function of reciprocal temperature. See
Appendix C for directions on graphing.
8. Obtain the slope from the equation of the line m. Since m = -Ea/R, then Ea = -mR.
9. Obtain the y-intercept from the equation of the line, b. Since b = ln A, then A = eb.
Bergen Community College 58 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
Data Sheet:
Experiment Time to color change, seconds
Temperature, °C
Trial 1 Trial 2
1
2
----------------
3
----------------
4
----------------
5
----------------
Bergen Community College 59 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
Results: Show calculations on a separate sheet. Determination of Reaction Order and Rate Constant
Experiment 1 Experiment 2 Experiment 3
Initial concentrations: [Fe+3]0 ____________ ____________ ____________
[I-]0 ____________ ____________ ____________
[S2O3 -2]0 ____________
Concentration Changes:
Δ[S2O3 -2] _____________
Δ[I3 -] _____________
Average time to color change, Δt, s ____________ ____________ _____________
Reaction Rate, Δ[I3
-]/Δt, M/s ____________ ____________ _____________
Ratio of rates, r1/r2 ______________
Reaction order for iron(III), x ______________
Order x rounded to nearest integer _______
Ratio of rates, r3/r2 ______________
Reaction order for iodide, y ______________
Order y rounded to nearest integer _______
Rate constant, k ____________ ____________ _____________
Average rate constant (include units) __________________
Bergen Community College 60 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
Results (p2):
Rate Law:
Determination of Activation Energy and Collision Frequency Factor
Initial concentrations: [Fe+3]0 ___________ [I -]0 ___________ [S2O3
-2]0 ___________
Concentration Changes: Δ[S2O3 -2] _____________ Δ[I3
-] _____________
Experiment 1 Experiment 4 Experiment 5
Time to color change, Δt, s ____________ ____________ ____________
Reaction Rate, Δ[I3
-]/Δt, M/s ____________ ____________ ____________
Rate constant, k ____________ ____________ ____________
ln k ____________ ____________ ____________
Temperature, °C ____________ ____________ ____________
Temperature, K ____________ ____________ ____________
1/T, K-1 ____________ ____________ ____________
slope of line,m, (from Arrhenius Plot) ______________
Activation Energy, Ea, J/mol ______________
y-intercept, b, (from Arrhenius Plot) ______________
Collision frequency factor, A, 1/Ms ______________
Bergen Community College 61 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
POST-LABORATORY QUESTIONS
1. Shown below are the results of four separate experiments measuring reaction rate for the
following reaction:
2 A + 3 B + C 2 D + 3 E
Trial [A]o [B]o [C]o Initial rate (M/s)
1 0.01 0.20 0.10 1.67 × 10-4
2 0.02 0.20 0.20 1.33 × 10-3
3 0.02 0.20 0.10 3.33 × 10-4
4 0.01 0.40 0.10 1.67 × 10-4
a. Determine the rate law of the reaction.
b. Determine the rate constant, including the units.
c. What is the rate when the initial concentrations for all reactants is 0.15 M?
Bergen Community College 62 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 6: Chemical Kinetics
POST-LABORATORY QUESTIONS (continued)
2. If 0.20 M Fe3+ had been used instead of 0.020 M Fe3+, how would the numerical value of
the rate constant and the activation energy of the reaction have been affected?
3. If another experiment had been performed in water bath containing ice-water, how would
the rate of the reaction have been affected?
4. With two values of the rate constant, the activation energy can be found using the following
form of the Arrhenius equation:
𝑙𝑛 𝑘1 𝑘2
= −𝐸𝑎
𝑅 (
1
𝑇1 −
1
𝑇2 )
The rate constant for the conversion of cyclopropane to propene is 1.10 x 10-4 s-1 at 470 °C
and 1.02 x 10-3 s-1 at 510 °C. What is the activation energy?
Bergen Community College 63 General Chemistry II Laboratory
Experiment 7
Le Châtelier’s Principle
OBJECTIVE: To investigate how the relative equilibrium quantities of reactants and products are affected by changes in concentration and temperature.
BACKGROUND:
Many chemical reactions are reversible, and therefore do not go to completion. Not
only do reactants change to form products- the forward direction of the reaction, but products
also change to form reactants- the reverse direction. As a reaction proceeds, the concentration
of the reactants decreases, and thus the rate of the forward reaction decreases. However, the
concentration of the products increases, causing the rate of the reverse reaction to increase
from zero. Eventually, the rate of the forward reaction becomes equal to the rate of the reverse
reaction. At this point, a dynamic equilibrium is established, and the concentrations of all
species have no net change over time- provided that the temperature of the system remains
constant. This relationship can be seen in the figure below, which shows the concentrations of
the reactants and products over time for the following homogeneous system (one in which the
reactants and products are in the same phase):
a A + b B c C + d D
A chemical equation for a reversible reaction
uses two arrows, pointing in opposite directions,
as shown. The rate equations of the forward
reaction and reverse reactions are as follows:
forward rate = k1[A] a [B]b
reverse rate = k2[C] c [D]d
At equilibrium, the rates are equal, giving
k1 [A] a [B]b = k2[C]
c [D]d Rearranging,
The ratio of the rate constants is also a constant, which is temperature dependent, as are rate
constants. This is known as the equilibrium constant, Keq, and is expressed in the following
relationship, known as the law of mass action first postulated by Guldberg and Waage in 1864:
ba
dc
2
1
[B][A]
[D][C] =
k
k
ba
dc
[B][A]
[D][C] =eqK
Experiment 7 Le Châtelier’s Principle
Bergen Community College 64 General Chemistry II Laboratory
If a system at equilibrium has a change made to one of the factors that affects the
position of the equilibrium (the relative amounts of reactants and products), the system is no
longer at equilibrium. Therefore, the reaction must proceed in either the forward or the reverse
direction in order to establish a new equilibrium position. The direction in which the reaction
proceeds can be predicted using Le Châtelier’s Principle:
If a system at equilibrium experiences a disturbance, or a change to a factor that affects the
equilibrium, then the reaction proceeds in the direction that counteracts this change.
This concept is similar to that of Newton’s third law: For every action, there is equal reaction
in the opposite direction. The following factors affect the position of the equilibrium:
Temperature: The numerical value of Keq depends on temperature. Thus, when the temperature
of the system is changed, the reaction proceeds in either the forward or reverse direction until
a new equilibrium is established according to the new value of the equilibrium constant. In
exothermic reactions, heat is released to the surroundings in the forward direction, and
absorbed by the system in the reverse direction. If the temperature of such a system is
increased, the system counteracts this disturbance by proceeding in the reverse (endothermic)
direction, forming reactants from products, in order to absorb heat and decrease the
temperature. The scenario is reversed for endothermic reactions; an increase in temperature
causes the equilibrium to respond in the forward direction, forming products from reactants.
Pressure: In equilibrium systems involving gases, pressure changes due to changes in volume
also affect the position of the equilibrium. Pressure is inversely related to volume; when the
volume of the system is decreased, the pressure is increased. Thus, the system will respond in
the direction that then decreases the pressure. Since the pressure is due to the collisions of the
molecules with the container walls, this direction depends on the relative number of moles of
gas in the reactants and the products. The system responds in the direction that gives fewer
moles of gas, reducing the pressure. Conversely, if the volume is increased, the pressure is
decreased, and the system responds in the direction that gives more moles of gas. If the number
of moles of gas in the reactants and products is the same, there is no influence of changes in
pressure on equilibrium.
Concentration: An increase in concentration of a reactant causes the system to proceed in the
forward direction, consuming reactants, thus reducing its concentration. On the other hand, a
decrease in reactant concentration causes the system to proceed in the reverse direction. The
scenario is reversed for products; an increase in product concentration causes the reaction to
proceed in the reverse direction, and a decrease results in the forward reaction occirring.
Catalyst: The presence of a catalyst has no effect on the equilibrium position. It does, however,
bring the system to the equilibrium faster.
In this experiment, five equilibrium systems will be investigated. Results will be
interpreted in terms of Le Châtelier’s Principle.
Bergen Community College 65 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
PRE-LABORATORY QUESTIONS
1. In the chemical industry, ammonia is manufactured by the Haber process according to the
following chemical equation.
N2(g) + 3 H2(g) 2 NH3(g) + heat
This is an exothermic reaction. How can the yield of ammonia production be improved?
2. How do the following changes affect the equilibrium of the following exothermic reaction?
Explain using LeChâtlier’s Principle.
H2(g) + I2(g) 2 HI(g)
a. increasing the temperature
b. removing HI(g)
c. adding H2(g)
d. increasing the pressure by decreasing the volume
Bergen Community College 66 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
PRE-LABORATORY QUESTIONS (continued)
3. Hemoglobin (Hb) and oxygen gas form a complex (HbO2) that carries oxygen throughout
the human body. Unfortunately, carbon monoxide also binds to hemoglobin so that an
equilibrium is established. Carbon monoxide poisoning occurs when the concentration of
HbO2 in the blood is reduced.
HbO2 + CO HbCO + O2
The first aid for a person suffering from carbon monoxide poisoning is to (1) remove them to
an area of fresh air, and (2) administer oxygen. Using the principles of equilibrium, explain
how each of these helps to restore the HbO2 concentration.
Experiment 7 Le Châtelier’s Principle
Bergen Community College 67 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Solid sodium chloride Test tubes
distilled water 250-mL beakers
0.020 M FeCl3 balance
0.10 M KSCN stoppers
0.10 M acetic acid wash bottle
bromothymol blue indicator ice
0.10 M NaOH hot plate
1.0 M NaOH
0.10 M Cr(NO3)2
12 M HCl
6 M HCl
0.10 M AgNO3
Prepared closed test tube containing NO2 gas
Safety Alert
Eye Goggles must be worn at all times, no eating and drinking are allowed in the laboratory,
and hands must be washed at the conclusion of the experiment. 12 M and 6 M HCl solutions
are CORROSIVE and can cause severe burns. Notify the instructor in the event of a spill or
contact with these solutions. Dispose of all the chemicals in the designated container. NO2 gas
is toxic. Be careful when handling this test tube.
PROCEDURE
Part A: Saturated NaCl Solution
1. Place 1-2 gram of NaCl into a test tube using a spatula. Using a wash bottle, add
approximately 5 mL of distilled water. Stopper the test tube and shake vigorously.
2. If all of the NaCl has dissolved, add a small amount more (tip of spatula) and shake the
test tube again. Continue in this fashion until undissolved solid NaCl is visible after shaking.
3. Add a couple of drops of 6 M of HCl solution, watching the solution closely. Record the
observation. Disposal: NaCl solution may be disposed of down the drain.
Part B: Aqueous Co(H2O)6 +2 (pink) and CoCl4
-2 (blue)
1. Add approximately 2 mL Co(NO3)2 solution to a test tube. Observe and record the color.
2. Slowly add 12 M HCl solution to the test tube until the color changes. Observe and record.
3. Add distilled water dropwise until the color changes. Observe and record.
Experiment 7 Le Châtelier’s Principle
Bergen Community College 68 General Chemistry II Laboratory
4. Prepare an ice bath by adding 100 mL of water to ice in a 250-mL beaker. Immerse the
test tube in the ice bath. Record any observation.
5. Add 150 mL water to a 250-mL beaker, place on a hot plate and bring to a boil. Immerse
the test tube in the boiling water. Record any observation.
Disposal: Dispose of cobalt (II) solutions in the appropriate waste container.
Part C: Formation of [Fe(SCN)]+2 (red) from Fe+3 (yellow) and SCN- (colorless)
1. Add approximately 2 mL FeCl3 solution to a test tube. Observe and record the color.
2. Slowly add KSCN solution until the color changes. Observe and record the color.
3. Add AgNO3 solution dropwise to the test tube until a change is seen. Observe and record.
4. Add approximately 2 mL of FeCl3 solution to a second test tube; slowly add KSCN
solution until the color changes.
5. Add 1.0 M NaOH dropwise to the test tube until a change is seen. Observe and record.
Disposal: Dispose of cobalt (II) solutions in the appropriate waste container.
Part D: Dissociation of Acetic Acid
1. Add approximately 2 mL of CH3COOH solution to the test tube and add A couple of
drops of bromthymol blue indicator. Observe and record the color.
2. Add 0.10 M of NaOH solution to the test tube until a color change is seen.
Disposal: Dispose of CH3COOH solutions in the appropriate waste container.
Part E: Dimerization of NO2 (brown) forming N2O4 (colorless)
THIS PART OF THE EXPERIMENT MUST BE RUN IN THE FUME HOOD.
1. Prepare two water baths- one ice bath and one water bath at 80oC.
2. Obtain a test tube containing NO2 gas from the instructor and immerse the test tube in the
ice bath. Observe and record the color.
3. Remove the test tube from the ice bath and immerse into the hot water bath. Observe and
record the color.
4. Remove the test tube from the hot water bath and immerse it again in the ice bath.
Observe and record the color.
