Element	Symbol	Atomic Number	Atomic Weight
Hydrogen	H	1	1
Carbon	C	6	12
Nitrogen	N	7	14
Oxygen	O	8	16
Sodium	Na	11	23
Magnesium	Mg	12	24.3
Phosphorus	P	15	31
Sulfur	S	16	32.1
Chlorine	Cl	17	35.5
Potassium	K	19	39.1
Calcium	Ca	20	40.1
Iron	Fe	26	55.8
Cobalt	Co	27	58.9
Copper	Cu	29	63.5
Zinc	Zn	30	65.4
Iodine	I	53	126.9

All atoms have the same fundamental structure consisting of a center, or atomic nucleus, and surrounding shells (Figure 2.1), but because of the different numbers of subatomic particles, each element has its own characteristic atomic structure. Located in the center of the atom is the atomic nucleus, which consists of positively charged particles called protons, and particles without charge called neutrons. The atomic weight (atomic mass) of an atom is equal to the sum of protons and neutrons. The atomic number indicates the number of protons in the atomic nucleus. Surrounding the atomic nucleus in shells are negatively charged subatomic particles called electrons. Electrons travel around the nucleus at high speed and occupy positions in a volume of space called an orbital or electron cloud. These orbitals form an energy level also referred to as shells, in which the electrons usually remain. Electrons fill the orbitals and shells in pairs, and each orbital within a shell can carry two electrons.

FIGURE 2.1 Models of atomic structure.
These simplified diagrams of hydrogen and carbon atoms show the atomic nucleus containing protons and neutrons in the center of the atom, and the electrons in the surrounding shells.

The nucleus of a given atom is surrounded by successive shells spaced further and further away from the nucleus. The energy level of electrons increases with the distance of their shells from the nucleus. The innermost (first) shell can be occupied by up to two electrons within one orbital, the second shell with up to eight within four orbitals, and each consecutive shell can potentially hold more electrons. However, most elements with biological significance need eight electrons to fill the outermost shell. The shells always fill sequentially from the inside out: two electrons in the first shell, eight in the next, and so on. For example, carbon with 6 electrons carries 2 electrons in the first shell and 4 in the second shell, and sodium with 11 electrons has 2 electrons in the first shell, 8 in the second, and 1 in the third shell (Figure 2.2).

FIGURE 2.2 Atomic structure of several different atoms important in living systems.
The diagrams illustrate that the number of protons equals the number of electrons in each of these atoms, which are found in association with living systems.

In general, the number of protons and electrons of an atom are equal, making the atom an electrically neutral unit. The stability of an atom depends on the number of electrons in the outermost shell. For example, an atom is most stable if the outermost shell is filled to its capacity. Hydrogen is the simplest element, with the atomic number of 1, and therefore has one electron in the outermost shell. Helium, with the atomic number of 2, has two electrons in the outermost shell. This shell is fully occupied and is stable. Helium atoms will not react with each other and also cannot combine with atoms of other elements. Helium is therefore called an inert gas.

If the outermost shell is not complete, the atom can participate in a chemical reaction and form a chemical bond. Electrons in the outermost shell of an atom that are available for chemical bonding are called valence electrons. These electrons determine what kind of chemical bonds, if any, the atom can form.

Isotopes are atoms with the same number of protons but a different number of neutrons. The atomic number of isotopes is unchanged because the number of protons remains the same and only the atomic weight is different. For example, the element hydrogen has two isotopes (Figure 2.3):

FIGURE 2.3 Isotopes of hydrogen.
Hydrogen has one proton only, deuterium has one proton and one neutron, and tritium has one proton and two neutrons.

• Hydrogen (one proton)

• Deuterium (one proton and one neutron)

• Tritium (one proton and two neutrons)

Radioisotopes are unstable because of their imbalance of energy within the nucleus. When the nucleus loses a neutron it gives off energy and is said to be radioactive. Radioactivity is the release of energy and matter that results from changes in the nucleus of an atom. Tritium is an example of a radioactive isotope that is used in research and clinical procedures.

MEDICAL HIGHLIGHTS

Radioactive Isotopes

Though radiation from radioactive isotopes can be dangerous to living organisms by causing tissue damage, several of the isotopes have medical uses and are important to research as well. Microbiological studies rely on the use of radioisotopes to identify proteins, label the surface components of bacteria, and trace the transcription and translation steps in nucleic acid and protein synthesis. Furthermore, radioisotopes are used in immunological studies and in the identification of specific pathogens.

In therapeutics radioactive isotopes are used to label or tag molecules, which can then be located in a patient’s body tissues by health professionals using specialized equipment. This methodology enables physicians to locate damaged tissue for more precisely focused treatment, including the targeting and killing of certain cancers.

Ions

Ions are electrically charged atoms, molecules, or subatomic particles that are formed when one or more valence electrons are transferred from one atom to another (see Formation and Classification of Chemical Bonds and Forces, below). If an atom loses one or more electrons to another atom, it becomes positive (+), whereas the atom that gains the electron becomes negative (−). Positively charged ions are called cations, and in an electric field move toward the negative pole, the cathode. Negatively charged ions, referred to as anions, move toward the positive pole, or anode, of an electric field.

