2 Small Molecules and the Chemistry of Life. 2 Small Molecules and the Chemistry of Life 2.1 How Does Atomic Structure Explain the Properties of Matter?

2Small Molecules and the

Chemistry of Life

2 Small Molecules and the Chemistry of Life

2.1 How Does Atomic Structure Explain the Properties of Matter?

2.2 How Do Atoms Bond to Form Molecules?

2.3 How Do Atoms Change Partners in Chemical Reactions?

2.4 What Makes Water So Important for Life?

2 Small Molecules and the Chemistry of Life

Opening Question:Can isotope analysis of water be used to detect climate change?

Elements have naturally occurring variants called isotopes, which have slightly different weights.

Figure 2.1 The Helium Atom

All matter is composed of atoms.


Electrons—negligible mass; negative charge

In the nucleus:

Protons—have mass; positive charge

Neutrons—have mass; no charge


Atoms have volume and mass.

Mass of one proton or one neutron =

1 atomic mass unit (amu), or 1 dalton,

or 1.7 × 10–24 grams.

Mass of one electron = 9 × 10–28

(usually ignored).


Element: pure substance containing only one kind of atom.

The number of protons identifies an element.

Number of protons = atomic number.


The number of protons and electrons determines how an element behaves in chemical reactions.

Elements are arranged in the periodic table.

Figure 2.2 The Periodic Table (Part 1)

Figure 2.2 The Periodic Table (Part 2)


All elements except hydrogen have one or more neutrons.

Mass number = number of protons + number of neutrons.

Mass number = mass of atom in daltons.


Isotopes: forms of an element with different numbers of neutrons, and thus different mass numbers.

Example:12C has 6 neutrons13C has 7 neutrons14C has 8 neutrons


The isotopes of some elements have special names:


Atomic weight (relative atomic mass): average of mass numbers of isotopes in their normally occurring proportions.

Atomic weight of hydrogen = 1.0079

Given as a range = 1.00784 – 1.00811


Some isotopes are unstable:

Radioisotopes give off energy in the form of alpha, beta, and gamma radiation from the nucleus.

This radioactive decay transforms the atom, sometimes resulting in a change in the number of protons.


The radiation can be used to detect the presence of radioisotopes.

Radioisotopes can be incorporated into molecules to act as a tag or label.

Molecule: stable association of atoms.

Figure 2.3 Tagging the Brain


Radioisotopes are useful in research and in medicine, but radiation can damage cells and tissues.

Radiation is sometimes used to treat cancer.


The number of electrons determines how atoms will interact.

Chemical reactions involve changes in the distribution of electrons between atoms.


Locations of electrons are described by orbitals.

Orbital: region where an electron is found at least 90% of the time.

Orbitals have characteristic shapes and orientations and can be occupied by two electrons.

Orbitals are filled in a specific sequence.

Figure 2.4 Electron Shells and Orbitals


Orbitals occur in a series called electron shells, or energy levels.

First shell: one s orbital (holds 2 electrons).

Second shell: 1 s and 3 p orbitals (holds 8 electrons).

Additional shells: 4 orbitals (each holds 8 electrons).


The outermost electron shell (valence shell) determines how the atom behaves.

If the outermost shell is full, the atom is stable; it won’t react with other atoms.

Reactive atoms have unpaired electrons in their outermost shell.

Figure 2.5 Electron Shells Determine the Reactivity of Atoms


Reactive atoms can share electrons, or lose or gain electrons, resulting in atoms being bonded together to form molecules.

Octet rule: tendency of atoms to form stable molecules resulting in 8 electrons in their outermost shells.


Chemical bond: the attractive force that links atoms together to form molecules.

There are several kinds of chemical bonds.

Table 2.1


Covalent bonds: atoms share one or more pairs of electrons so that the outer shells are filled.

Figure 2.6 Electrons Are Shared in Covalent Bonds


Compound: a molecule made up of two or more elements bonded together in a fixed ratio.

Example: H2O

The molecular weight of a compound is the sum of the atomic weights of all atoms in the molecule.

