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Chapter 2

Lecture Outline

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Chapter 2

The Chemical Basis of Life, I: Atoms, Molecules, and Water

Key Concepts:

• Atoms

• Chemical Bonds and Molecules

• Properties of Water

• pH and Buffers

Atoms

• The smallest functional units of matter that form all chemical substances

• Cannot be further broken down into other substances by ordinary means

• Each specific type of atom is a chemical element

Three subatomic particles

Protons

• positive charge (+) • found in nucleus

Neutrons

• neutral • found in nucleus

Electrons

• negative charge (−) • found in orbitals

Protons and electrons are present in equal numbers, giving the atom no net charge

The number of neutrons can vary

Table 2.1

Table 2.1 Characteristics of Major Subatomic Particles

Particle Location Charge Mass relative to electron

Proton Nucleus +1 1,836

Neutron Nucleus 0 1,839

Electron Around the nucleus

−1 1

Electrons occupy orbitals

Scientists initially visualized an atom as a miniature solar system

• This is an oversimplified but convenient image

Electrons travel within regions surrounding the nucleus (orbitals) in which the probability of finding that electron is high

Can be depicted as a cloud

Figure 2.1

Figure 2.2 panels 1 to 3

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Figure 2.2 panels 4 to 6

Orbitals • s orbitals are spherical

• p orbitals are propeller or dumbbell shaped

• Each orbital can hold only 2 electrons

Orbital name 1s 2s 2p

Number of electrons per electron shell

2 2 per orbital; 8 total

2 per orbital; 8 total

Orbital shape

Spherical First orbital: spherical

Second to fourth orbital: dumbbell- shaped

Electron Shells

Atoms with more electrons have orbitals within electron shells that are at greater and greater distances from the center of the nucleus

1st shell

• 1 spherical orbital (1s) - holds one pair of electrons

2nd shell

• 1 spherical orbital (2s) - holds one pair of electrons

• 3 dumbbell-shaped orbitals (2p) - three pairs of electrons

• Can hold four pairs of electrons = 8 electrons

Example: Nitrogen atom

7 protons and 7 electrons

2 electrons fill 1st shell

• 2 in the 1s orbital

5 electrons in 2nd shell

• 2 fill the 2s orbital • 1 in each of the three 2p orbitals

Note: the outer 2nd shell is not full

Electrons in the outer shell available to combine with other atoms are called valence electrons

Figure 2.4 a

a) Simplified depiction of a nitrogen atom

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Figure 2.4 b

b) Nitrogen atom showing electrons in orbitals

Protons

Number of protons is what distinguishes one element from another

Atomic number

• Equals number of protons

• Also equal to the number of electrons in the atom so that the net charge is zero

Periodic table

• Organized by atomic number

• Rows correspond to number of electron shells

• Columns, from left to right, indicate the numbers of electrons in the outer shell (the number of valence electrons)

• Similar properties of elements within a column occur because they have the same number of electrons in their outer shells, and therefore they have similar chemical bonding properties

Figure 2.5

Atomic mass

Protons and neutrons are nearly equal in mass, and both are more than 1,800 times the mass of an electron

Atomic mass scale indicates an atom’s mass relative to the mass of other atoms

Most common form of carbon has six protons and six neutrons, is assigned an atomic mass of exactly 12

• Hydrogen atom (atomic mass of 1) has 1/12 the mass of a carbon atom

• Magnesium atom (atomic mass of 24) has twice the mass of a carbon atom

Atomic mass in relation to mass of an electron

Table 2.1 Characteristics of Major Subatomic Particles

Particle Location Charge Mass relative to electron

Proton Nucleus +1 1,836

Neutron Nucleus 0 1,839

Electron Around the nucleus

−1 1

Mass versus weight

Weight is derived from the gravitational pull on a given mass

A man weighs 154 pounds on Earth

• On the moon he weighs about 25 pounds

• On a neutron star’s surface he would weigh 21 trillion pounds

His mass is the same in all locations

Units

Dalton

• Unit of measurement for atomic mass

• Also known as atomic mass unit (amu)