A substance that dissociates into free ions when dissolved in a solvent such as water is called an electrolyte. The solvent in which it is dissolved can then conduct an electric current and is referred to as an electrically conductive medium. Because these solvents contain ions or electrolytes they are called ionic solutions. Chemically they are acids, bases, or salts (see Acids, Bases, and the pH Scale, below).

Several cations and anions (Table 2.2) are important components of higher life forms. All of these higher life forms require a complex electrolyte balance, called an osmotic gradient, between their intercellular and extracellular fluid compartments (see Chapter 3, Cell Structure and Function). This maintenance of a precise internal balance of electrolytes to maintain the osmotic gradient is called homeostasis. It is required to regulate the hydration, blood pH, and nerve and muscle function of an organism.

TABLE 2.2

Common Ions in Living Organisms

Cations	Anions
Sodium (Na⁺)	Chloride (Cl⁻)
Potassium (K⁺ )	Bicarbonate (HCO₃⁻)
Calcium (Ca²⁺ )	Phosphate (PO₄³⁻)
Magnesium (Mg²⁺ )	Sulfate (SO₄²⁻)

MEDICAL HIGHLIGHTS

Rehydration Therapy

To treat microbial diseases in which diarrhea, vomiting, starvation, or other extreme conditions result in dehydration, rehydration therapy is needed and may be instituted orally or intravenously. In addition to an intravenous drip, electrolyte drinks containing sodium and potassium salts can be taken to replenish the body’s water and electrolyte balance. At present, sports drinks, which are electrolyte drinks to which large amounts of carbohydrates such as glucose are added to provide energy, are readily obtainable and have become popular. However, because of the high concentration of sugar in these drinks, they are not recommended for regular use by children. Instead, specially formulated pediatric electrolyte solutions are available.

Chemical Bonds and Molecules

Molecules are two or more atoms linked together by chemical bonds formed by their valence electrons. As stated above, atoms are most stable when their outermost shell is filled with eight electrons. This is the octet principle. The number of bonds a single atom can have is dependent on how many electrons are needed to complete the outermost shell. Hydrogen with one electron in the outermost shell can form one chemical bond; oxygen with eight electrons (2 + 6) (and therefore six in the outermost shell) can form two bonds; and carbon with six electrons (2 + 4) (four in the outermost shell) can form up to four chemical bonds to fill the outermost shell. When the outermost shell is not completely occupied with electrons, the atom has the tendency to interact with other atoms forming chemical bonds to achieve higher stability. These atoms then become stable and cannot react with others.

Formation and Classification of Chemical Bonds and Forces

Molecules made from atoms of different elements are called compounds. Compounds are new chemicals with properties that are different from those of the atoms of which they are composed. Groups of atoms that consistently form specific groups within compounds are referred to as functional groups. They have specific characteristics that are different from those of the individual participating atoms of that given group. Some molecules have more than one functional group, which may differ from one another. The most common functional groups found in molecules important to living organisms are shown in Table 2.3.

TABLE 2.3

Common Functional Groups in Living Organisms

Functional Group	Formula	Functional Group	Formula
Acetyl	CH₃	Ethyl	C₂H₅
Aldehyde	CHO	Hydroxyl	OH
Amino	NH₂	Keto	CO
Ammonium	NH₄	Methyl	CH₃
Bicarbonate	HCO₃	Nitrate	NO₃
Carbonate	CO₃	Phosphate	PO₄
Carboxyl	COOH	Sulfate	SO₄

The principal types of chemical bonds formed by the interactions of atoms and/or molecules are covalent bonds, ionic bonds, hydrogen bonds, and those based on van der Waals forces. Covalent and hydrogen bonds occur between atoms to form a molecule, whereas hydrogen bonds and van der Waals forces are intermolecular connections. Chemical bonds vary in their strength but, in general, covalent bonds are considered the strongest bond, followed by ionic bonds, hydrogen bonds, and—with the weakest connection—the van der Waals forces.

Covalent bonds result from a sharing of electrons between two atoms of the same element or between atoms of different elements. In bonds between identical atoms such as oxygen and hydrogen, the electrons are shared equally by each atom. Covalent bonds usually are the strongest chemical bonds. Because the electrons are equally distributed, the resulting molecule is nonpolar and the bond is called a nonpolar covalent bond (Figure 2.4, A). Carbon atoms play a significant role in large organic molecules because they form stable nonpolar covalent bonds with each other. This stable framework is the backbone of organic carbon-based molecules, providing the chemical foundation of organic chemistry and of life.

FIGURE 2.4 Covalent bonds: a sharing of electrons.
A, Nonpolar covalent bond between two hydrogen atoms (H₂), formed by an equal sharing of electrons. B, Polar covalent bond between an oxygen atom and two hydrogen atoms (H₂O), formed by an unequal sharing of electrons.

The covalent bonds between atoms of two different-sized elements are polar covalent bonds, in which the electrons are unequally distributed because they are pulled toward the larger atom. As a result, one end of the molecule becomes more negative compared with the other end (Figure 2.4, B). Oxygen, nitrogen, and phosphorus atoms have a tendency to form polar covalent bonds. Polar covalent bonds are somewhat weaker than nonpolar covalent bonds. Coordinate covalent bonds, such as occurs in the formation of the ammonium ion from ammonia, are formed when both electrons are from one atom. The molecule no living organism can exist without is water, in which the atoms hydrogen and oxygen are held together by polar covalent bonds. Some properties of water result from this type of bond.