Bergen Community College 69 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
Data and Results Sheet: (p 1)
Part A: Saturated NaCl Solution
1. Write the equilibrium equation for the dissolution of NaCl in a saturated solution.
2. Observation upon the addition of 6 M HCl solution:
3. Explain the observation in terms of Le Châtelier’s Principle.
Part B: Aqueous Co(H2O)6
+2 (pink) and CoCl4 -2 (blue)
1. Write the equilibrium equation for the addition of Cl- to an aqueous Co(H2O)6 +2 solution.
2. Initial color of solution:
3. Color upon the addition of HCl solution:
4. Explain the observation in terms of Le Châtelier’s Principle.
Bergen Community College 70 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
Data and Results Sheet: (p 2)
5. Color upon addition of water:
6. Explain the observation in terms of Le Châtelier’s Principle.
7. Observation when placed in ice bath:
8. Explain the observation in terms of Le Châtelier’s Principle.
9. Observation when placed in boiling water bath:
10. Explain the observation in terms of Le Châtelier’s Principle.
Bergen Community College 71 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
Data and Results Sheet: (p 3)
Part C: Formation of aqueous [Fe(SCN)]+2 (red) from Fe+3 (yellow) and SCN- (colorless)
1. Write the equilibrium equation for the addition of SCN- to aqueous Fe+3.
2. Initial color of solution:
3. Color upon the addition of SCN- solution:
4. Explain the observation in terms of Le Châtelier’s Principle.
5. Observations upon addition of AgNO3:
6. Explain the observation in terms of Le Châtelier’s Principle.
7. Observation upon addition of NaOH:
8. Explain the observation in terms of Le Châtelier’s Principle.
Bergen Community College 72 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
Data and Results Sheet: (p 4)
Part D: Dissociation of acetic acid.
1. Write the equilibrium equation for the dissociation of CH3COOH in water.
2. Color upon the addition of bromthymol blue.
3. Color upon the addition of NaOH solution.
4. Explain the observation in terms of Le Châtelier’s Principle. (Bromthymol blue is an
indication that is blue when [H+] is low and yellow when [H+] is high.
Part E: Dimerization of NO2 (brown) forming N2O4 (colorless)
1. Write the equilibrium equation for the dimerization of NO2.
2. Observation upon placement in ice bath:
3. Observation upon placement in hot water bath:
4. Explain the observations in terms of Le Châtelier’s Principle.
5. Write the equilibrium equation including the heat of the reaction.
Bergen Community College 73 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
POST-LABORATORY QUESTIONS
1. Consider the following equilibrium:
Cr2O7 2-(aq) + H2O(l) 2CrO4
2-(aq) + 2H+(aq)
a. Predict the shift in system at equilibrium if a solution of AgNO3 is added dropwise to the
system at equilibrium. Briefly explain.
b. Predict the change in equilibrium if a solution of NaOH is added dropwise to the system at
equilibrium. Briefly explain.
c. Predict the change in equilibrium if the system at equilibrium is diluted by distilled water.
Briefly explain.
Bergen Community College 74 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 7: Le Châtelier’s Principle
POST-LABORATORY QUESTIONS (continued)
2. Consider the following equilibrium:
Mg(OH)2(s) Mg 2+(aq) + 2OH-(aq)
a. Predict the shift in system at equilibrium if a solution of HCl is added dropwise to the
system at equilibrium. Briefly explain.
b. Predict the change in equilibrium if a solution of NaOH is added dropwise to the system at
equilibrium. Briefly explain.
c. Predict the change in equilibrium if the system at equilibrium is diluted by distilled water.
Briefly explain.
Bergen Community College 75 General Chemistry II Laboratory
Experiment 8
Coordination Number
OBJECTIVE: To determine the coordination number of a copper-ammonia complex.
BACKGROUND:
An Arrhenius acid is a substance that, when dissolved in water, produces hydrogen
ions, H+. These hydrogen ions then form hydronium ions, H3O +, when a lone pair of electrons
on a water molecule occupies the empty orbital of a hydrogen ion, forming a covalent bond.
. .
H+ ˸ O — H
|
Arrhenius Acid H
In this sense, the Arrhenius definition of an acid is quite limited. The lone pair of electrons
must come from a water molecule, and the empty orbital is in a hydrogen ion.
The Brønsted-Lowry definition of an acid is more broad, but the essence of the
interaction is the same. A covalent bond is formed when a lone pair of electrons occupies the
empty orbital of a hydrogen ion. Because the lone pair of electrons need not come from a water
molecule, the definition describes the role of the hydrogen ion, rather than the water, and a
Brønsted-Lowry acid is defined as a proton donor. (A hydrogen ion is often referred to as a
proton, as the most common isotope of hydrogen has no neutrons.) The substance supplying
the lone pair of electrons is called the Brønsted-Lowry base.
. .
H+ Brønsted-Lowry Base
Brønsted-Lowry Acid
The Lewis definition of an acid is even less restrictive, but again, the essence of the
interaction is the same. A covalent bond is formed when a lone pair of electrons occupies an
empty orbital. In the Lewis concept, any lone pair of electrons can occupy any empty orbital,
not just that in a hydrogen ion. Thus, the definition describes the role of the electron pair, and
a Lewis acid is defined as an electron pair acceptor. The substance supplying the electron pair
is called the Lewis base.
. .
Lewis Base
Lewis Acid
Experiment 8 Coordination Number
Bergen Community College 76 General Chemistry II Laboratory
One example of Lewis acid-base behavior occurs when transition metal ions (Lewis
acids) bond with Lewis bases called ligands. The resultant species are called complexes, or,
if they carry a net charge, complex ions. Because the covalent bonds in a complex have two
electrons that come from the same atom, rather than one from each atom in the bond, the bonds
are called coordinate covalent bonds. The number of coordinate covalent bonds the central
metal ion makes with the ligands is called the coordination number. Compounds containing
complex ions, and thus coordinate covalent bonds, are called coordination compounds. Some
examples of complex ions and their compounds are listed below. Note that the charge on the
ion is the sum of the charges of the Lewis acids and bases that compose the complex.
Complex Ion
Central
Metal Ion
(Lewis Acid)
Ligands
(Lewis Bases)
Coordination
Number
Coordination
Compound
[PtCl4] ˉ 2 Pt+2 Cl ˉ 4 K2[PtCl4]
[Ni(NH3)6] +2 Ni+2 NH3 6 [Ni(NH3)6]Cl2
[Co(NH3)4Cl2] + Co+3 NH3, Cl ˉ 6 [Co(NH3)4Cl2]Cl
In this experiment, the coordination number for Cu+2 in a copper-ammonia complex
will be determined. The general formula for the complex is [Cu(NH3)x]SO4∙H2O, where x represents the coordination number. The complex will be prepared by reacting concentrated
ammonia with copper(II) sulfate pentahydrate in aqueous solution, and recovered by vacuum
filtration. After drying, it will then be analyzed for copper and ammonia independently. From
this, x, the ratio of ammonia to copper(II), i.e. the coordination number, can be determined.
The copper content will be determined spectrophotometrically. The absorbance of five
standard solutions, as well as a solution made from the complex, will be determined at the
analytical wavelength, 675 nm. As the complex is not water-soluble, it will be dissolved in
nitric acid. Thus, the blank used to calibrate the spectrophotometer will be nitric acid. A Beer’s
Law Plot will be constructed from which the copper(II) concentration in the prepared solution
can be determined. See experiment 3 for information on spectroscopy and Beer’s Law. The
number of moles of copper(II) per gram of the complex will be calculated.
The ammonia content will be determined by titration with standardized hydrochloric
acid solution using methyl orange indicator. One mole of acid neutralizes one mole of
ammonia. From this, the number of moles of ammonia per gram of the complex will be
calculated. To determine the coordination number, the number of moles of ammonia per gram
of the complex is divided by the number of moles of copper(II) per gram of the complex.
Bergen Community College 77 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 8: Coordination Number
PRE-LABORATORY QUESTIONS
1. A 0.862 g sample of transition metal-ammonia complex was titrated with 0.6859 M
HCl, requiring 25.65 mL to reach the end point.
a. How many moles of HCl were used?
b. How many moles of ammonia are in the sample?
c. How many moles of ammonia are there per gram of complex?
2. When 0.828 g of Cu+2 complex was dissolved in 20.0 mL of nitric acid, spectroscopy
showed that the concentration of the solution was 0.175 M.
a. How many moles of Cu+2 are contained in the solution?
b. How many moles of copper(II) are there per gram of the complex?
3. A nickel-ammonia complex was found to contain 0.0259 mol ammonia per gram of
the complex, and 0.00431 mol Ni+2 per gram of the complex. What is the
coordination number?
Bergen Community College 78 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Part A: CuSO4∙5H2O 2 100-mL beakers
14 M NH3 10-mL graduated cylinder
Ice bath Stirring Rod and Rubber policeman
ethanol Vacuum Filtration apparatus (Appendix E)
acetone Paper towel and rubber band
Part B: Cu+2 standard solutions Spectrophotometer, Cuvette, and Kim wipes
1 M HNO3 50-mL Erlenmeyer flask
Part C: Standardized 0.5 M HCl Burette, clamp, ring stand, stir station and stirrer
Methyl orange 2 250-mL Erlenmeyer flasks
SAFETY: CuSO4∙5H2O is toxic and an irritant. Ethanol is toxic and flammable.
Concentrated ammonia is toxic and corrosive. Use extreme caution.
PROCEDURE:
Part A: Preparation of the complex
1. Tare a 100-mL beaker. Add 5 grams CuSO4∙5H2O.
2. Add 5 mL distilled water and stir. Any undissolved solid will dissolve in step 3.
3. In the fume hood, add 8 mL 14 M NH3. Stir for 8 minutes.
4. Add 15-20 mL ethanol. A precipitate should form.
5. Collect the precipitate by vacuum filtration.* Rinse with acetone.
6. Transfer the precipitate to a clean 100-mL beaker.
7. Cover with a paper towel secured with a rubber band. Store as directed.
* See Appendix E for directions on filtration.
Disposal: All solutions are disposed of in the waste container in the hood.
Part B: Determination of copper(II) content
1. Turn on the spectrophotometer and allow to warm up for 10 minutes.
2. Set the wavelength to 675 nm. Calibrate using nitric acid as a blank.
3. Condition a cuvette, and fill 3/4 full with the first standard solution.
4. Measure and record the absorbance per the manufacturer’s instructions.
5. Repeat steps 3 and 4 with the remaining four standard solutions.
6. Tare a 50-mL Erlenmeyer flask. Add 0.5 g of the complex. Record the mass.
Bergen Community College 79 General Chemistry II Laboratory
7. Add 20-mL 1M HNO3 from the communal burette; recording initial and final readings.
8. Stir to dissolve. Line and fill the cuvette. Measure and record the absorbance.
Disposal: All solutions are disposed of in the waste container in the hood.
Part C: Determination of ammonia content
1. Line and fill a burette with the standardized HCl solution. Record the molarity.
2. Tare a 250-mL Erlenmeyer flask. Add 0.8 g of the complex. Record the mass.
3. Add 35-50 mL d-water and stir to dissolve. Add 10 drops methyl orange indicator.
4. Record the initial burette reading.
5. Titrate** to a salmon-pink endpoint. Record the final burette reading.
6. Perform a second trial by refilling the burette and repeating steps 2-5.
**See Appendix D for general instructions on titration.
Disposal: All solutions are disposed of in the waste container in the hood.
CALCULATIONS:
Part B: 1. Determine the concentrations of the standard solutions using M1V1=M2V2.
2. Determine the volume of the solution (assuming no change due to dissolution)
by subtracting the initial buret reading from the final buret reading.
3. Prepare a Beer’s Law Plot for Cu+2. Determine the concentration (molarity)
Cu+2 of from the slope of the plot. See Experiment 3 for information on a Beer’s
Law Plot and Appendix C for directions on graphing.
4. Determine the number of moles of Cu+2 in the spectroscopy sample from the
volume of the solution (data) and the molarity (calculated above).
5. Determine the number of moles of Cu+2 per gram of complex from the number
of moles of Cu+2 (calculated above) and the mass of the complex (data).
Part C: 1. For each titration, determine the number of moles of HCl used in the titration
from volume of the HCl solution used (data) and its molarity (data).
2. Again, for each titration, determine the moles of NH3 from the moles of HCl
(calculated above) and the mole ratio for the reaction- 1 mol HCl to 1 mol NH3.
3. Again, for each titration, determine the number of moles of ammonia per gram
of complex from the number of moles (calculated above) and the mass (data).
4. Determine the average number of moles of ammonia per gram of complex from
the values calculated above.
Part D: Divide the number of moles of ammonia per gram of complex (Part C) by the
number of moles of copper(II) per gram of complex (Part B) to get the ratio.
Bergen Community College 80 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 8: Coordination Number
DATA and Results:
Part B: Determination of copper(II) content Initial buret reading: _____________
Mass of complex: ______________ Final buret reading: _____________
Volume of nitric acid: _____________
Sample#
Volume (mL)
of
0.165 M Cu2+
Volume (mL)
of
1.00M HNO3
Total Volume
(mL) Molarity (M)
Absorbance
(A)
Blank --- --- ---
1 20.0 35.9
2 20.0 26.5
3 20.0 19.8
4 20.0 14.7
5 20.0 10.0
Complex --- --- --- ---
Part C: Determination of ammonia content Molarity of HCl solution: ______________
Titration #1 Titration #2 Titration #3
if necessary
Mass of copper-ammonia complex used, g
Final buret reading, mL
Initial buret reading, mL
Volume of HCl used, mL
Moles of NH3, mol
Bergen Community College 81 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 8: Coordination Number
Results: Show calculations on a separate sheet.