• One Dalton (Da) equals 1/12 the mass of a carbon atom

• Carbon has an atomic mass of 12 Daltons

Mole

• 1 mole of any element contains the same number of atoms— 236.022 10

• Avogadro’s number

Isotopes

• Multiple forms of an element that differ in the number of neutrons

• 12 C contains 6 protons and 6 neutrons

• 14

C contains 6 protons and 8 neutrons

• Atomic masses are averages of the weights of different isotopes of an element

Figure 2.6

©Steven Needell/Science Source

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Hydrogen, oxygen, carbon, & nitrogen

Make up about 95% of the atoms in living organisms

• Hydrogen and oxygen occur primarily in water

• Nitrogen is found in proteins

• Carbon is the building block of all living matter

Mineral elements - less than 1%

Trace elements - less than 0.01%

• Yet they are essential for normal growth and function

Table 2.2

Table 2.2 Chemical Elements Essential for Life in Many Organisms*

Most abundant in living organisms (approximately 95% of total mass)

Element Symbol % Human body mass

% All atoms in human body

Oxygen O 65 25.5

Carbon C 18 9.5

Hydrogen H 9 63.0

Nitrogen N 3 1.4

Mineral elements (less than 1% of total mass)

Calcium Ca Potassium K

Chlorine Cl Sodium Na

Magnesium Mg Sulfur S

Phosphorus P

Trace elements (less than 0.01% of total mass)

Boron B Manganese Mn

Chromium Cr Molybdenum Mo

Cobalt Co Selenium Se

Copper Cu Silicon Si

Fluorine F Tin Sn

Iodine I Vanadium V

Iron Fe Zinc Zn

*Although these are the most common elements in living organisms, many other trace and mineral elements have reported functions. For example, aluminum is believed to be a cofactor for certain chemical reactions in animals, but it is generally toxic to plants.

Chemical Bonds and Molecules

Molecule • Two or more atoms bonded together

Molecular formula

• Contains chemical symbols of the elements in the molecule (C6H12O6)

• Subscript indicates how many of each atom are present (H2O has two hydrogens, 1 oxygen)

Compound

• Any molecule composed of two or more elements • N2 and O2 are examples of molecules that are not

compounds

Three types of bonds

Covalent Bond

• Electrons are shared to fill valence shells • Can be polar covalent or nonpolar covalent

Hydrogen Bond

• Hydrogen atom from one polar molecule is attracted to an electronegative atom from another molecule

Ionic Bond

• Electrons are transferred, forming ions that are attracted to each other

Covalent bonds

Atoms share a pair of electrons

Occurs between atoms with unfilled valence electron shells

Covalent bonds are strong chemical bonds, because the shared electrons behave as if they belong to each atom

Can share …

• 1 pair of electrons – single bond, example H-F • 2 pairs of electrons – double bond, example O=O • 3 pairs of electrons – triple bond, example N≡N

Figure 2.7

Octet rule

• Atoms are stable when their outer shell is full

• For many atoms, the outer shell is filled with 8 electrons (“the octet rule”)

• An exception is hydrogen, which fills its outer shell with just 2 electrons

Figure 2.8

Figure 2.9

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Polar covalent bonds

• Form between atoms of different electronegativity (attraction to electrons)

• Shared electrons are more likely to be close to the more electronegative atom

• The unequal distribution of electrons creates a polarity (difference in electric charge) across the molecule

Water has polar covalent bonds

• The classic example of polar covalent bonds

• Electrons tend to be near the more electronegative oxygen atom rather than the less electronegative hydrogen atoms

• Water molecule has a partial negative charge

 δ around the oxygen and a partial positive charge  δ around the hydrogens

Figure 2.10

Nonpolar covalent bonds

• Between atoms with similar electronegativities (attraction to electrons)

• Equal sharing of electrons

• No charge difference across molecule

Hydrogen bonds The hydrogen atom from one polar molecule is attracted to an electronegative atom of another

Represented as dashed or dotted lines

Individually, these are weak bonds that can form and break easily

Collectively, many H bonds can be strong overall

• Holds DNA strands together

Figure 2.11 a and b

Ionic bonds

An ion is an atom or molecule that has gained or lost one or more electrons

• Cations – have a net positive charge (+)

• Anions – have a net negative charge (−)