Depending on the number of electrons shared, molecules can be formed from a single covalent bond by sharing one pair of electrons, such as the bond between hydrogen atoms. Single covalent bonds are indicated by one solid line (HH). Double covalent bonds are formed by sharing two pairs of electrons, as seen between oxygen atoms. These bonds are indicated by two solid lines (OO). Triple covalent bonds may occur through the sharing of three pairs of electrons, such as between nitrogen atoms. These bonds are identified by three solid lines (NN).

Ionic bonds are formed when one or more electrons from one atom are transferred to another. If an atom loses one electron in the process it will have a charge of +1; if two electrons are lost the charge will be +2, because the protons in the nucleus will be unbalanced by the remaining electrons. The resulting anions and cations in an ionic bond are held together by attraction of their opposite charges and form an ionic compound. Ionic bonds can easily dissociate (break down) in water to form electrolyte solutions. For example, in water metals such as Na⁺, readily give up electrons, and nonmetals such as Cl⁻ readily take up electrons (Na⁺ + Cl⁻ → NaCl). If the water is evaporated, a solid crystal of NaCl, common table salt, is formed. Sodium with a total of 11 electrons (2 + 8 + 1) has only 1 electron in its outermost shell, whereas chlorine with a total of 17 electrons (2 + 8 + 7) only needs 1 electron to fill its outermost shell. The only electron in sodium’s outermost shell is therefore attracted to chlorine’s outermost shell and its transfer forms an ionic compound (Figure 2.5, A). The charged sodium chloride molecules and other salts form characteristic large crystal structures in which the atoms of the molecules alternate in a regular, geometric pattern (Figure 2.5, B). In water, NaCl readily dissociates to form an electrolyte.

FIGURE 2.5 Ionic bond.
A, The electron in the outermost shell of the sodium atom is transferred to the chlorine atom, resulting in a sodium (+) and a chloride (−) ion. B, Crystal of sodium chloride (table salt). The sodium and chloride ions alternate in a definite, regular, geometric pattern.

Hydrogen bonds are weak chemical bonds with only about 5% of the strength of covalent bonds. However, when many hydrogen bonds are formed between two molecules, the resulting union can be strong enough to be stable. These bonds are formed by attraction forces between charged atoms within a large molecule or between adjacent molecules (Figure 2.6). Hydrogen bonds always involve a hydrogen atom with a slight positive charge and an oxygen or nitrogen atom with a slightly negative charge. Although hydrogen bonds do not form molecules they can alter the shapes of molecules or hold together different molecules. Examples of hydrogen bonds include bonds between water molecules, acetic acid molecules, amino acid molecules, and nucleic acid molecules. Hydrogen bonds are always indicated by dotted lines (—). The attraction created by hydrogen bonds keeps water in the liquid state over a wide range of temperatures.

FIGURE 2.6 Hydrogen bond between water molecules.
The positively charged hydrogen portion of one water molecule is slightly attracted to the negatively charged oxygen portion of another water molecule. Hydrogen bonds are always indicated by dotted lines.

Van der Waals forces are the weakest of the intermolecular forces in all chemical reactions. The van der Waals force of attraction is inversely proportional to the seventh power of the interatomic distance, whereas the force of an ionic bond diminishes as the square of the distance. Therefore a very slight increase in the interatomic distance between atoms markedly reduces the van der Waals force of attraction. In terms of the “lock and key” concept of reactants and their receptors, such as between antigens and antibodies and between drugs and their receptors, van der Waals forces determine the final molecular arrangements that define selectivity and specificity properties. In other words, if there is an interaction (selectivity), how well it fits (specificity) is a function of the van der Waals forces. Van der Waals forces explain how a spider can hang upside down from a ceiling and why a gecko can hang by one toe from a glass surface. For drug–receptor interactions, ionic bonds based on electrostatic attraction probably come into play as the drug approaches the immediate vicinity of the receptor, followed by additional attraction based on van der Waals forces. Covalent bonding is more of a factor in longer lasting drug actions.

Types of Chemical Reactions

Pathways of chemical reactions trace metabolic activities in living organisms. Within these pathways, several chemical reactions occur that are essential for the survival of living organisms, including microbes. These reactions are as follows:

• Dehydration synthesis (condensation)

• Hydrolysis (decomposition)

• Endergonic (energy-requiring) reactions

• Exergonic (energy-producing) reactions

• Oxidation and reduction (redox)

• Exchange reactions

Dehydration synthesis, or condensation, is the formation of a larger compound (polymer) from smaller ones (monomers). Monomers are the unit molecules (building blocks) of these larger molecules, called polymers. These reactions require specific enzymes and the removal of water from the reactants, that is, a hydroxyl group (OH) is removed from one monomer and combined with a hydrogen (H) from the other (Figure 2.7). Enzymes (see Chapter 3, Cell Structure and Function) are biological catalysts and function to speed up the rate of chemical reactions without changing themselves. The synthesis of new compounds within a cell occurs during anabolism, which utilizes energy provided by catabolism (see Cellular Metabolism in Chapter 3, Cell Structure and Function). An example of synthesis is the production of complex sugars from simple sugars (see Carbohydrates, below).