Part B:
Molarity of Cu+2 solution (from Beer’s Law Plot) ___________________
Number of moles of Cu+2 ___________________
Number of moles of Cu+2 per gram of complex ___________________
Part C: Titration 1 Titration 2
Number of moles of NH3 ______________ ______________
Number of moles of NH3 per gram of complex ______________ ______________
Average number of moles NH3 per gram of complex __________________
Part D:
Ratio of moles of NH3 to moles Cu +2, including significant digits ____________________
Ratio, rounded to nearest whole number ____________________
Formula of complex ________________________________
Bergen Community College 82 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 8: Coordination Number
POST-LABORATORY QUESTIONS
1. Consider the complex ion [Pt(NH3)4Cl2] +2. Note: Include charges for any ions.
a. Identify the Lewis acid(s). ____________________
b. Identify the Lewis base(s). ____________________
c. What is the coordination number? ____________________
2. Consider the analysis of the copper content in this experiment.
If the cuvets used in the spectrophotometer were rinsed with distilled water, but not
dried or conditioned (lined) before use, would the calculated number of moles of
copper per gram of complex be too high, too low, or unaffected? Briefly explain.
3. Consider the analysis of the ammonia content in this experiment.
If the buret used in the titration was rinsed with distilled water, but not dried or
conditioned before use, would the calculated number of moles of ammonia per gram
of complex be too high, too low, or unaffected? Briefly Explain.
If the titration flask was rinsed with distilled water but not dried before use, would the
calculated number of moles of ammonia per gram of complex be too high, too low, or
unaffected? Briefly Explain.
Bergen Community College 83 General Chemistry II Laboratory
Experiment 9
Identification of a Weak Acid
OBJECTIVE: To identify an unknown weak acid by its molar mass and pKa value.
BACKGROUND:
The field of analytical chemistry involves the determination of the amount of a
substance present in a mixture by allowing the substance to react with a known amount of
another substance. From the stoichiometry of the reaction, the amount of the unknown can
then be found. In a titration, a solution containing one reagent, the titrant, is slowly added to
the other. A process or device is used to signal that the reaction is complete. There must be a
rapid quantitative reaction taking place as the titrant is added. Although many types of
chemical reactions are used in titration, acid-base (neutralization) titrations are quite common.
This type of titration can also be used to identify a weak acid by allowing the determination
of the acid’s equivalent mass and pKa value. The equivalent mass of an acid is the mass
required to supply one mole of hydrogen ions. For a monoprotic acid, it is equal to its molar
mass. The pKa value of an acid is the negative logarithm of the acid dissociation constant, Ka.
The neutralization reaction is one in which an acid and a base react to produce salt and
water. During this process, a hydrogen ion, H+ (or hydronium ion, H3O +), from the acid and a
hydroxide ion, OH−, from the base form a water molecule, H2O. A salt is also formed from
the anion of the acid and the cation of the base. Neutralization reactions are generally
exothermic, releasing heat to the surroundings as the covalent bond in water is formed.
HX(aq) + MOH(aq) → MX(aq) + H2O(l)
A visual indicator, usually an alcohol-soluble organic compound, is used in acid-base
titrations to signal the end of the reaction. Its color is dependent on the pH of the solution.
Phenolphthalein, a commonly used visual indicator, is pink at higher pH and colorless at lower
pH, changing in the pH range of 8 – 10. The point in a titration at which the indicator changes
color is called the endpoint. A pH meter can also be used to identify the completion of an acid-
base reaction. This figure shows the titration curve (pH as a function of titrant volume) for the
titration of a strong acid (e.g. HCl) with a strong
base (e.g. NaOH). The equivalence point occurs
when a chemically equivalent amount of OH‾
has been added to the acid, and can be identified
by the inflection point, the point at which the
slope is a maximum. Here, the equivalence point
occurs at a pH of 7. At the beginning of the
titration, the pH changes very slowly, and
reflects the concentration of the free strong acid.
After the equivalence point, the pH reflects the
concentration of excess strong base.
Experiment 9 Identification of a Weak Acid
Bergen Community College 84 General Chemistry II Laboratory
The titration curve of a weak acid with a strong base differs from that of a strong acid
with a strong base. First, a jump is observed at the beginning of the titration which
subsequently levels off. Secondly, the pH of the equivalence point is higher than 7.0. The
titration curve of a weak acid with a strong base is shown below. At the half-equivalence point
(the point at which half of the acid has been converted to the conjugate base) the pH of the
solution is equal to the pKa of the acid. The presence of these two species (the acid and its
conjugate base) form a buffer, and the pH can be calculated by the Henderson-Hasselbalch
equation. The titration curve for such a titration is shown below.
Henderson-Hasselbalch
Equation:
pH = pKa + log [CB]
[Acid]
Note: An alternate plot to find the
equivalence point is the first
derivative (ΔpH/ΔV) of the data .
The slope of the titration curve is
steepest at the equivalence point,
and the first derivative plot shows
a maximum.
The stoichiometry is the same as described above in case of a titration of a strong acid with a
strong base;one mol of NaOH is required to neutralize one mol of monoprotic weak acid, HA.
HA(aq) + NaOH(aq) → NaA(aq) + H2O(l)
In the titration of a diprotic acid, pictured
to the left, two inflection points are observed on
the pH curve. The first is used to determine the
dissociation constant of the first ionization, pKa1,
while the second inflection point is used to
determine pKa2.
In this experiment, an unknown weak
acid will be identified from the molar mass and
the pKa value(s) by comparison to the acids
listed in the table on the following page. The pKa
value(s) will be obtained from the titration curve,
and the molar mass will be determined from the
stoichiometry of the reaction.
Experiment 9 Identification of a Weak Acid
Bergen Community College 85 General Chemistry II Laboratory
Common Weak Acids
Common
name
IUPAC name Chemical
formula
Molar
Mass g/mol
pKa1 pKa2
Oxalic acid ethanedioic
acid
Dihydrate
(COOH)2
(COOH)2∙2H2O
90
126
1.27 4.27
Malonic acid propanedioic
acid
HOOC-
(CH2)-
COOH
104 2.85 5.70
Succinic acid butanedioic
acid
HOOC-
(CH2)2-
COOH
118 4.21 5.41
Adipic acid hexanedioic
acid
HOOC-
(CH2)4-
COOH
146 4.41 5.41
Acetic acid Ethanoic acid CH3COOH 60 4.75
Chloroacetic
acid
Chloroethanoic
acid
ClCH2COOH 94.5 2.87
Fluoroacetic
acid
Fluoroethanoic
acid
FCH2COOH 78.0 2.60
Sodium
bisulfite
Sodium
hydrogen
sulfite
NaHSO3 104 6.97
Potassium
hydrogen
phthalate
Potassium
hydrogen
phthalate
C8H5KO4 204.2 5.70
Bergen Community College 86 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
PRE-LABORATORY QUESTIONS
A student titrated 0.180 g of an unknown acid with 0.10 M NaOH solution and the following
data set was obtained:
mL of 0.1M
NaOH pH
mL of 0.1M
NaOH pH
0 2.20 20 6.32
2 3.25 22 6.70
4 4.00 24 7.00
6 4.25 26 9.00
8 4.50 28 11.20
10 4.81 30 11.70
12 5.25 32 12.00
14 5.55 34 12.22
16 5.75 36 12.40
18 6.10 38 12.60
1. Prepare a titration curve using Microsoft Excel. See Appendix C for directions on
graphing.
2. Is this acid monoprotic or diprotic? Briefly explain.
3. Is this acid strong or weak? Briefly explain.
4. Determine the molar mass of the unknown acid.
5. Determine the pKa(s) of the acid.
6. Identify the acid using the table on the previous page.
Experiment 9 Identification of a weak Acid
Bergen Community College 87 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Unknown acid (monoprotic or diprotic) Buret
NaOH solution Stir station and magnetic stirrer
Phenolphthalein Ring stand and clamp
potassium hydrogen phthalate (KHP), dried pH meter
distilled water 25-mL graduated cylinder
Buffer solutions (pH = 4 and pH = 7) 2 125-mL Erlenmeyer flasks
2 150-mL beakers
Safety Precautions
NaOH is corrosive and toxic. Handle with care. Wear protective eye goggles at all times
throughout the experiment. DO NOT FIL THE BURET ABOVE THE EYE LEVEL. In case
of an emergency, have an eye wash available. Avoid skin contact with solids and solutions.
Rinse spills with large amounts of water. Dispose of all solutions into the designated waste
containers located in the fume hoods. Wash hands before leaving the laboratory.
First aid
In case of eye contact, immediately wash eyes with large amounts of water and notify the
instructor immediately. In case of skin contact, immediately flush the contaminated skin
with water for at least 15 minutes. Do not rub the skin. Notify the instructor.
PROCEDURE:
Part A: Standardization of NaOH
1. Tare a 125-mL Erlenmeyer flask. Add approximately 0.25 g KHP. Weigh and record.
2. Add approximately 25 mL of distilled water and 2-3 drops of phenolphthalein.
3. Record the initial buret reading, and titrate* to a pale pink that persists for 15 seconds.
4. Record the final buret reading.
5. Repeat steps 1 through 4 for a second trial.
*See appendix D for general instructions for titration.
Part B: Identification of Unknown Acid
1. Calibrate the pH meter according to the instructor.
2. Tare a 150-mL beaker and add 0.25 – 0.50 g of the unknown acid. Weigh and record.
3. Add 25-30 mL of distilled water and dissolve using a magnetic stirrer and stir station.
4. Place the pH probe in the beaker and secure in place with a clamp. Record the initial pH.
Experiment 9 Identification of a weak Acid
Bergen Community College 88 General Chemistry II Laboratory
5. Refill the buret with the standardized NaOH solution.
6. Add 0.50 mL of the NaOH solution to the beaker and record the pH.
7. Continue in this fashion until a pH of 4.0 is reached.
8. Add 0.20 mL of NaOH; record the pH. Continue in this fashion to a pH of 8.0.
9. Add 0.50 mL NaOH; record the pH. Continue in this fashion until to a pH of 12.
10. Prepare a titration curve in Microsoft Excel;** determine the acid’s molar mass and pKa.
11. Repeat steps 2 to 10 for a second trial.
12. Compare the molar mass and the pKa determined in the first trial. If the values are not in
close proximity, repeat the experiment for a third time with another sample of acid.
Dispose of all solutions in the appropriate container.
CALCULATIONS:
Part A: Standardization of NaOH
1. For each titration, determine the number of moles of KHP used. The molar mass of KHP
is 204.2 g/mol.
2. Since KHP is a monoprotic acid, the mole ratio for the reaction is one mole acid to one
mole NaOH. Thus, the number of moles a NaOH is equal to that of the KHP.
3. Determine the molarity of the NaOH solution by dividing the number of moles by the
number of liters used in the titration.
4. Determine the average molarity.
Part B: Identification of Unknown Weak Acid
For each titration:
1. Prepare a titration curve (pH as a function of titrant volume) in Microsoft Excel.**
2. Determine by the shape of the titration curve whether the acid is monoprotic or diprotic.
3. Locate and mark the equivalence point on the graph, and determine the corresponding
NaOH volume. (For a diprotic acid, use the second inflection point.)
4. Determine the number of moles of NaOH at the equivalence point using the molarity
obtained in Part A.
5. Determine the number of moles of the acid: If monoprotic, it is equal to the number of
moles of NaOH. If diprotic, it is equal to half the number of moles of NaOH.
6. Determine the molar mass by dividing the number of grams by the number of moles.
7. Determine the pKa value(s) by reading and marking the pH at half of the volume at the
inflection point.
8. Compare the molar mass and pKa value(s) to those in the table and identify the acid.
**See Appendix C for insructions on graphing.
Bergen Community College 89 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
Part A: Standardization of NaOH
Trial 3,
DATA: Trial 1 Trial 2 if necessary
Mass of the KHP, g ____________ ____________ ____________
Initial buret reading, mL ____________ ____________ ____________
Final buret reading, mL ____________ ____________ ____________
Volume of NaOH, mL ____________ ____________ ____________
Results:
Show sample calculation below or on a separate sheet.
Mole of KHP used ____________ ____________ ____________
Molarity of NaOH ____________ ____________ ____________
Average Molarity of NaOH _______________
Bergen Community College 90 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
Part B: Identification of Unknown
DATA: Trial 1
Volume NaOH, mL pH Volume NaOH, mL pH
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
Bergen Community College 91 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
Part B: Identification of Unknown
DATA: Trial 2
Volume NaOH, mL pH Volume NaOH, mL pH
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
________ _____ ________ _____
Bergen Community College 92 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
Part B: Identification of Unknown
Results: Show sample calculations below or on a separate sheet.
Type of acid (monoprotic or diprotic) ____________________________
Trial 1 Trial 2
Volume of NaOH at equivalence point ____________ ____________
Number of moles NaOH at equivalence point ____________ ____________
Number of moles of unknown acid ____________ ____________
Molar mass of unknown acid ____________ ____________
pKa value(s) of unknown acid ____________ ____________
Identity of unknown acid __________________
Bergen Community College 93 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 9: Identification of a Weak Acid
POST-LABORATORY QUESTION
The following titration data is provided for the titration of a weak acid with a strong base.