Ionic bond occurs when a cation binds to an anion by electrostatic attraction

Ionic compounds are called salts

• Example NaCl, KCl, CaCl2

Figure 2.12 a and b

a) Formation of ions and an ionic bond

b) Sodium chloride (NaCl) crystals

Molecules May Change Their Shapes

• Atoms combine to form a molecule with three dimensional shape

• The shape is determined by the arrangement and number of bonds between atoms

• Angles that form between atoms give molecules specific shapes

• Covalent bonds are not rigid and rotation around single covalent bonds allows molecules to change shape

Figure 2.13

Figure 2.14

Free radicals

• Highly reactive molecules

• Can form by exposure to radiation and some toxins

• Can cause cell damage

• Can kill invading bacteria

• Benefits of antioxidants

Chemical Reactions

When one or more substances are changed into other substances

• Reactants → products

Properties of chemical reactions

• Require a source of energy

• In living organisms, they often require an enzyme as catalyst

• Tend to proceed in a particular direction but will eventually reach equilibrium

• Occur in liquid (water)

Properties of Water

Solution = solutes in a solvent

• Solutes are dissolved substances

• Solvent is the liquid

In an aqueous solution, water is the solvent

Ions and molecules with polar covalent bonds will dissolve in water

These are hydrophilic

Figure 2.16

Solutes Hydrophilic – “water-loving”

• Readily dissolve in water

• Molecules with ionic and/or polar covalent bonds

Hydrophobic – “water-fearing”

• Do not dissolve in water

• Nonpolar molecules like hydrocarbons, oils

Amphipathic – “both loves”

• Have both polar/ionized and nonpolar regions

• May form micelles in water

• Detergent is an amphipathic molecule

Figure 2.17

• Polar (hydrophilic) regions at the surface of the micelle

• Nonpolar (hydrophobic) ends are oriented toward the interior of the micelle

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Measuring solutions

Concentration

• Amount of a solute dissolved in a unit volume of solution

• 1 gram of NaCl dissolved in 1 liter of water = 1 gram/Liter

Molarity

• Number of moles of a solute dissolved in 1 Liter of water

• 1 mole of a substance is the amount of the substance in grams equal to its atomic or molecular mass

H2O in three states of matter

Solid (ice), liquid (water), and gas (water vapor)

Changes in state, such as changes between the solid, liquid, and gas states of H2O, involve an input or release of energy

• Heat of vaporization – energy to boil

• Heat of fusion – energy to melt

Specific heat is the amount of heat energy to raise

temperature 1 Celsius

Water is extremely stable as a liquid, due to high heats of vaporization and fusion, and high specific heat

Figure 2.18

Colligative properties of water

Temperature at which a solution freezes or boils is influenced by amounts of dissolved solutes

Addition of solutes to water

• lowers the freezing point below 0 Celsius

• raises the boiling point above 100 Celsius

Some animals produce antifreeze molecules, lowering the freezing point of body fluids to prevent blood and cells from freezing

Not just a solvent

Water has many important functions in living organisms:

• Participates in chemical reactions (hydrolysis or condensation)

• Provides force or support

• Removes toxic waste components

• Evaporative cooling

• Cohesion and adhesion

• Surface tension

• Lubrication

Figure 2.19

b: ©Aaron Haupt/Science Source; d: ©Chris McGrath/Getty Images; e: ©Dana Tezarr/Getty Images; f: ©Gallo Images-Anthony Bannister/DigitalVision/Getty Images; g: ©Matti Suopajarvi/mattisj/Getty Images

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Acids and Bases 1

• Pure water ionizes to a very small extent into

hydrogen ions  H and hydroxide ions  OH

• In pure water 7 7 14

H OH 10 M 10 M 10 M                    

Acids and Bases 2

Acids are molecules that release hydrogen

ions in solution

• A strong acid releases more H 

than a weak acid

Bases lower the H concentration

• Some release OH

• Others bind H

The pH scale

• 10

pH log Η     

• Acidic solutions are pH 6 or below

• pH 7 is neutral

• Alkaline solutions are pH 8 or above

Figure 2.20

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Effects of pH

The pH of a solution can affect

• The shapes and functions of molecules

• The rates of many chemical reactions

• The ability of two molecules to bind to each other

• The ability of ions or molecules to dissolve in water

Buffers • Organisms usually tolerate only small changes

in pH

• Buffers help to maintain a constant pH

• An acid-base buffer system can shift to

remove or release H to adjust for changes in

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