FIGURE 2.7 Dehydration synthesis and hydrolysis.
Dehydration synthesis is an anabolic reaction (requires energy) producing polymers (large molecules) from monomers (small molecules) by removing water. Hydrolysis is a catabolic reaction (releasing energy) in which polymers are broken down into monomers. These reactions require water.

< ?xml:namespace prefix = "mml" />Glucose+glucose→maltose

Monomer+monomer→polymer

Hydrolysis (decomposition) breaks down large molecules (polymers) into their unit molecules (monomers). An example of hydrolysis is the breakdown of nutrient molecules such as carbohydrates, proteins, and lipids into smaller molecules during the digestive process. Hydrolysis is the reverse of dehydration synthesis and occurs during the metabolic process of catabolism. This complex reaction of breaking down polymers requires water and the resulting monomers can be used in cellular metabolism for the generation of energy.

Reactions that yield energy are called exergonic reactions. Reactions that utilize energy are endergonic reactions. Hydrolysis that occurs during catabolism is an exergonic reaction and it releases energy. Endergonic reactions require energy such as occurs in the dehydration synthesis of nutrient molecules during anabolism.

Redox (reduction–oxidation) reactions are chemical reactions in which atoms have their oxidation number (oxidation state) changed. An oxidation reaction does not occur without a reduction reaction happening at the same time. Oxidation is loss of electrons and is a common reaction in the production of cellular energy. The transfer of electrons is catalyzed by enzymes within the metabolic pathways. The electrical state of an atom is identified as the oxidation state or by its oxidation number. Atoms are neutral and their oxidation state is therefore zero. The oxidation state changes because of a loss or gain of electrons. Loss of electrons results in a positive oxidation state of an atom and a gain of electrons results in a negative oxidation state. Oxidation reactions are indicated as follows, using calcium as an example:

Ca→Ca2++2e−

Oxidation–reduction (electron transfer) reactions are coupled. In other words, oxidation reactions occur with reduction reactions (Figure 2.8). Reduction is the gain of electrons. It is also catalyzed by enzymatic reactions and is written as follows:

FIGURE 2.8 Oxidation–reduction (electron transfer) reactions.
Oxidation refers to the loss of electrons, whereas reduction involves gaining electrons.

O+2e−→O2−

LIFE APPLICATION

Bleach

The decolorizing action of bleaching agents is partly due to their ability to remove the electrons that are activated by light to produce the various colors. Most chemical bleaches contain sodium hypochlorite (NaOCl) or hydrogen peroxide (H₂O₂), which are both oxidizing agents. Hypochlorite (OCl⁻) is reduced to chloride ions and hydroxyl ions.

OCl−+2e−+HOH→Cl−+2OH−

Bleaches are sometimes combined with other compounds that are different from bleaches but are capable of absorbing wavelengths of ultraviolet light, which is invisible to the human eye. They convert these wavelengths to visible blue or blue-green light that is reflected, making the fabric appear brighter.

Exchange reactions transfer the same molecules in a reaction but in a different combination. In other words, the components of the reaction and the products remain the same but their combination results in a different product.

AB+CD→AD+CB

The balancing of chemical activities in living cells is a continuous, dynamic process. In general, however, breaking of bonds releases more energy than is required for synthesis. A living cell continuously undergoes hydrolysis and synthesis for the purpose of energy production such as occurs in cells during the breakdown of complex molecules (catabolism). Chemical bonds contain potential energy that is released when the bonds are broken. This energy is captured and used in reactions for the essential functions of a cell. Catabolic reactions are primarily oxidation reactions and are divided into different pathways, namely, glycolysis, the Krebs cycle (citric acid cycle), and the electron transport chain (see Chapter 3, Cell Structure and Function).

LIFE APPLICATION

Nitrogen Fixation

Nitrogen, the most abundant element in the atmosphere of the earth, is a vital element in many of the compounds essential to living systems. For example, in all green plants nitrogen is a primary nutrient. However, it must be modified by a process called nitrogen fixation before it can be used. Nitrogen fixation is a complex process that reduces nitrogen to nitrogen compounds such as ammonia, nitrate, and nitrogen dioxide. This process is performed by different nitrogen-fixing bacteria present in soil. Nitrogen fixation involves a number of oxidation–reduction reactions that occur sequentially to yield ammonium ions. Also, nitrogen in the form of ammonia is used by living systems in the synthesis of many organic compounds. Ammonia can be formed by the process of nitrification, during which nitrates and nitrites released by decaying organic matter are converted to ammonium ions by nitrifying bacteria, again present in the soil. This process can also be achieved by a series of oxidation–reduction reactions. Denitrifying bacteria, acting on ammonia and nitrates produced by decay, recycle these compounds to free nitrogen to complete what is called the nitrogen cycle.