Prepare a titration curve on the graph paper below. (See Appendix C for instructions on
graphing.) Determine the pKa and Ka of the acid.
Volume of NaOH, mL pH Volume of NaOH, mL pH Volume of NaOH, mL pH
0.00 3.02 8.00 5.75 11.50 11.30
2.00 4.14 8.50 6.20 11.80 11.40
3.00 4.34 9.00 7.35 12.00 11.50
4.00 4.60 9.50 8.05 12.50 11.60
5.00 4.85 10.00 8.71 12.80 11.70
5.50 5.00 10.30 10.50 13.00 11.75
6.00 5.25 10.50 10.75 13.50 11.85
6.50 5.45 10.80 10.90 14.00 11.90
7.00 5.55 11.00 11.00
Bergen Community College 94 General Chemistry II Laboratory
Experiment 10
Solubility Product
OBJECTIVE: To determine the solubility product constant for potassium hydrogen tartrate and investigate the effect on solubility of a common ion.
BACKGROUND:
Although some salts are said to be insoluble in water, they do dissolve to a very slight
extent. A dynamic equilibrium is formed between the solid salt and its aqueous ions when the
rate of dissolution of the solid and the rate of precipitation of the ions are equal. Such a solution
is said to be saturated, and the concentration of the salt in this saturated solution is referred to
as its solubility.
Figure 1. Solution equilibrium in a heterogeneous system.
A saturated solution of an insoluble salt is represented by a net ionic equilibrium
equation, as shown in the following examples for silver chloride and lead(II) bromide.
AgCl(s) Ag + (aq) + Clˉ(aq)
PbCl2(s) Pb +2
(aq) + 2 Clˉ(aq)
Equilibrium constant expressions can be written for these equilibrium systems of
insoluble salts. Because the reactant is in the solid phase, there is no factor in the denominator
of the expression. This means that the equilibrium constant is the product of the molar
concentrations of the aqueous ions, raised to their balancing coefficients. It is accordingly
called the solubility product constant and given the symbol Ksp. Expressions for the solubility
product constants for the salts in the above examples are as follows.
Ksp = [Ag+][Clˉ]
Ksp = [Pb+2][Brˉ]2
Experiment 10 Solubility Product
Bergen Community College 95 General Chemistry II Laboratory
The value of a solubility product can be determined experimentally by measuring the
concentration of one of the ions in a saturated solution. The concentration of the counter ion can
be found from the stoichiometry of the salt. Concentration can be determined using various
methods, such as spectroscopy for colored species, or pH for insoluble hydroxides.
In this experiment, the molar solubility and the solubility product constant for
potassium hydrogen tartrate (KHC4H4O6), often called KHT, will be determined. The
hydrogen tartrate ion, HC4H4O6ˉ, is a monoprotic weak acid, and its concentration can be
obtained by titration with sodium hydroxide. The effect on the solubility of KHT caused by
the presence of other salts, specifically potassium nitrate and sodium nitrate, will also be
investigated. Le Chatelier’s principle will be used to interpret these results.
REAGENTS: EQUIPMENT:
KHC4H4O6 (KHT) Three 250-mL Erlenmeyer flasks
0.10 M KNO3 100-mL graduated cylinder
0.10 M NaNO3 Gravity Filtration Apparatus and 250-mL beakers
0.10 M NaOH, standardized 25-mL volumetric pipet
Phenolphthalein Titration Apparatus† and 125-mL Erlenmeyer flasks
SAFETY:
NaOH is corrosive and toxic. Handle with caution.
Wear eye goggles at all times. No drinking, eating, and use of cell phones are allowed in
chemistry laboratories.
Wash your hand before leaving the laboratory.
MSDS are available.
Disposal: Solutions may be disposed of down the drain.
Bergen Community College 96 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 10: Solubility Product
PRE-LABORATORY QUESTIONS
1. Write the net ionic equilibrium equation for the dissolution of KHC4H4O6 in water.
2. Write the solubility product constant expression for KHC4H4O6.
3. Why is it not necessary to record the exact mass of KHT placed into the flasks?
4. The solubility of calcium oxalate, CaC2O4, is 4.8×10
-5 mol/L. What is its Ksp?
5. The solubility product constant for MgF2 is 3.4×10
-11. What is its solubility?
Experiment 10 Solubility Product
Bergen Community College 97 General Chemistry II Laboratory
PROCEDURE:
1. Add 1 gram of KHT to each of three labeled 250-mL Erlenmeyer flasks.
2. Add 100 mL of distilled H2O the first, 100 mL 0.10 M KNO3 to the second, and
100 mL 0.10 M NaNO3 to the last.
3. Swirl, or stir, the contents of the flasks for twenty minutes, being certain that excess
solid KHT remains for the entire period. Add additional KHT if needed.
4. Filter* the first solution directly into a funnel lined with dry filter paper into a clean, dry 250-mL beaker.
5. Transfer one aliquot of the filtered solution to a 125-mL Erlenmeyer flask using a
25-mL volumetric pipet. ‡
6. Titrate† with sodium hydroxide, using 2 drops of phenolphthalein indicator.
7. Perform a second trial.
8. Repeat steps 4 to 7 for the other two solutions.
Disposal: Solutions may be disposed of down the drain.
* See appendix E for directions for filtration. ‡ See appendix B for directions on using volumetric glassware. † See appendix D for directions for titration.
CALCULATIONS:
For each of the three solutions:
1. Determine the number of moles of NaOH used from the molarity and average volume.
2. Determine the number of moles of HC4H4O6¯ from the acid-base mole ratio.
3. Determine the molar solubility of HC4H4O6¯, and thus KHT, in the saturated solutions.
For the solution made from distilled water only:
4. Determine Ksp for KHT from its molar solubility in water, calculated above.
Bergen Community College 98 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 10: Solubility Product
Data:
Molarity of NaOH: __________ Volume of KHT solution titrated: __________
Titration Data: Flask 1 Flask 2 Flask 3
KHT in d-H2O KHT in KNO3 KHT in NaNO3
Trial 1
Initial buret reading ____________ ____________ ____________
Final buret reading ____________ ____________ ____________
Volume NaOH used ____________ ____________ ____________
Trial 2
Initial buret reading ____________ ____________ ____________
Final buret reading ____________ ____________ ____________
Volume NaOH used ____________ ____________ ____________
Trial 3 (if needed)
Initial buret reading ____________ ____________ ____________
Final buret reading ____________ ____________ ____________
Volume NaOH used ____________ ____________ ____________
Average volume
NaOH used ____________ ____________ ____________
Results: Show calculations on a separate sheet.
Solubility of KHT ____________ ____________ ____________
Ksp of KHT ____________
Bergen Community College 99 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 10: Solubility Product
POST-LABORATORY QUESTIONS
1. If the undissolved KHT is not filtered out before the titration with NaOH, will the
calculated value of Ksp be too high, too low, or unaffected? Briefly explain.
2. If the buret is rinsed with distilled water but not conditioned with NaOH, will the
calculated value of Ksp be too high, too low, or unaffected? Briefly explain.
3. If the titration flask is rinsed with distilled water and not dried, will the calculated
value of Ksp be too high, too low, or unaffected? Briefly explain.
4. The usual procedure in a gravity filtration calls for the filter paper to be wet with
distilled water before filtering. Why was the filter paper dry for this procedure?
Bergen Community College 100 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 10: Solubility Product
POST-LABORATORY QUESTIONS (continued)
5. Explain the effect of the presence of potassium nitrate and sodium nitrate on the
solubility of KHT using Le Chatelier’s principle. (Begin by writing the chemical
equation for the equilibrium. Consider only the magnitude of the solublility.)
6. A laboratory experiment designed to determine the solubility product constant for
magnesium hydroxide involved a similar procedure to establish a saturated solution.
The hydroxide ion concentration was then determined using a pH meter. The pH was
recorded as 10.52. Determine Ksp.
Bergen Community College 101 General Chemistry II Laboratory
Experiment 11
Qualitative Analysis of Cations
OBJECTIVE: To analyze an unknown solution for the presence of six cations.
BACKGROUND:
Qualitative analysis involves the determination of whether or not a chemical species is
present in a sample. This scheme includes the analysis of six cations: barium ion, calcium ion,
copper(II) ion, iron(III) ion, manganese(II) ion, and nickel(II) ion. In order to correctly identify
the presence of the cations, the solution to be analyzed is separated into three cation groups,
based on the solubility of their salts, in a process known as selective precipitation. Insoluble
salts (precipitates) are separated from the soluble ions by centrifugation- rapid spinning of the
sample to compact the solid in the bottom of the tube- and decantation- pouring off the liquid
to leave the solid behind.
The first step is to separate the hydroxide group, which includes Fe3+ and Mn2+ ions,
as their hydroxide salts are insoluble. Once separated, nitric acid is added to dissolve
precipitates. The solution is tested for iron(III) ion by the addition of potassium thiocyanate;
a deep red color indicates the presence of Fe3+. The solution is also tested for manganese(II)
ion by the addition of sodium bismuthate; a deep purple color that may fade with time indicates
the presence of Mn2+. The net ionic equations for these reactions are as follows:
Fe3+(aq) + SCN ˉ(aq) → Fe(SCN)2+(aq, red)
2 Mn2+(aq) + 5 BiO3ˉ(aq) + 14 H +
(aq) → 2 MnO4ˉ(aq, purple) + 5 Bi3+(aq) + 7 H2O(l)
Next, the carbonate group is separated. Both Ba2+ and Ca2+ form insoluble carbonates.
Once separated, the precipitates are dissolved in hydrochloric acid. The solution is tested for
barium ion by the addition of sodium sulfate; the formation of a white precipitate indicates the
presence of Ba2+. The solution is then tested for calcium ion by the addition of ammonium
oxalate; a white precipitate indicates the presence of Ca2+. The equations are as follows:
Ba2+(aq) + SO4 2ˉ(aq) → BaSO4(s)
Ca2+(aq) + C2O4 2ˉ(aq) → CaC2O4(s)
The remaining ions, Cu2+ and Ni2+, comprise the soluble group. After heating to
dryness and re-dissolving in phosphoric acid and sodium hydrogen phosphate, the solution is
tested for copper(II) ion by the addition of potassium iodide; the formation of a light tan
precipitate and reddish brown solution indicates the presence of Cu2+. The solution is also
tested for nickel(II) ion by the addition of dimethylglyoxime (HDMG); a red precipitate
indicates the presence of Ni2+. The net ionic equations for the reactions are as follows:
2 Cu2+ (aq) + 4 I ˉ (aq) → 2 CuI (s, light tan) + I2 (aq) I2 (aq) + I ˉ (aq) → I3ˉ (aq, reddish brown)
Ni2+ (aq) + 2 HDMG (aq) → Ni(DMG)2 (s, red) + 2 H+ (aq)
Experiment 11 Qualitative Analysis of Cations
Bergen Community College 102 General Chemistry II Laboratory
In order to be certain that all of the intended ions have been removed by precipitation,
the sample is tested for completeness of precipitation by the addition of additional reagent. In
order to be certain that all of the soluble ions have been removed from the precipitate, the
precipitates are washed by adding distilled water and mixing, followed by centrifugation and
decantation. In this experiment, a solution containing all of the six cations will first be tested. Then, the
scheme will be used to identify which of these ions are present in an unknown solution.
Flow chart for the separation and identification of six cations:
Ba2+, Ca2+, Cu2+, Fe3+, Mn2+, Ni2+
NaOH, HNO3
Precipitate Liquid
Fe(OH)3, Mn(OH)2 Ba 2+, Ca2+, Cu2+, Ni2+
Wash, HNO3 (NH4)2CO3
Precipitate Liquid
Fe3+, Mn2+ BaCO3, CaCO3 Cu 2+, Ni2+
Heat to dryness
HCl H3PO4, Na2HPO4
KSCN NaBiO3
Ba2+, Ca2+ Cu2+, Ni2+
Fe(SCN)2+ MnO4ˉ
(dark red) (dark purple) Na2SO4
Fe3+ ✓ Mn2+✓
Precipitate Liquid
NH3,
(NH4)2C2O4
BaSO4 CaC2O4
(white ppt) (white ppt)
Ba2+✓ Ca2+✓
KI HDMG
CuI (tan ppt), Ni(DMG)2
I3ˉ (red-brown) (red ppt)
Cu2+✓ Ni2+✓
Bergen Community College 103 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 11: Qualitative Analysis of Cations
PRE-LABORATORY QUESTIONS
1. Briefly explain the meaning of the following terms in the context of this laboratory exercise.
centrifugation-
decantation-
2. What is the purpose of checking for completeness of precipitation? Briefly explain how
this is done.
3. What is the purpose of washing a precipitate? Briefly explain how this is done.
4. In addition to the step-by-step instructions, the flow chart found on the previous page in
quite useful in keeping track of the cations during the experiment. Sketch the flow chart so
that the scheme is familiar before beginning the exercise.
Experiment 11 Qualitative Analysis of Cations
Bergen Community College 104 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
Known solution distilled water dropper bottles Test tubes
6M NH3 1M Na2SO4 Test tube rack
6M HNO3 0.30M (NH4)2C2O4 Glass stirring rods
1M (NH4)2CO3 1M H3PO4 Centrifuge
KSCN 1M Na2HPO4 50-mL beaker
NaBiO3 KI Hot plate
6M HCl HDMG Blue litmus paper
Unknown solutions
SAFETY:
6M NH3, 0.3M (NH4)2C2O4 and 6M HNO3 are toxic and corrosive. 1M H3PO4 is corrosive.