Table 2.2


Covalent bonds are very strong—a lot of energy is required to break them.

At temperatures in which life exists, the covalent bonds of biological molecules are very stable (although changes can still occur).


Orientation of bonds:

The length, angle, and direction of bonds between any two elements are always the same.

Example: Methane always forms a tetrahedron.

Figure 2.7 Covalent Bonding Can Form Compounds


The shape of molecules can change as atoms rotate around a covalent bond:

In-text art, pg 7, bottom


Covalent bonds can be:

• Single, sharing 1 pair of electrons

• Double, sharing 2 pairs of electrons

• Triple, sharing 3 pairs of electrons C C

C H

N N


Sharing of electrons in a covalent bond is not always equal.

Electronegativity: the attractive force that an atomic nucleus exerts on electrons.

Electronegativity depends on the number of protons and the distance between the nucleus and electrons.

Table 2.3


Polar covalent bond: electrons are drawn to one nucleus more than the other because that atom has greater electronegativity.

Nonpolar covalent bond: electrons are shared equally (atoms have similar electronegativity).

Figure 2.8 Water’s Covalent Bonds Are Polar


Because of unequal sharing, a molecule with a polar bond has a slightly negative charge on one end and a slightly positive charge on the other.

The partial charges result in polar molecules or polar regions of large molecules.


When one atom is much more electronegative than the other, a complete transfer of electrons may occur.

This results in two ions with full outer shells.

Figure 2.9 Formation of Sodium and Chloride Ions


Ions: electrically charged particles formed when atoms lose or gain electrons.

Cations—positive

Anions—negative

Complex ions: groups of covalently bonded atoms that carry a charge, e.g., NH4

+ and SO42–.


Ionic attractions: bonds formed by the electrical attraction of positive and negative ions.

Salts are ionically bonded compounds.


In a solid, ions are close together and the ionic attractions are strong.

In water, the ions are far apart and the attraction is much weaker.

Ions can interact with polar molecules (e.g., salts dissolve in water).

Figure 2.10 Water Molecules Surround Ions


Hydrogen bonds: attraction between the δ– end of one molecule and the δ+ hydrogen end of another molecule.

Hydrogen bonds form between water molecules and are important in the structure of DNA and proteins.

Figure 2.11 Hydrogen Bonds Can Form between or within Molecules


Polar molecules that form hydrogen bonds with water are hydrophilic (“water-loving”).

Nonpolar molecules, such as hydrocarbons that interact with each other but not with water, are hydrophobic (“water-hating”).

Figure 2.12 Hydrophilic and Hydrophobic


van der Waals forces: attractions between nonpolar molecules that are close together.

Individual interactions are brief and weak, but can be substantial when summed over a large molecule.


Chemical reactions occur when atoms collide with enough energy to combine or change their bonding partners.

Figure 2.13 Bonding Partners and Energy May Change in a Chemical Reaction


This is an oxidation–reduction reaction.

Electrons and protons are transferred from propane (the reducing agent) to oxygen (the oxidizing agent) to form water.

C3H8 + 5 O2 3 CO2 + 4 H2O + Energy

Reactants Products


In all chemical reactions, matter and energy are neither created nor destroyed.

Energy: capacity to do work or the capacity for change. Energy usually changes form during chemical reactions.


Some chemical reactions release energy, whereas others require it. These reactions are often linked.

Many reactions in living cells are oxidation–reduction reactions.


Water has a unique structure and special properties:

• Polar molecule

• Forms hydrogen bonds

• Tetrahedral shape

In-Text Art, Ch. 2, p. 32


Ice floats: Each molecule is hydrogen-bonded to four other molecules in a rigid, crystalline structure.

Solid water is less dense than liquid water.

Figure 2.14 Hydrogen Bonding and the Properties of Water


Ice requires a lot of energy to melt, to break the hydrogen bonds.

When water freezes, a lot of energy is released to the environment.


Water has high specific heat—the amount of heat energy required to raise the temperature of 1 gram of water by 1°C.