Chemical Notations

Chemical compounds and reactions are shown by “chemical shorthand” or chemical notation. The rules of the chemical notation are as follows (Box 2.1):

BOX 2.1 Chemical Notation

Symbols

H: An atom of hydrogen	C: An atom of carbon
O: An atom of oxygen	N: An atom of nitrogen

Prefixes

2H: Two individual atoms of hydrogen

2O: Two individual atoms of oxygen

3C: Three individual atoms of carbon

4N: Four individual atoms of nitrogen

Number of Atoms in a Molecule

H₂: A molecule with two hydrogen atoms

O₂: A molecule with two oxygen atoms

C₆: A molecule with six carbon atoms

N₄: A molecule with four nitrogen atoms

Superscripts

Na⁺: A sodium atom that has lost one electron, resulting in a positively charged ion

Cl⁻ : A chlorine atom that has gained one electron, resulting in a negatively charged ion

K⁺: A potassium ion

Ca²⁺: A calcium atom that has lost two electrons

Chemical Equation

Balanced: 2H + O → H₂O

2H₂ + O₂ → 2H₂O

NaOH + HCl → NaCl + H₂O

Unbalanced: H₂O₂ → H₂O

2NaOH + HCl → NaCl + H₂O

• The abbreviation of an element represents one atom of that element and is its chemical symbol.

• The number before the chemical symbol is the number of atoms; the number before the chemical formula is the number of molecules.

• The subscript after the chemical symbol of an element shows the number of that atom in the molecule.

• The reaction of the chemicals describes the interaction of the participants, called the reactants. Chemical reactions form one or more products. Arrows in the formula indicate the direction of the reaction, from reactant to product. Arrows in both directions indicate a reversible chemical reaction that can go in either direction. For example,

2H+O→H2O

The equation indicates that two atoms of hydrogen and one atom of oxygen combine to form water (H₂O). Another example:

NaOH+HCl⇄NaCl+H2O

Here one molecule of sodium hydroxide (NaOH) and one molecule of hydrochloric acid (HCl) form salt (NaCl) and water (H₂O); the reaction is reversible.

• A superscript of plus or minus after the atomic symbol indicates an ion. A single plus or minus shows the charge of an ion. If more than one electron has been lost or gained the charge of the ions is shown with a number before the plus or minus sign.

• Chemical reactions do not form or destroy atoms; they just rearrange them into new combinations. In any given chemical equation the number of atoms of each element must be the same on both sides, resulting in a balanced equation.

Inorganic Compounds

Inorganic compounds consist of molecules that do not contain carbon with the exception of a few molecules that are classified as inorganic compounds although they contain carbon, such as carbon dioxide (CO₂) and carbon monoxide (CO).

Acids, Bases, and the pH Scale

Some chemical compounds dissociate in water as ions, carry an electric current, and display an electrical charge. Substances that release hydrogen ions (H⁺) are acids, and those that release hydroxyl ions (OH⁻) are bases. The strength of acids and bases is determined by the hydrogen ion concentration of the water in which they dissociate. The higher the hydrogen ion concentration in the solution the more acidic the solution is. A low hydrogen ion concentration of a solution indicates a basic solution (Box 2.2).

BOX 2.2 Definitions

Acid: A molecule that can release H⁺ into a solution; it is called a “proton donor”

Base: A molecule that can combine with H⁺ and removes it from the solution; it is called a “proton acceptor”

pH: The negative logarithm of the H⁺ concentration of a solution, indicated in pH units on a pH scale that runs from 0 to 14 with pH 7 (pH of pure water at 25° C) as neutral

Buffer: A system of molecules and ions that stabilize the pH of a solution by resisting changes in the H⁺ concentration of the solution

This acidity or alkalinity of a solution is measured by the pH scale (“potential hydrogens”). It is a chemical symbol that ranges from 0 to 14 and is the negative logarithm of the hydrogen ion concentration (Figure 2.9). A solution that has a neutral pH is one in which the concentrations of H⁺ and OH⁻ ions are equal (10⁻⁷ M), and the chemical symbol for this negative logarithm is pH 7. The point of neutrality is standardized as the pH of pure water at 25° C. As the H⁺ concentration increases, the OH⁻ concentration decreases and vice versa. A change one of 1 unit on the pH scale (the negative logarithm scale) represents a 10-fold change in the hydrogen ion concentration. The higher the hydrogen ion concentration of a solution, the lower the number on the pH scale and the more acid the solution; the higher the number on the scale, the higher the hydroxyl ion concentration and the more basic (alkaline) the solution.

FIGURE 2.9 The pH scale.
With increasing hydrogen ion concentrations a solution becomes more acidic and the pH value decreases. When the hydrogen ion concentration decreases the solution becomes more basic and the pH value increases.

Chemical reactions in a living cell respond to slight changes in the pH of their environment. The majority of microbes as well as human cells survive better in a neutral or slightly basic environment. This sensitivity of microbes to changes in pH is used in the control of microbial growth and in food preservation (see Chapter 19, Physical and Chemical Methods of Control). However, as mentioned in Chapter 1, some microorganisms are found to exist successfully in all environments, such as sulfur-oxidizing bacteria, which prefer a very acidic environment, and yeast, which flourishes under slightly acidic conditions.

Buffers

Buffers are chemicals that can absorb hydrogen or hydroxyl ions and therefore resist changes in the pH of a solution. In a living organism buffers are essential to maintain the pH necessary for the survival of its cells. A common endogenous buffer in cells is the bicarbonate (HCO₃⁻) ion. Bicarbonate picks up hydrogen ions (H⁺), forming carbonic acid (H₂CO₃).

HCO3−+H+⇄H2CO3

This reaction is readily reversible so that adjustments in the pH can be made to maintain the required pH in a given environment. All biological fluids, both intracellular and extracellular, are heavily buffered.