6M HCl can cause severe burns. KSCN and NaBiO3 are toxic. MSDS are available.
Follow instructions for the safe use of a centrifuge.
Wear eye goggles at all times. No drinking, eating, or cell phone use is permitted.
Wash hands before leaving the laboratory.
PROCEDURE:
General instructions: Label all test tubes. Use a glass stirring rod to mix solutions. Be careful
not to contaminate solutions; always use clean test tubes, and rinse stirring rods with distilled
water after each use.
Use of centrifuge: Use only test tubes with a rim. Do not use chipped test tubes. The centrifuge
must be balanced with another test tube in the opposite position.
To check for completeness of precipitation: Add one drop of the precipitating reagent to the
solution. If additional precipitate forms, as evidenced by a cloudy appearance, centrifuge again
for two minutes. Continue this procedure until no additional precipitate forms.
To wash a precipitate: Add 10 drops distilled water and mix. Centrifuge for three minutes.
Decant and discard the wash liquid.
Separating cations:
1. Place 10 drops of the known solution (containing all six cations) in a test tube.
2. In the fume hood, add 15 drops 6M NH3 and 2 drops 6M HNO3. Mix. Make and record
observations.
3. Centrifuge for five minutes, check for completeness of precipitation, and decant into a clean
test tube. Label the test tube with the precipitate and save it for step 6.
Experiment 11 Qualitative Analysis of Cations
Bergen Community College 105 General Chemistry II Laboratory
4. To the test tube containing the liquid from step 3, add 20 drops 1M (NH4)2CO3. Mix. Make
and record observations.
5. Centrifuge for two minutes, check for completeness of precipitation, and decant into a 50-
mL beaker. Label the test tube with the precipitate and save for step 9. Heat the beaker to
dryness on a hot plate, about 20-30 minutes, then cool. Keep an eye on it while continuing.
Hydroxide group:
6. Wash the solid in the test tube from step 3, and add 20 drops distilled water. Place a piece
of blue litmus paper on the lab bench. Add 6M HNO3 to the test tube dropwise. After each
addition, stir the solution and then touch the stirring rod to the litmus paper. Continue until the
solution tests acidic to litmus by producing a pink color on the paper. If the precipitate has not
dissolved, add additional 6M HNO3 dropwise until it dissolves. Pour half of the solution into
a clean test tube.
7. To one of the tubes from step 6, add a small pea-sized amount of solid KSCN. Stir. Make
and record observations.
8. To the second test tube from step 6, add a small pea-sized amount of solid NaBiO3. Stir for
one minute. Centrifuge for two minutes. If no color change is observed, add 5 drops 6M HNO3
and stir. Make and record observations.
Carbonate group:
9. Add 10 drops distilled water to the solid from step 5. Add 6M HCl dropwise while stirring
until the precipitate dissolves. Add 5 drops 1M Na2SO4. Mix. Make and record observations.
10. Centrifuge for two minutes and decant into a clean test tube. Add 3 drops 6M NH3 and 10
drops 0.3 M (NH4)2C2O4. Mix. Make and record observations.
Soluble group:
11. Add 20 drops 1M H3PO4 and 20 drops 1M Na2HPO4 to the cooled beaker from step 5.
Divide the solution into two test tubes.
12. To one of the test tubes from step 11, and a small amount of solid KI. Mix. Wait two
minutes. Make and record observations.
13. To the other test tube from step 11, add ten drops of HDMG and 8 drops Na2HPO4. Make
and record observations.
Analyzing the unknown:
14. Obtain an unknown from the instructor. Record the identification code. Repeat the scheme
using the unknown solution. In this case, not all of the cations will be present. When in doubt,
compare the results to those obtained in the analysis of the known solution.
Bergen Community College 106 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 11: Qualitative Analysis of Cations
Data: Analyzing the Known
1. Observation after the addition of NH3 and HNO3:
2. Observation after the addition of (NH4)2CO3:
3. Observation after the addition of KSCN:
4. Observation after the addition of NaBiO3:
5. Observation after the addition of Na2SO4:
6. Observation after the addition of (NH4)2C2O4:
7. Observation after the addition of KI:
8. Observation after the addition of HDMG and Na2HPO4:
Bergen Community College 107 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 11: Qualitative Analysis of Cations
Data: Analyzing the Unknown Identification Code: __________
1. Observation after the addition of NH3 and HNO3:
2. Observation after the addition of (NH4)2CO3:
3. Observation after the addition of KSCN:
4. Observation after the addition of NaBiO3:
5. Observation after the addition of Na2SO4:
6. Observation after the addition of (NH4)2C2O4:
7. Observation after the addition of KI:
8. Observation after the addition of HDMG and Na2HPO4:
Results: Identification Code: _______ Cations present: _________________________
Bergen Community College 108 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 11: Qualitative Analysis of Cations
POST-LABORATORY QUESTIONS
1. What purpose does the addition of 6M NH3 in step 2 of the procedure serve?
2. An unknown solution was tested using the procedure in this experiment. Upon the addition
of 6M NH3 and 6M HNO3, a precipitate formed. The mixture was centrifuged and decanted.
The precipitate was dissolved in 6M HNO3 and separated into two test tubes. Solid KSCN was
added to the first; there was no change observed. Solid NaBiO3 was added to the second. After
30 seconds, the solution became purple, and then the color faded after ten minutes. (NH4)2CO3
was then added to the liquid decanted after the centrifugation. No change was observed. The
liquid was then transferred to a small beaker and heated to dryness. Once cool, 1M H3PO4 and
1M Na2HPO4 were added to re-dissolve the solid, and the solution was divided into two test
tubes. To the first, solid KI was added. A pale precipitate formed and the solution turned
brown. To the second, HDMG was added, and a red precipitate formed. Which cations are
present in the unknown?
3. Another unknown was tested using this procedure. Upon the addition of 6M NH3 and 6M
HNO3, no precipitate formed. (NH4)2CO3 was then added. A precipitate formed and the
mixture was centrifuged and decanted. The liquid was then transferred to a small beaker and
heated to dryness. Once cool, 1M H3PO4 and 1M Na2HPO4 were added to re-dissolve the solid,
and the solution was divided into two test tubes. To the first, solid KI was added; no change
was observed. To the second, HDMG was added; a red precipitate formed. The solid from the
centrifugation was then dissolved in 6M HCl and Na2SO4 was added. A white precipitate
formed. The mixture was centrifuged and decanted. 6M NH3 and (NH4)2C2O4 were added to
the liquid. A white precipitate formed. Which cations are present in the unknown?
Bergen Community College 109 General Chemistry II Laboratory
Experiment 12
Titration of Hydrogen Peroxide
OBJECTIVE: To analyze hydrogen peroxide solution by titration with standardized potassium permanganate.
BACKGROUND:
A. Permanganate (Oxidizing Agent)
Potassium permanganate is a powerful oxidizing agent commonly used as a reagent
in quantitative analysis. Manganese possesses several stable oxidation states. Those
encountered in the titration of hydrogen peroxide are Mn(VII), Mn(IV) and Mn(II), as seen
in the following half reactions:
MnO4 - (aq) + 4 H+ (aq) + 3e- MnO2 (s) + 2 H2O (l) (eq. 1)
MnO4 - (aq, purple) + 8 H+ (aq) + 5e- Mn2+ (aq, clear) + 4 H2O (l) (eq. 2)
The factors that govern these reduction products is complex. However, the following
general statement can be made. When permanganate is added to a neutral or an alkaline
solution as in equation 1, it is reduced to MnO2, or Mn(IV). When permanganate is added
to an acidic solution (greater than 0.5M), as in equation 2, then it is reduced to Mn(II).
Potassium permanganate solutions have limited stability. Their decomposition is
catalyzed by light, heat, acids, bases, Mn(II) and MnO2. Therefore, permanganate solutions
need to be standardized immediately prior to use to determine their concentration.
The color of permanganate solutions is an intense, deep purple, as shown in
equation 2 and the table below. When monitoring this reaction, the presence of
permanganate is readily identified because of the deep purple color of the permanganate
ion. Thus, the addition of a visual indicator is usually not required during titration; the
purple color of the permanganate is used to indicate its presence or absence.
FORMULA COLOR Analytical Wavelength
nm
Mn2+(aq) Colorless Does not absorb visible light
MnO4 -(aq) Purple 570nm
Also, it is difficult to read the bottom of the meniscus in a buret filled with
permanganate solution because of its intense color. In such cases, all readings are usually
made referring to the top of the meniscus. It is imperative to be consistent in this regard for
all readings.
Experiment 12 Titration of Hydrogen Peroxide
Bergen Community College 110 General Chemistry II Laboratory
Permanganate solutions can be standardized using Fe2+ and the redox reaction
between the permanganate and the iron(II) ions. The mechanism by which this reaction
proceeds is exceedingly complex, and reproducible analytical results can be obtained only
if certain empirical conditions are fulfilled. It is important to note that this reaction is
autocatalytic. In other words, the Mn2+ ions formed in this reaction act as a catalyst for the
reaction itself.
B. Mohr’s Reagent (Primary Standard Reducing Agent)
Mohr's salt, or ammonium iron(II) sulfate hexahydrate, is an inorganic compound
with the formula is (NH4)2Fe(SO4)2·6H2O. It contains two different cations, Fe2+ and NH4+,
and is classified as a double salt of iron(II) sulfate and ammonium sulfate. Since it is readily
available in pure form, and the crystals have a long shelf life and do not readily oxidize in
air, it can be used as a primary standard. Sulfuric acid is typically added to a solution
containing the ammonium iron(II) sulfate. The acid slows down the oxidation to Fe(III) in
water. An accurate and reliable determination of the amount of ammonium iron(II) sulfate
hexahydrate can be obtained via its mass.
When the ammonium iron(II) sulfate is oxidized, the Fe(II) is converted to Fe(III)
Fe+2(aq) e- + Fe+3(aq) (eq. 3)
The net reaction between the Mohr’s reagent and the Permanganate is seen below.
5Fe+2(aq) + MnO4 -(aq) + 8H+(aq) 5Fe+3(aq) + Mn2+(aq) + 4H2O(l) (eq. 4)
The ratio of permanganate to the ammonium iron(II) sulfate in the standardization
reaction is 1 mol permanganate to 5 mol ammonium iron(II) sulfate.
C. Hydrogen Peroxide, H2O2 (Unknown Reducing Agent)
Hydrogen peroxide is a thermodynamically unstable molecule. Over a period of
time the amount of peroxide in a solution will spontaneously decline. The rate of the decay
is affected by several factors, including storage temperature, the pH of the solution, and the
concentration of the solution. The stability of a peroxide solution can be increased if the
solution is dilute, stored at cool temperatures, and maintained at an acidic pH. Before
peroxide solution is used, it is best to determine its concentration with a standardized
oxidizing agent like permanganate, where the ratio of permanganate to the hydrogen
peroxide in the reaction is 2 mol MnO4 - to 5 mol H2O2.
5H2O2(aq) + 2MnO4 -(aq) + 6H+(aq) 5O2(g) + 2Mn
2+(aq) + 8H2O(l) (eq. 5)
In this experiment, a hydrogen peroxide solution with unknown concentration will be analyzed by titration with potassium permanganate. The potassium permanganate
solution will first be standardized with Mohr’s salt.
Bergen Community College 111 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
PRE-LABORATORY QUESTIONS
1. During the titration (standardization) of potassium permanganate, why it is
unnecessary to add an indicator?
2. Why is the solution of ammonium iron(II) sulfate acidified with sulfuric acid before
titration?
3. What is meant by the term autocatalytic in the context of this laboratory exercise?
4. Why is ammonium iron(II) sulfate hexahydrate a suitable primary standard for the
titration of the permanganate solution?
5. During the titration, the sides of the flask are washed down with deionized water from a
wash bottle. Why is this procedure necessary? And why it can be carried out without
affecting the result of the titration?
Bergen Community College 112 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
PRE-LABORATORY QUESTIONS (continued)
6. A student carried out the standardization of a potassium permanganate solution using
the procedure described in this experiment.
Mass of empty container 0.354 g
Mass of container with ammonium iron(II) sulfate hexahydrate 0.632 g
Initial potassium permanganate buret reading 0.38 mL
Final potassium permanganate buret reading 33.53 mL
a. Calculate the number of moles of (NH4)2Fe(SO4)2∙6H2O. ________
b. Calculate the number of moles of KMnO4. ________
c. Calculate the molarity of the KMnO4 solution. ________
7. A solution of potassium permanganate was standardized and its concentration was found
to be 0.0195M. A 10.00 mL sample of hydrogen peroxide was then titrated with the
potassium permanganate solution. Assume the density of the peroxide is 1.00 g/mL.
Initial buret reading 33.53 mL
Final buret reading 49.13 mL
Determine the % mass of the hydrogen peroxide in the solution.