The high heat capacity of water helps moderate climate and ocean temperatures.


Water has a high heat of vaporization: amount of energy required to change water from a liquid to a gas state.

The heat energy must be absorbed from the environment in contact with the water and results in cooling—e.g., when humans sweat.


Cohesion: water molecules resist coming apart from one another.

• Helps water move through plants

• Results in surface tension—water molecules at the surface are hydrogen-bonded to other water molecules below them.

Figure 2.15 Water in Biology


A solution is a substance (solute) dissolved in a liquid (solvent).

Many important biochemical reactions occur in aqueous solutions—water is the solvent.


Qualitative analysis: identification of the substances involved in chemical reactions.

Quantitative analysis: measuring concentrations of the substances.


Mole: The amount of a substance (in grams) that is numerically equal to its molecular weight.

1 mole of sodium ion Na+ = 23 g

1 mole of hydrogen gas H2 = 2 g

1 mole of table sugar (C12H22O11) = 342 g


The number of molecules in 1 mole is constant for all substances.

One mole contains 6.02 × 1023 molecules = Avogadro’s number


A 1 molar solution (1M) is 1 mole of a substance dissolved in water to make 1 liter of solution.

Example: 342 g table sugar dissolved in 1 liter of water.

In living tissues, substances occur in micromolar (μM), and smaller, concentrations.


When acids dissolve in water, they release hydrogen ions—H+ (protons).

H+ ions can attach to other molecules and change their properties.

Bases accept H+ ions.


HCl is a strong acid; it fully ionizes in water.

HCl H+ + Cl–


Organic acids have a carboxyl group:

—COOH —COO– + H+

Weak acids: not all the acid molecules dissociate into ions.


NaOH is a strong base (it accepts H+).

In water:

NaOH Na+ + OH–

OH– + H+ H2O


Weak bases:

• Bicarbonate ion

HCO3– + H+ H2CO3

• Ammonia

NH3 + H+ NH4+

• Compounds with amino groups

–NH2 + H+ NH3+


Acid–base reactions may be reversible.

CH3COOH CH3COO– + H+

Ionization of strong acids and bases is irreversible.

Ionization of weak acids and bases is somewhat reversible.


Water has a slight tendency to ionize.

H2O H+ + OH–

Two water molecules are actually involved:


Water acts as both a weak acid and a weak base.

Ionization of water is important to all organisms because of the abundance of water in living systems and the reactive nature of the H+ ions.


pH = negative log of the molar concentration of H+ ions.

pH = –log[H+]

H+ concentration of pure water is 10–7 M, its pH = 7.

Lower pH numbers mean higher H+ concentration, or greater acidity.


1 M HCl solution: H+ concentration = 1 M

pH = 0

1 M NaOH solution: H+ concentration = 10–14 M.

pH = 14

Figure 2.16 pH Values of Some Familiar Substances


pH influences the rate of biological reactions and can change the 3-dimensional structure of biological molecules, which impacts function.

Organisms use many mechanisms to minimize change in pH in their cells and tissues.


Living organisms maintain constant internal conditions (homeostasis).

Buffers help maintain constant pH.

A buffer is a weak acid and its corresponding base.

HCO3– + H+ H2CO3

Figure 2.17 Buffers Minimize Changes in pH


Buffers illustrate the law of mass action: addition of reactant on one side of a reversible equation drives the system in the direction that uses up that compound.

Addition of an acid drives the reaction in one direction; addition of a base drives the reaction in the other direction.

2 Answer to Opening Question

Heavy isotopes of H and O fall as precipitation more readily than the lighter isotopes.

As water vapor moves towards the poles, it becomes enriched in the lighter isotopes.

The cooler the climate, the lower the ratio of heavy to light isotopes because more water precipitates.

2 Answer to Opening Question

In ice cores, the ratio of heavy to light isotopes changes over time, allowing reconstruction of past climatic change.

2 Small Molecules and the Chemistry of Life. 2 Small Molecules and the Chemistry of Life 2.1 How Does Atomic Structure Explain the Properties of Matter?

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