Salts

Substances that dissociate in water and normally do not release hydrogen or hydroxyl ions are known as normal salts. Natron is a salt, composed of a mixture of sodium bicarbonate (common baking soda), sodium carbonate (soda ash), a small amount of sodium chloride (table salt), and sodium sulfate. Natron is called impure salt because it has lost its saltiness. Salts that contain a hydroxide ion are basic salts and salts that contain a hydrogen ion are acid salts. Salts are formed when acids and bases react. Solutions of salts in water are called electrolytes and they conduct electricity.

Salts are formed by chemical reactions between a base and an acid or between a metal and an acid. The name of a salt starts with the name of the cation (ammonium, magnesium, etc.) followed by the name of the anion (chloride, sulfate, etc.).

For example:

NH3+HCl→NH4Cl(ammonium chloride)

Mg+H2SO4→MgSO4+H2(magnesium sulfate)

Sometimes salts are referred to less specifically by the name of the cation (e.g., sodium salt, ammonium salt) or by the name of the anion (e.g., chloride, acetate). Common salt-forming cations and anions are shown in Table 2.4.

TABLE 2.4

Common Salt-forming Anions and Cations

Salt-forming Anions*		Salt-forming Cations
Name	Formula	Name	Formula
Acetate (acetic acid)	CH₃COO⁻	Ammonium	NH₄⁺
Carbonate (carbonic acid)	CO₃²⁻	Calcium	Ca²⁺
Chloride (hydrochloric acid)	Cl⁻	Iron	Fe²⁺ and Fe³⁺
Hydroxide	OH⁻	Magnesium	Mg²⁺
Nitrate (nitric acid)	NO₃⁻	Potassium	K⁺
Oxide	O₂⁻	Pyridinium	C₅H₅NH⁺
Phosphate (phosphoric acid)	PO₄³⁻	Quaternary ammonium	NR₄⁺^†
Sulfate (sulfuric acid)	SO₄²⁻	Sodium	Na⁺

^*Parent acid in parentheses.

^†A side chain/group.

Salts are usually solid crystals but can exist as a liquid at room temperature and are called ionic liquids. Different salts can stimulate sensations of all five basic tastes: salty (sodium chloride), sweet (lead diacetate), sour (potassium bitartrate), bitter (magnesium sulfate), and umami (monosodium glutamate). Pure salts are odorless whereas impure salts may smell acidic (acetates) or basic (ammonium salts). Salts can be clear and transparent (sodium chloride), opaque (titanium dioxide), or metallic (iron disulfate). They also exist in different colors (Table 2.5). Most mineral, inorganic pigments and many synthetic organic dyes are salts.

TABLE 2.5

Colors of Salts

Name of Salt	Formula	Color
Sodium chloride	NaCl	Clear, transparent
Titanium dioxide	TiO₂	Opaque, white
Iron disulfide	FeS₂	Metallic
Sodium chromate	Na₂Cr₂O₄·2H₂O	Yellow
Sodium dichromate	Na₂Cr₂O₇·2H₂O	Orange
Mercury sulfide	HgS	Red
Cobalt dichloride hexahydrate	CoCl₂·6H₂O	Mauve
Copper sulfate pentahydrate	CuSO₄·5H₂O	Blue
Ferric hexacyanoferrate	Fe₄[Fe(CN)₆]₃	Blue
Nickel oxide	Ni₂O₃	Green
Magnesium sulfate	MgSO₄	Colorless
Manganese dioxide	MnO₂	Black

Water

Life on earth most likely exists because of the abundance of liquid water. Water is unique because it can exist in three different temperature-dependent states: gas (steam), liquid, and solid (ice). Water is a molecule that consists of hydrogen and oxygen at a ratio of 2 to 1 and is absolutely necessary for all life forms. The bond between the oxygen atom and the hydrogen atom is a polar covalent bond, resulting in the molecule having a slightly positive side and a slightly negative side. The water molecules in water are held together by hydrogen bonding. Water molecules can quickly break down and reform their hydrogen bonds, which results in the property of cohesion. Water’s high level of cohesion, or “stickiness,” allows things to float easily on its surface at the air–water interface, and also causes water to form beads when dispersed. It is the most abundant molecule in the human body and in microorganisms, which contain at least 70% water, and the earth’s surface is 71% covered by water as well.

LIFE APPLICATION

Water Insects

Many insects can skate across the surface of water because of the water’s high surface tension. One of the most common examples is the water strider, an insect that never breaks the surface as it skates on the water of streams, rivers, ponds, and even the ocean. Long-distance water movement is essential to the survival of land plants. On a dry, warm, sunny day a leaf may lose almost 100% of its water in a very short time. Therefore, the water evaporated from the leaves must be replaced by the uptake of water from the soil. This water transport can be explained by the cohesion–tension theory that states that the driving force of this transport is the evaporation (transpiration) of water from the leaf surfaces. The water molecules stick together (cohere) and are pulled up the plant by capillary action (tension) exerted by the evaporation at the leaf surface.

Water’s freezing and boiling points at sea level are the baseline from which temperature is measured. Zero degrees Celsius (0° C) is water’s freezing point and 100° C is its boiling point. The solid form of water (ice) is less dense than the liquid form and for that reason ice floats on water. Whereas most liquids contract as they become colder, water expands until it is solid.