Experiment 12 Titration of Hydrogen Peroxide
Bergen Community College 113 General Chemistry II Laboratory
REAGENTS: EQUIPMENT:
4 M H2SO4 250-mL Erlenmeyer flask
(NH4)2Fe(SO4)2∙6H2O 10-mL graduated cylinder
0.2M KMnO4 solution 250-mL beaker
Unknown H2O2 solution buret, stir station and magnetic stirrer
150-mL Erlenmeyer flask
25-mL pipette
SAFETY: SAFETY ALERT
Hydrogen peroxide- corrosive
Potassium permanganate- irritant
Sulfuric acid- corrosive
Material Safety Data Sheets of these chemicals can be found at www.msds.com
Caution: This version of the KMnO4 standardization is modified. The more common
method uses sodium oxalate, which is poisonous. In place, ammonium ferrous sulfate
hexahydrate, (NH4)2FeSO4∙6H2O, known as Mohr's salt is used. WEAR DEPARTMENTALLY APPROVED EYE PROTECTION and APRON WHILE DOING THIS
EXPERIMENT.
PROCEDURE:
Part A: Standardization of Permanganate
1. Tare a 250-mL Erlenmeyer flask and add 0.25 - 0.30 grams (NH4)2Fe(SO4)2∙6H2O,
Mohr’s Salt. Weigh and record the mass. (Alternatively, a clean, dry piece of weighing
paper or a weighing boat may be used and the salt transferred to the flask.) DO NOT
PLACE THE MOHR’S SALT DIRECTLY ONTO THE BALANCE.
2. Add 50 – 60 mL distilled water and 10 mL 4.0 M sulfuric acid solution. Swirl the flask
until the solid dissolves.
3. Obtain 100 mL (maximum) KMnO4 solution in a 250-mL beaker. Place a watch glass
(a cover) over the top of the beaker.
4. Prepare a buret with the permanganate solution and record the initial buret reading,
remembering to read the buret referring to the top of the meniscus. Titrate the Mohr’s Salt
slowly to a faint pink color (indicating excess permanganate) that persists for 15 seconds.
See Appendix D for general instructions on titration.
5. Perform a second determination.
Discard all titrated solutions in the appropriate container.
Experiment 12 Titration of Hydrogen Peroxide
Bergen Community College 114 General Chemistry II Laboratory
Part B: Analysis of Hydrogen Peroxide
1. Pour approximately 50 mL of the assigned hydrogen peroxide solution into a clean, dry
150-mL beaker and cover with a watch glass. Record the unknown identification code.
2. Condition a 25-mL pipet twice with the peroxide solution, and pipet 25.00 mL of the
solution into clean, but not necessarily dry, 250-mL Erlenmeyer flask. See Appendix B for
general instructions for volumetric glassware.
3. Add 10 mL of 4.0 M H2SO4 to the flask and mix thoroughly.
4. Record the initial buret reading and titrate the peroxide solution with the permanganate
solution until a pink coloration persists for 15 seconds. Record the final buret reading. See
Appendix D for general instructions on titration.
5. Perform a second determination. If the time allows, perform a third determination.
Successive determinations should agree to within one tenth of a milliliter.
Discard all titrated solutions in the appropriate container. Wash hands thoroughly with
soap before leaving the laboratory.
CALCULATIONS
Part A: For each determination:
1. Calculate the number of moles (NH4)2Fe(SO4)2∙6H2O, Mohr’s Salt, and hence Fe 2+ used.
2. Calculate the number of moles of MnO4ˉ from the stoichiometry in equation 4.
3. Calculate the molarity of the permanganate solution, and obtain the average.
Part B: For each determination:
1. Using the molarity of the permanganate solution determined in Part A and its volume,
calculate the number of moles of MnO4ˉ used.
2. Calculate the number of moles of H2O2 from the stoichiometry in equation 5.
3. Calculate the mass of the H2O2 using its molar mass.
4. Calculate the mass of the unknown hydrogen peroxide sample from its volume. Dilute
hydrogen peroxide solutions such as these have a density of 1.00 g/mL.
5. Calculate the percent by mass of H2O2 in the unknown sample, and obtain the average.
% H2O2 = (mass of H2O2 / mass of the unknown sample) x 100
Bergen Community College 115 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
Part A: Standardization of Permanganate
DATA: Determination 1 Determination 2
Mass of weighing paper, g (if using) ____________ ____________
Mass of weighing paper + (NH4)2Fe(SO4)2∙6H2O, g ____________ ____________
Mass of (NH4)2Fe(SO4)2∙6H2O, g ____________ ____________
Initial buret reading, mL ____________ ____________
Final buret reading, mL ____________ ____________
Volume of permanganate, mL ____________ ____________
Results: Show calculations below or on a separate sheet.
Number of moles of (NH4)2Fe(SO4)2∙6H2O, mol ____________ ____________
Number of moles of MnO4ˉ, mol ____________ ____________
Molarity of permanganate solution, mol/L ____________ ____________
Average molarity, mol/L ___________________
Bergen Community College 116 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
Part B: Analysis of Hydrogen Peroxide
Data: Unknown Code: _________
Determination Determination Determination
1 2 3, if needed
Volume of peroxide solution, mL ___________ ___________ ___________
Initial buret reading, mL ___________ ___________ ___________
Final buret reading, mL ___________ ___________ ___________
Volume of permanganate, mL ___________ ___________ ___________
Results: Show calculations below or on a separate sheet.
Number of mol MnO4ˉ, mol ___________ ___________ ___________
Number of mol of H2O2, mol ___________ ___________ ___________
Mass of H2O2, g ___________ ___________ ___________
Mass hydrogen peroxide sample, g ___________ ___________ ___________
Percent H2O2 in the sample, % ___________ ___________ ___________
Average percent, % ______________
Bergen Community College 117 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
POST-LABORATORY QUESTIONS
1. Assume the solutions used in this experiment have actual concentrations of 0.15% H2O2
and 0.020 M KMnO4. Determine the percent error in the experimental results for each.
% Error = (|accepted value – experimental value| / accepted value) × 100
2. In order to analyze a commercial H2O2 solution using the procedure in this experiment, it
must first be diluted with distilled water. Consider 10.0 mL of a commercial peroxide solution
measured with volumetric pipet diluted to 100.0 mL in a volumetric flask. If 25.00 mL of the
diluted solution is titrated with 0.0202 M MnO4, calculate the percent H2O2 in the commercial
peroxide solution, i.e. before dilution. The initial buret reading was 0.08 mL, and the final
reading was 27.59 mL. Assume that the diluted peroxide solution had a density of 1.00 g/mL.
Bergen Community College 118 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 12: Titration of Hydrogen Peroxide
POST-LABORATORY QUESTIONS (continued)
3. Sodium oxalate is generally used for the standardization of KMnO4 solution where it reacts
in strongly acidic solution according to the following chemical equation:
5 C2O4 2-
(aq) + 2 MnO4 - (aq) + 16 H
+ (aq) 10 CO2(g) + 2 Mn
2+ (aq) + 8 H2O(l)
A student carried out the standardization of a potassium permanganate solution using a
procedure similarly described in this experiment, except using sodium oxalate as a standard.
The following data were obtained.
Mass of the weighing paper 0.354 g
Mass of sodium oxalate & weighing paper 0.632g
Initial buret reading 0.38 mL
Final buret reading 33.53 mL
a. Calculate the number of moles of Na2C2O4.
b. Calculate the number of moles of KMnO4.
c. Calculate the molarity of KMnO4.
Bergen Community College 119 General Chemistry II Laboratory
Experiment 13
Electrochemistry
OBJECTIVE: To build a galvanic cell, measure cell potential and perform electrolysis.
BACKGROUND:
Although energy cannot be created nor destroyed, it can be transformed from one form
to another. A device that converts chemical energy into electric energy is called a voltaic cell
(named for Alessandro Volta) or a galvanic cell (named for Luigi Galvani). The chemical
reactions that occur in this type of electrochemical cell are oxidation-reduction reactions, and
are spontaneous, i.e. the change in Gibbs Free Energy, ΔG, is negative. Electrons are
transferred from one species to another due to a difference in potential energy, moving from a
position of higher to lower potential energy.
In a voltaic cell, the reactants for the oxidation half-reaction and the reduction half-
reaction are placed in two separate compartments, or half-cells. Electrodes, which provide a
surface for electron transfer, are placed in the half-cells and connected with a wire. The species
being oxidized gives up electrons at the negative electrode, the anode, and travel across the
external circuit to the positive electrode, the cathode, in the other half-cell. Here, the electrons
are gained by the species being reduced. A salt bridge connects the two half-cells so that ions
can migrate from one to the other in order to maintain electrical neutrality. Useful work could
potentially be done by the current generated in the external circuit. A schematic of a voltaic
cell is shown below.
The difference in the potential energy of the electrons at the cathode and at the anode,
E, is measured in volts, V, and is primarily dependent on the species involved. Voltage also
depends on the concentrations of the reacting species and gas pressures. Standard half-cell
potentials, E°, are measured relative to an assigned reduction potential of zero for the standard
hydrogen electrode (SHE), where PH2 = 1 atm and [H +] = 1.0 M. Some half-cell potential are
shown on the following page.
Experiment 13 Electrochemistry
Bergen Community College 120 General Chemistry II Laboratory
Some standard half-cell
potentials, Eo:
The metal with the lower standard reduction potential serves as the anode. Thus, in the
galvanic cell represented on the previous page, the zinc electrode is the anode as it has a lower
reduction potential than copper, as seen above. The zinc anode is oxidized to aqueous zinc ion,
causing the zinc to dissolve over time. Copper(II) ions are reduced to copper metal at the
cathode, increasing the mass of the copper over time.
ANODE
Zn dissolves and [Zn+2] increases.
CATHODE
[Cu+2] decreases and Cu accumulates.
The chemical equations for the reactions are as follows:
Half-cell Reactions:
At the anode: Zn(s) Zn2+(aq) + 2 e- (oxidation: losing electrons) At the cathode: Cu2+(aq) + 2 e- Cu(s) (reduction: gaining electrons)
The overall reaction: Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)
The cell potential is given by EoCell = E o cathode - E
o anode
Eocell = E o cathode – E
o anode = 0.34 V – (- 0.76 V) = 1.10 V
A diagram of a functioning galvanic cell using this redox reaction is found on the next page.
EoH+/H2= 0.0 V reference potential
at 25 oC and 1 atm
EoAg+/Ag = +0.80 V
EoCu2+/Cu = +0.34 V
EoZn2+/Zn = -0.76 V
Lighter
blue
color
Experiment 13 Electrochemistry
Bergen Community College 121 General Chemistry II Laboratory
Electrons move from
the zinc anode to the
copper cathode.
In the salt bridge, the
cations ions move to
the cathode while the
anions move to the
anode.
The salt bridge compensates for the ionic gradients. Since [Zn2+] increases in the
oxidation half-cell and [Cu2+] decreases on the reduction half-cell, the K+ ions will move to
the cathode and the NO3ˉ ions to the anode to keep the electroneutrality in check.
In another type of electrochemical cell called an electrolytic cell, the reverse process
of a galvanic cell occurs. Energy is supplied from an outside source to reverse the reactions.
In electrolysis, electrical energy is used to decompose a compound into its constituents. With
all electrochemical cells, oxidation occurs at the anode.
The figure to the left shows the electrolysis of
acidified water, using inert electrodes. The water
decomposes into hydrogen and oxygen gases. The
volume of the H2 is twice that of the O2 under the same
conditions. The following reactions occur:
CATHODE: 2 H2O(l) + 2 eˉ → H2(g) + 2 OHˉ(aq)
ANODE: 2 H2O(l) → O2(g) + 4 H+(aq) + 4 eˉ
OVERALL: 2 H2O(l) 2 H2(g) + O2(g)
During the electrolysis of an aqueous solution of an alkali
metal salt such as NaCl(aq), depicted to the right, chloride ions are
oxidized to chlorine gas at the anode. However, sodium is not reduced
at the cathode. Since the reduction potential of hydrogen (0.0 V) is
higher than that of sodium (-2.71 V), hydrogen is reduced rather than
sodium ion, and H2 gas is evolved at the cathode. Alkali metals can
only be produced when the molten halide salts of sodium or potassium
are electrolyzed using inert electrodes, such as graphite or platinum.
CATHODE: 2 H2O(l) + 2 e - H2(g) + 2 OHˉ(aq)
ANODE: 2 Clˉ(aq) Cl2(g) + 2 e -
OVERALL: 2 NaCl(aq) +2 H2O(l) → 2 NaOH(aq) + H2(g) + Cl2(g)
Cu deposit
Anode Cathode
Experiment 13 Electrochemistry
Bergen Community College 122 General Chemistry II Laboratory
The amount of metal deposited at the cathode during electrolysis is proportional to the
quantity of electricity used, measured in coulombs, C. A coulomb is the quantity of charge
passing a point in a circuit in one second when the current is one ampere. Hence, an ampere
is a coulomb per second, 1 A = 1 C/s. The amount of metal deposited also depends on the
number of electrons required per ion (i.e. the charge on the metal ion) and its molar mass. The
charge of a mole of electrons is referred to as Faraday’s constant, and has a value of 96,485
C/mol eˉ. The mass of metal deposited can thus be determined using dimensional analysis.
Example 1. What is the cell potential of a voltaic (galvanic) cell made of a zinc electrode
immersed in 1.0 M Zn(NO3)2 and a gold electrode immersed in 1.0 M Au(NO3)3?