Water is a contributor in most chemical reactions of cells and is essential to break down polymers into monomers by the process of hydrolysis. The amount of water needed for metabolic activities varies among different microorganisms. The availability of water influences microbial growth rates (see Chapter 6, Bacteria and Archaea). Some of the properties of water are shown in Box 2.3.

Water is a solvent and can dissolve many different substances, and is therefore often called the “universal solvent.” The substances dissolved in water or another solvent are solutes, and the combination of a solvent and its solutes is referred to as a solution (Figure 2.10). The solubility of molecules is determined by their molecular structure. Molecules that exhibit local differences in electrical properties, or polar areas, are water soluble and those that do not are insoluble in water. The survival of living cells depends on the appropriate concentration of solutes in a solution. Most organisms do not tolerate environments where the concentration of solutes is much higher than that in their intracellular environment. Depending on the amount of solutes within or outside a cell, the environment can be:

FIGURE 2.10 Solvent, solutes, and solution.
In this diagram water acts as the solvent, sodium and chloride ions are the solutes, and when solutes are dissolved (dissociated) in a solvent, a solution is formed.

• Isotonic: The solute concentration and hence the osmotic pressure within the cell (intracellular) is the same as it is outside of the cell (extracellular). A cell placed in an isotonic solution will not change its cell volume.

• Hypertonic: The solute concentration in the cell is less than in the extracellular environment, which causes a net loss of water from the cell, resulting in cell shrinkage. The cell shape becomes notched or crenated.

• Hypotonic: The solute concentration in the extracellular environment is less than that inside the cell (intracellular), causing the uptake of water into the cell, resulting in the bursting of the cell (Figure 2.11)

FIGURE 2.11 Tonicity.
Cells placed into a hypotonic environment will take up water, swell, and may eventually burst. Cells placed into a hypertonic environment will lose water and shrink. Cells placed in an isotonic environment will remain unchanged.

Because of their polarity, ions attract the polar water molecules that surround the ions, which in turn attract other water molecules to form hydration spheres around each ion (Figure 2.12). Formations of hydration spheres are responsible for the solubility of ions in water. Organic molecules such as glucose, amino acids, and others are water soluble if the covalent bonding pattern permits the formation of hydration spheres around their atoms of oxygen, nitrogen, and phosphorus. These molecules are hydrophilic (water loving) water-soluble compounds. Molecules held together by nonpolar covalent bonds are hydrophobic (water repelling) and insoluble in water because of their inability to form hydration spheres. Parts of drug molecules may be hydrophilic, conferring water solubility properties on them, and vice versa for hydrophobic parts of drug molecules (see Chapter 21, Pharmacology).

FIGURE 2.12 Hydration spheres.
Water molecules form hydration spheres around ions.

Organic Molecules

All organic molecules contain atoms of carbon and hydrogen. Organic molecules have a backbone of chains or rings formed by the carbon and hydrogen atoms, referred to as a hydrocarbon backbone. Carbons commonly form covalent bonds not only with hydrogen, but also with oxygen, nitrogen, sulfur, and phosphorus. It is this covalent bonding of carbons with other carbons that yields the immense number and variety of organic molecules and allows different arrangements of chains or rings (Figure 2.13). The major organic molecules in living organisms are carbohydrates, proteins, lipids, and nucleic acids. Each of these compounds is composed of specific unit molecules or monomers (Table 2.6).

TABLE 2.6

Organic Molecules and Their Monomers

Organic Molecule	Monomer
Carbohydrates	Monosaccharides
Proteins	Amino acids
Lipids	Glycerol and fatty acids
Nucleic acids	Nucleotides

FIGURE 2.13 Carbon backbone of organic molecules.
Carbon atoms can be arranged in **(A)** chains, **(B)** rings, or **(C)** branched chains.

Carbohydrates

Carbohydrates (sugars) include monosaccharides (monomer), disaccharides (two monosaccharides), and polysaccharides (many monosaccharides—polymer), all of which have a characteristic ratio (2 : 1 : 2) of carbon, hydrogen, and oxygen atoms. The name “carbohydrate” (hydrates [water] of carbon) is derived from this ratio. Sugars store carbon as well as large amounts of energy that are extracted during catabolism. Many microorganisms prefer sugars, when they are available, as their source of energy. Carbohydrates are also present in a large variety of cellular structures.

Monosaccharides are simple sugars that contain three to seven carbon atoms and an aldehyde group or a keto group. Monosaccharides represent the unit molecules (monomers) of carbohydrates. Monosaccharides include glucose (C₆H₁₂O₆), fructose, galactose, ribose (C₅H₁₀O₅), and deoxyribose (C₅H₁₀O₄).

Disaccharides (Figure 2.14, A) are compounds formed when two monosaccharides combine with the loss of a water molecule. Disaccharides include the following:

FIGURE 2.14 Carbohydrates.
A, Disaccharide (lactose). B, Polysaccharides (starch and cellulose).

• Sucrose: Composed of glucose and fructose

• Lactose: Composed of glucose and galactose

• Maltose: Composed of two glucose molecules

Polysaccharides (Figure 2.14, B) are formed when many monosaccharides combine to form a larger compound. Starch in plants and glycogen in animals are polysaccharide storage forms of glucose. The most abundant polysaccharide is cellulose, a major component of the cell walls of plants, fungi, and most algae.