EoAu 3+
/Au = 1.42 V and E o Zn
2+ /Zn = - 0.76 V
Answer. The standard reduction potential of the gold is higher than that of the zinc. Therefore,
the Au is the cathode and the metallic Zn is the anode.
Eocell = E o cathode – E
o Anode = 1.42 V - (- 0.76 V) = 2.18 V
Example 2. How many grams of copper will be deposited from a copper(II) sulfate solution,
CuSO4, in 16 minutes and 5 seconds using a current of 10. A? Cu = 63.5 g/mol
Answer. 16 minutes five seconds is equal to 965 seconds. Cu2+ requires 2 eˉ to be reduced.
965 s 10. C mol eˉ mol Cu 63.5 g Cu = 3.2 g Cu
s 96,485 C 2 mol eˉ mol Cu
In this laboratory exercise, a voltaic cell will be constructed from copper metal in a
solution of copper(II) sulfate, and zinc metal in a solution of zinc sulfate. Additionally, an
electrolytic cell will be constructed using an aqueous solution of sodium chloride.
REAGENTS: EQUIPMENT:
1.0 M CuSO4 sand paper copper and zinc electrodes
1.0 M ZnSO4 2 250-mL beakers connecting wires and alligator clips
Saturated NaCl solution thermometer voltmeter
Phenolphthalein salt bridge graphite electrodes
9-V battery
SAFETY:
Chlorine gas is toxic and corrosive. Do NOT breathe the gas. Perform this experiment in the
fume hood or well-ventilated area. Copper(II) sulfate is toxic and an irritant.
Bergen Community College 123 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 13: Electrochemistry
PRE-LABORATORY QUESTIONS
1. What is a galvanic cell?
2. In a galvanic cell, at which electrode does reduction occur?
3. Identify the parts of a galvanic cell.
4. Consider the following experiment conducted with unknown metals X, Y, and Z, along
with small beakers containing solutions of X2+(aq) and Z2+(aq). When metal X was dropped
into the beaker containing Z2+(aq), no reaction occurred. When metal Y was dropped into the
beaker containing the Z2+(aq), a reaction did occur. Based on these observations, rank the
metals from least active to most active.
Experiment 13 Electrochemistry
Bergen Community College 124 General Chemistry II Laboratory
PROCEDURE:
Part A: Galvanic Cell
1. Obtain copper and zinc electrodes; clean with sand paper until shiny; wash with water, and
dry.
2. Obtain two 250-mL beakers. Add 100 mL of 1.0 M CuSO4 to one, and 100 mL 0.10 M
ZnSO4 to the other.
3. Fill a salt bridge with a saturated
solution of NaCl and stopper it.
4. Obtain connecting wires, alligator clips,
and a voltmeter. Assemble the galvanic
cell as shown to the right.
5. Measure and record the temperature of
the solutions.
6. Record the voltage displayed on the voltmeter.
7. Compare the experimental value to the accepted value of 1.10 V; discuss how these differ.
8. Continue to observe the voltage, and record after 10 minutes.
9. Compare the new voltage to the initial value; discuss how or why they differ.
10. Disassemble the galvanic cell. Discard the solutions in the appropriate waste container.
Clean and return the glassware and the electrodes.
Part B: Electrolytic Cell
Caution: Small amounts of chlorine gas will evolve at the anode. Do NOT breathe the gas.
This experiment may be done in the fume hood or well-ventilated area.
1. Add 150 mL of NaCl solution to a 250-mL beaker.
2. Immerse two graphite electrodes into the NaCl solution, and connect to a 9-V battery
using alligator clips.
3. Observe and record the gases released at the cathode and at the anode.
4. After two minutes, add two drops of phenolphthalein into the beaker. Observe and record
the color change.
5. Write chemical equations for the reactions at the cathode, the anode, and overall.
6. Dispose of the electrolytic solution in the sink with running water. Clean and return the
electrolytic cell.
7. Wash hands before leaving the laboratory.
Bergen Community College 125 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 13: Electrochemistry
DATA:
Part A: Galvanic Cell
1. Temperature _______________
2. Initial voltage _______________
3. Comparison to 1.10 V:
4. Voltage after 10 minutes _______________
5. Comparison to initial voltage:
Part B: Electrolytic Cell
1. Observation of Gases:
2. Color with phenolphthalein _______________
3. Chemical equations:
Anode: _____________________________________________
Cathode: _____________________________________________
Overall: _____________________________________________
Bergen Community College 126 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 13: Electrochemistry
POST-LABORATORY QUESTIONS
1. A cell is constructed using a silver electrode and a copper electrode in their appropriate
solutions. EoAg +
/Ag = +0.80 V and EoCu2+/Cu = +0.36 V. What is E°cell?
2. Which metal serves as the anode in the above cell?
3. What is the oxidation number of Cr in chromate, CrO4 2-?
4. In the following reaction, which element is reduced and which is oxidized?
Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)
Bergen Community College 127 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 13: Electrochemistry
POST-LABORATORY QUESTIONS (continued)
5. What are the products of the electrolysis of an aqueous solution of NaCl?
6. What are the products of the electrolysis of a molten NaCl?
7. A current of 5 A is passed through a solution of CuSO4 for 30 minutes. How many grams
of Copper are deposited at the cathode? Cu = 63.54 g/mol.
8. In Part B of this experiment, why did the color of the phenolphthalein change during the
electrolysis of the NaCl solution?
Bergen Community College 128 General Chemistry II Laboratory
Experiment 14
Determination of Mass Composition of a
Mixture of Na2CO3 and NaHCO3
OBJECTIVE: To determine the mass composition of a sample of sodium carbonate and sodium bicarbonate mixture using small scale laboratory methods.
BACKGROUND:
Titration is a procedure used in analytical chemistry to determine the amount or
concentration of a substance. In a titration one reagent, the titrant, is added to another slowly.
As it is added a chemical stoichiometric reaction occurs until one of the reagents is exhausted,
and some process or device signals that this has occurred. The purpose of a titration is
generally to determine the quantity or concentration of one of the reagents, that of the other
being known beforehand (e.g. standardized). In any titration there must be a rapid quantitative
reaction taking place as the titrant is added, and in acid-base titrations this is a stoichiometric
neutralization. The type of titration is simply the type of chemical reaction taking place. This
section utilizes an acid-base titration.
Na2CO3 is a stronger base than NaHCO3, with pKa1 = 10.32 and pKa2 = 6.37. The pKa
of NaHCO3 is 6.37. The titration uses two different indicators. The first is phenolphthalein,
which has a pKa of 8.2. In this part of the titration, Na2CO3 is titrated to NaHCO3. The second
indicator is bromocresol green, which has a pKa of 3.8. Therefore in this part, NaHCO3 (both
that present in the original sample and that formed from the first titration) will be titrated.
In this experiment, an unknown mixture of
Na2CO3-NaHCO3 will be titrated by a standardized
HCl solution. Since the concentration can change
during storage, its concentration is determined
before titrating the mixture. The standardization
should be performed against a standardized NaOH
solution. NaOH can react with atmospheric CO2 and
turn into Na2CO3. Therefore, the NaOH solution
must first be titrated with KHP. Then, the accurate
concentration of the HCl solution can be determined
after being titrated with the NaOH solution. It is a
tedious work, but it has to be done correctly.
Figure 1. pH profile of the carbonate
and bicarbonate mixtures.
Experiment 14 Determination of Mass Composition of a Mixture of Na2CO3 and NaHCO3
Bergen Community College 129 General Chemistry II Laboratory
REAGENTS:
Unknown sodium carbonate and sodium bicarbonate mixture 1.0 M HCl solution
Phenolphthalein and bromocresol green indicators
potassium hydrogen phthalate (KHP) (dried)
CO2 free water: Boil 100-mL distilled water in a 250-mL beaker, for 1 minute. Cool.
triple distilled water
EQUIPMENT:
Thin-stem pipettes (2) 100-mL beaker 50-mL beaker
(2) 10 mL graduated cylinder 125-mL Erlenmeyer flask ring stand and clamp
magnetic stirrer and stirrer hot plate analytic balance buret
SAFTEY:
Hydrochloric acid is toxic and corrosive. Wear protective glasses at all times throughout the
experiment. In case of eye contact, immediately wash the eyes with large amounts of water,
occasionally lifting the lower and upper lids. Get medical attention immediately. In case of
skin contact, immediately flush the contaminated skin with water for at least 15 minutes. Do
not rub the skin. Notify the instructor. Rinse spills with large amounts of water. Do not dispose
of any solutions before asking the instructor. Wash hands before leaving the laboratory.
PROCEDURE:
I. Standardization of NaOH solution
1. Select clean and dry clean 125-mL Erlenmeyer.
2. Add about 0.150 g of KHP to this beaker.
3. Add about 25 mL of distilled water to KHP.
4. Add a drop of phenolphthalein to KHP solution and place the magnetic stirrer
into the beaker.
5. Select a clean 50-mL buret and fill it with 50.00 mL of NaOH solution.
6. Using the buret, add small portions of NaOH solution into the KHP solution,
drop by drop. Add NaOH solution continuously to the KHP solution until the
indicator retains a light pink color for about 15 seconds.
7. Record the volume of NaOH solution used.
8. Calculate the concentration of the NaOH solution provided. (See calculations
A below.)
Experiment 14 Determination of Mass Composition of a Mixture of Na2CO3 and NaHCO3
Bergen Community College 130 General Chemistry II Laboratory
II. Standardization of HCl solution
1. Using a 10-mL graduated cylinder add 10.00 mL
of HCl solution to a 125-mL Erlenmeyer.
2. Add one drop of phenolphthalein indicator
(observe the colorless solution).
3. Record the volume of NaOH in the buret. Start
adding drop wise the standardized NaOH
solution to the Erlenmeyer until a pale pink
color.
4. Record the volume of NaOH solution added for
neutralization of the HCl solution. And,
determine the concentration of HCl solution.
1. (See calculations B below.) Figure 2. Titration setup.
III. Determination of the composition of the unknown mixture:
Titration of the unknown mixture with the standardized HCl solution.
1. Weigh about 2.5 grams of the sample provided. Record the mass: msample
2. Place in a 125-mL Erlenmeyer flask and dissolve completely by adding CO2
free water.
3. Add a single drop of phenolphthalein indicator to the Erlenmeyer flask.
4. Fill the buret with 1.0 M HCl solution.
5. Start titrating by adding HCl solution until the pink color of the solution
disappears. Record the volume of HCl added: Vphth
6. Add 2-3 drop of Bromocresol Green indicator. Observe the blue color of the
solution.
7. Start adding again the HCl solution. After adding 2.00 mL of HCl solution,
remove the Erlenmeyer flask and place on a hot plate and start heating the
solution. Boil for 1 minute and remove from the heater. Observe the dark blue
color of the solution.
Experiment 14 Determination of Mass Composition of a Mixture of Na2CO3 and NaHCO3
Bergen Community College 131 General Chemistry II Laboratory
8. Wait until the solution is cooled down.
9. Continue the titration until the indicator changes its color to yellow.
10. Record the total volume: Vtotal in milliliters.
11. Repeat the experiment with two more samples . Record the volumes of HCl
used for each titration.
12. Take the average of the HCl volumes of the titrations.
13. Determine the volume used for Na2CO3 titration which is equal to (2×Vphth)
and the volume needed for NaHCO3 as VBCG = Vtotal – (2×Vphth) in milliliters.
14. After the experiment finished, clean the workbench and all samples were
returned.
Calculations
A. KHP titrated with NaOH
KHP (Potassium hydrogen phthalate) is a weak monoprotic acid. The reaction involved in
this experiment is represented by the following equation:
KHP(aq) + NaOH(aq) KNaP(aq) + H2O(l)
To neutralize 1 mol of KHP (monoprotic), 1 mol of NaOH (monovalent) is required.
nKHP = nNaOH (eq. 1)
mKHP
MWKHP = MNaOH×VNaOH (eq. 2)
mKHP is the mass of KHP used. The molecular weight of KHP is 204 g/mol. VNaOH is the
volume of NaOH, in liters, used to titrate the KHP solution.
Equation 2 is rearranged so that it yields the molarity of NaOH, MNaOH:
MNaOH= mKHP
MWKHP×VNaOH (eq. 3)
Molarity of NaOH solution: ________________
B. Concentration of HCl provided.
The hydrochloric acid reacts to 1:1 ratio with NaOH solution according to the
following reaction.
2. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
nHCl = nNaOH
Experiment 14 Determination of Mass Composition of a Mixture of Na2CO3 and NaHCO3
Bergen Community College 132 General Chemistry II Laboratory
𝑴𝑯𝑪𝒍 = 𝑴𝑵𝒂𝑶𝑯
𝑽𝑯𝑪𝒍 × 𝑽𝑵𝒂𝑶𝑯 (eq. 4)
Molarity of HCl Solution: _________________
C. Titration of the unknown carbonate/bicarbonate mixture with an HCl solution.
Na2CO3 titration requires two (2) moles of HCl solution. However, the titration of
NaHCO3 requires only one (1) mole of HCl solution.
Na2CO3 + 2HCl 2NaCl + CO2 + H2O
NaHCO3 + HCl NaCl + CO2 + H2O
x = the mass of the Na2CO3 (MW = 106 g/mol) is the sample.
y = the mass of the NaHCO3 (MW = 84 g/mol) in the sample.
1. Vphth: Average volume of HCl needed to titrate Na2CO3: ______________ L
2. VBCG: Average volume of HCl needed to titrate NaHCO3: ______________ L
(VBCG = Vtotal of HCl - 2 × Vptht)
3. Number of mol of HCl to titrate Na2CO3 = Vptht × Molarity of HCl = __________
mol
4. Number of mol of HCl to titrate Na2CO3 = VBCG × Molarity of HCl = __________
mol
5. x, mass of Na2CO3 in the sample: (line 3) × 2 × 106 g/mol = _______________ g
6. y, mass of NaHCO3 in the sample: (line 4) × 1 × 84 g/mol = _______________ g
D. Determining the mass composition of the mixture. % Na2CO3 =
(𝑥)
𝑚𝑠𝑎𝑚𝑝𝑙𝑒 × 100 = %
% NaHCO3 = (𝑦)
𝑚𝑠𝑎𝑚𝑝𝑙𝑒 × 100 = %
Note: The sum of the percent compositions of these two compounds should be equal to 100. A
value which is smaller than 100% indicates the existence of possible crystal water present in
the sample constituents.
Bergen Community College 133 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 14: Determination of Mass Composition of a Mixture of
Na2CO3 and NaHCO3
POST-LABORATORY QUESTIONS
1. Graph the titration of the mixture and attach to the laboratory report.
2. Determine the volume of 0.250 M of NaOH required to neutralize 0.125 grams of dried
KHP.
3. A student standardizes a solution NaOH using a KHP sample which is not dried
completely. Will the calculated concentration of NaOH be too high, too low or unaffected?
4. Why are different indicators needed to titrate a mixture of NaHCO3 and Na2CO3?
Bergen Community College 134 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 14: Determination of Mass Composition of a Mixture of
Na2CO3 and NaHCO3
POST-LABORATORY QUESTIONS (continued)
5. A sample of a washing
powder that contains a
mixture of Na2CO3 and
NaHCO3 is titrated with
aqueous HCl and the
result to the right is
obtained: What is the
mole ratio of CO3 2- to
HCO3 -in the washing
powder?
6. A titration graph (pH versus volume of
NaOH used is given to the right. Is this a
titration graph of a diacid or a two different
monoprotic acids with different Ka's and
different concentrations? Justify your
answer.
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Appendix A: Common Laboratory Equipment
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Appendix B: Volumetric Glassware
Although graduated cylinders are commonly used
in the chemistry laboratory, volumetric glassware often
provides the better option. Graduated cylinders are very
convenient, as they allow any volume up to the size of the
cylinder to be measured. Volumetric cylinders and
pipettes only allow for one volume to be measured, but it
is a much more precise measurement. Unlike a graduated
cylinder, whose thickness is constant, volumetric
glassware hold the bulk of a liquid in a wide portion and
the measurement mark is on a much thinner portion. The
precision in measuring liquids in either case is limited by
the visual task of ensuring that the meniscus lies on the
mark. Because the mark on volumetric glassware is on a
tube with a much smaller diameter than it is on a
graduated cylinder, the same error in height corresponds
to a much smaller volume. This provides a much more
precise measurement.
Using Volumetric Flasks
If the solute is a solid, add it to a small beaker and dissolve in roughly half the amount
of solvent needed. Transfer to the volumetric flask. Rinse the beaker with a small amount of
solvent and add to the flask as well. If the solute is a liquid, or the solution is being prepared
from a more concentrated one, pipette the correct amount directly to the volumetric flask.
Carefully add solvent to the mark using a dropper. Cap and invert several times to mix.
Using Volumetric Pipettes
Attach a pipette filler to the end of the pipette and place the tip in the solution to be
measured. Draw the liquid into the pipette, filling to roughly an inch above the mark, being
certain that the liquid does not enter the pipette filler. Carefully slide the filler off the end of
the pipette and replace it with an index finger to seal. Raise the tip of the pipette above the
liquid surface. Holding it vertically, allow the seal to loosen, but not to break. The liquid will
slowly drain from the pipette. When the meniscus just sits on the marked line, tighten the seal
to prevent further draining of the liquid. If there is a bead of liquid clinging to the tip of the
pipette, touch it to the wall of the beaker to remove it. Move the pipette to the desired container
and allow the seal to be broken so that the liquid is delivered. Once drained, if there is a bead
of liquid clinging to the tip, touch it to the beaker wall so that it is transferred. Volumetric
pipettes are labeled either TD or TC. If it is labeled TD, as is more common, the pipette is
calibrated to deliver the correct volume, and the last bit of liquid in the pipette should be left
behind. However, if it is labeled TC, it is calibrated to contain the correct volume, and all of
the liquid in the pipette should be transferred.
Bergen Community College 137 General Chemistry II Laboratory
Appendix C: Graphing
Graphing by Hand
Drawing the Axes
Using a ruler and graph paper, draw in the x and y axes, filling the page as much as possible
while leaving space between the axes and the edge of the paper for labels.
Choosing a Scale
For each axis, choose a scale that is easy to read in that 1 box should represent a nice whole
number or a simple fraction of units. Avoid threes and thirds! Some examples of good scales
are:
1 box = 1 unit, 10 units, 0.1 unit, or any power of 10
1 box = 2 units, 0.2 units, etc.
1 box = 5 units, 0.5 unit, etc.
Not every grid line should be marked on the graph. Mark enough to make the graph easy to
read but not so many that the graph becomes cluttered, and thus difficult to read.
Plotting Data Points
Find the x and y coordinates that correspond to the data points and draw them as small dots.
If the origin, (0,0), is a data point, be sure to include it.
Drawing the Curve
If the graph is linear, use a ruler to add the best-fit line. To do this, draw the line with points
lying above the line and below the line to balance out. If the curve is not linear, draw a smooth
line between the data points by hand. (A simple plastic French Curve can be used as an aid.)
Labels and Titles
Add labels, including units, to the axes. Add a title indicating what the graph represents.
Slope Determination
To determine the slope of a linear graph, select two points on the line that are NOT data points
(to increase accuracy) and are NOT close to one another (to increase precision). Indicate the
selected points on the graph, using a different symbol than used for data points. Calculate the
slope as the ratio of the rise to the run, ∆y/∆x. Remember to include units.
Appendix C Graphing
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Graphing in Excel
Preparing the Graph
1. Enter the data in columns, with the x values in the first column and the y values in
the second.
2. With data highlighted, go to the “insert” tab and select “Scatter” in the Charts
section of the ribbon.
3. Choose the Scatter Chart with NO LINES. A graph should appear.
4. Click on the “plus” sign at the upper right hand corner of the graph to access Chart
Elements.
Adding the Trendline
1. In “Chart Elements,” click the arrow that appears by “Trendline,” and select
“More Options...”
2. Choose the regression type. Most graphs prepared in General Chemistry are linear.
3. If linear, check the “Display Equation on Chart” and “Display R-squared value
on Chart” boxes.
4. If the curve must include the (0,0) point, check the “Set Intercept” box.
Adding a Title
1. In “Chart Elements,” select “Chart Title.”
2. “Above the Chart” is standard; other options can be found under “More Options...”
Labeling the Axes
1. In “Chart Elements,” select “Axis Titles.”
2. The standard for the Horizontal Axis is below the axis, and for the vertical axis is a
rotated title.
3. Include units in the titles of the axes.
Finishing Touches
1. To change the range of values displayed on an axis, in “Chart Elements,” click the
arrow that appears by “Axes,” check the desired axis, and select “More Options...”
At the top, select “Axis Options” and select “Axis Options” again in the list. Here
the minimum and/or maximum values can be “fixed.”
2. Be certain that the data points can be readily seen and axis scales are easy to read.
Printing the Graph
1. The Chart Area must be highlighted (selected) in order to print the Chart only.
Bergen Community College 139 General Chemistry II Laboratory
Appendix D: Titration
Titration is a method used in quantitative analytical analysis that allows for the
determination of the amount of a substance present in a mixture. The substance is chemically
reacted with a known amount of another substance. By the stoichiometry of the reaction, the
amount of the unknown can be obtained. Many types of reactions are utilized in titrations,
including acid-base reactions and oxidation-reduction reactions.
The titrant, a solution containing one of the reactants, is placed in a buret. This common
laboratory glassware is a cylindrical tube that precisely measures volume, as it can be read to
the nearest 0.01 mL. It has a stopcock at the bottom to allow for controlled introduction of the
solution to the titration flask. It is important to note that a buret does not directly measure the
volume it contains. The volume used is obtained by subtracting the initial buret reading from
the final one. These readings must lie within the 0.00 mL – 50.00 mL markings on the buret.
Also, a buret is marked with the 0 mL mark at the top, and the 50 mL mark towards the bottom.
The other reactant, which can be a solid or a solution, is placed directly into the flask.
The titrant is slowly added to the titration flask while its contents are stirred until the reaction
is complete. There are several methods used to signal the completion of the reaction. If the
titrant is colored whereas the reaction products are colorless, the presence of excess titrant is
easily observed. Acid-base titrations commonly use a visual indicator, a substance whose color
depends on the pH of the solution. Acid-base titrations can also be monitored with a pH meter.
General instructions for a titration are given below.
1. Set up a ring stand with a buret clamp, and a magnetic stir station. (The titration flask may
also be swirled by hand.)
2. Line, or condition, a buret by adding 5-10 mL of the titrant and carefully tilting and twisting
the buret so that the titrant has made contact with all of the inside surfaces. Drain the titrant
through the tip into a waste beaker. Repeat.
3. With the top of the buret below eye level, fill the buret with the titrant, and attach to the
buret clamp. Allow the titrant to run through the tip into the waste beaker until the tip is filled.
Be certain that there are no air pockets in the tip. This would give an inaccurate volume
measurement, as the size of the air pocket will change as the weight of the solution changes.
Add additional titrant, if needed, to between the 0 and 5 mL marks. Record the initial reading.
4. Add the substance to be titrated to a clean 125-mL Erlenmeyer flask. Add 25-35 mL
distilled water and, if using, the indicator. (Some titrations also require other components.)
5. Turn on the stir station, and slowly add titrant to the reaction flask. When the reaction nears
completion, rinse down the walls of the titration flask with distilled water from a squeeze
bottle, and add the titrant one drop at a time, continuing to rinse down the walls of the flask.
When it becomes necessary, the titrant can be added half a drop at a time. Allow a bead of
titrant to collect on the tip of the buret and close the stop cock. Touch the bead to the inner
wall of the titration flask, and rinse down with distilled water.
6. When the endpoint has been reached, record the final buret reading. For the subsequent
determination, be sure to add sufficient titrant to the buret if necessary.
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Appendix E: Filtration
Filtration is the process of separating a liquid from a solid by allowing the liquid to
pass through a porous barrier, such as paper, while leaving the solid behind. The liquid that
passes through is called the filtrate, and the solid that remains is called the residue. There are
two common types of filtration used in the laboratory: gravity and vacuum. Gravity filtration
utilizes a conical funnel with a larger piece of filter paper that is folded to fit. Vacuum filtration
uses a funnel with a flat bottom, called a Buchner funnel, with a piece of filter paper the same
size as the bottom of the funnel. In either method, if the mass of the residue is desired, the
filter paper is weighed before the procedure so that the mass can be easily obtained once the
residue has dried. Following are directions for both types of filtration.
Gravity Filtration
1. Place a funnel in a ring clamped to a stand, and a beaker underneath to collect the filtrate.
2. Fold a piece of filter paper in half once, and then in half again, forming a quarter of a circle
containing four layers of paper. Open the filter paper with three sheets on one side and the
fourth on the other side. The paper should resemble a cone. From the side with the three layers,
rip a small corner of the filter paper from the corner of the outer two layers so that the inner
layer of filter paper can adhere to the funnel once it is wet.
3. Place the filter paper in the funnel, and moisten it with the solvent to be used in the filtration.
4. Using a glass rod as a guide, slowly pour the mixture into the filter paper. Be careful not to
touch the filter paper with the rod, as it could rip. Also be careful not to allow the liquid to
overfill the filter paper, as some of the solid could pass around the paper and be collected with
the filtrate. Wait for the first amount to filter before adding a second.
5. Once all of the mixture has been filtered, rinse the container by adding 5-10 mL of solvent.
Swirl and filter. Rinse a second time with another 5-10 mL of solvent.
Vacuum Filtration
1. Place a Buchner funnel in a vacuum filtration flask (an Erlenmeyer flask with an arm to
allow connection to the vacuum) that is clamped to a ring stand. Connect the arm of the flask
to the trap connected to the vacuum pump, vacuum line on the benchtop, or aspirator.
2. Place a piece of filter paper in the funnel and moisten it with the solvent. Turn on the
vacuum, and be certain the filter paper is securely held on the bottom of the funnel.
3. Slowly pour the liquid onto the center of the paper. Be careful not to add too much liquid,
as the vacuum seal could break allowing some of the solid to pass around the paper.
4. Once the liquid has been added, transfer the solid to the center of the filter paper using a
rubber policeman as a guide.
5. Rinse the container (and rubber policeman) twice by adding 5-10 mL of solvent, swirling,
and adding slowly to the funnel each time. Allow the vacuum to continue for a few minutes to
assist with the drying of the residue. A rinse with acetone is sometimes used to assist as well.
Bergen Community College 141 General Chemistry II Laboratory