Proteins

There are 20 known naturally occurring amino acids (Table 2.7), and they are the monomers of proteins. All amino acids consist of an amino group, a carboxyl group, and a variable side chain designated chemically as R (the R group; Figure 2.15). Two amino acids joined together form a dipeptide; 3 amino acids form a tripeptide; and a chain of 10 or more amino acids form a polypeptide.

TABLE 2.7

Naturally Occurring Amino Acids

Amino Acid	Abbreviation	Amino Acid	Abbreviation
Alanine	Ala	Leucine	Leu
Arginine	Arg	Lysine	Lys
Asparagine	Asn	Methionine	Met
Aspartic acid	Asp	Phenylalanine	Phe
Cysteine	Cys	Proline	Pro
Glutamic acid	Glu	Serine	Ser
Glutamine	Gln	Threonine	Thr
Glycine	Gly	Tryptophan	Trp
Histidine	His	Tyrosine	Tyr
Isoleucine	Ile	Valine	Val

FIGURE 2.15 General structure of amino acids.
Amino acids contain amino and carboxyl groups as well as the variable R group side chain.

Recalling Crick’s central dogma that DNA makes mRNA and mRNA makes protein, the sequence of amino acids in a polypeptide is determined in the codons of messenger ribonucleic acid (mRNA), which translates codons (three-letter genetic words) into a polypeptide chain. The mRNA codons are transcribed from codons in deoxyribonucleic acid (DNA). All living things are dependent on proteins for structure and function (Box 2.4). Proteins can contain up to 10,000 amino acids. The sequence and folding of the amino acid chains determine the shape, which in turn determines the function and specificity of a protein. Different combinations of amino acids yield an infinite variety of polypeptides, which provides the chemical basis for the incredible biological diversity in structure and function among living organisms.

Proteins occur in four different structural arrangements (Figure 2.16):

FIGURE 2.16 Structural arrangements of proteins.
The primary structure consists of a chain of amino acids (AA), the secondary structure can be either an α helix or a β sheet, the tertiary structure has a globular shape, and the quaternary structure contains several different polypeptides forming a functional unit.

• A primary structure is represented by a single chain of amino acids. Examples of primary structure proteins are small hypothalamic hormones such as gonadotropin-releasing hormone (GnRH) and thyrotropin-releasing hormone (TRH).

• A secondary structure is made of polypeptide chains that are either folded into a β sheet (β conformation) or form an α helix. The right-handed helix makes a complete turn in a clockwise direction every 3.6 amino acids. The α helix is held together by hydrogen bonds. β Sheets consist of two or more polypeptide chains lying side by side and are also stabilized by hydrogen bonds. Examples of such helical proteins are keratin, myosin, and collagen; silk represents a β-pleated sheet.

• A tertiary structure has a globular shape because of the additional coiling of secondary structure proteins. This structure is stabilized by the formation of additional hydrogen, ionic, and disulfide bonds. Properties of solubility are determined by hydrophobic nonpolar chains, which are generally positioned on the inside of the protein, and by hydrophilic polar chains, which are positioned on the outside of protein molecules. Examples of tertiary structure proteins are enzymes and some peptide hormones such as insulin.

• A quaternary structure contains several polypeptides that form a functional unit. Such complexes can consist of several copies of the same polypeptide or different polypeptides. An example of a quaternary structure protein is the hemoglobin molecule.

Heat, pH, salts, radiation, and heavy metals can change the shape of a protein, causing protein denaturation. This process results in nonfunctioning protein compounds. Denatured enzymes (proteins) can no longer function as biological catalysts, and their metabolic reactions come to a stop. Denatured antibodies can no longer bind to an antigen and fail to produce the all-important antigen–antibody complex in immune reactions (see Chapter 20, The Immune System). Denatured hormones are no longer able to act on their target cells. And the denaturing of bacterial or viral proteins often will eliminate the microbe (see Chapter 22, Antimicrobial Drugs).

When proteins are combined with inorganic or organic nonprotein compounds they are called conjugated proteins. These compounds are named accordingly as glycoproteins, lipoproteins, nucleoproteins, and phosphoproteins.

Lipids

Lipids are molecules that vary markedly in their chemical structures. With the exception of phospholipids they are hydrophobic and are soluble in organic solvents such as ether, acetone, chloroform, benzene, and alcohols. They consist of hydrocarbon chains and rings, as triglycerides, phospholipids, steroids, cholesterol, prostaglandins, or leukotrienes.

Triglycerides (fats and oils) consist of glycerol and fatty acid chains (neutral fats). At room temperature fats are solid whereas oils are liquid. Structurally a fatty acid has a tail portion, which is a long hydrocarbon chain, and a head portion that consists of a carboxyl group (COOH). The tails of the fatty acids are hydrophobic and the heads are hydrophilic. When the head portion of the fatty acid is attached to a glycerol molecule to form fat, the entire molecule becomes hydrophobic and therefore insoluble in water (Figure 2.17).

FIGURE 2.17 Triglycerides.
The triglyceride molecule consists of glycerol and fatty acid side chains.

Depending on the absence or presence of double bonds between the carbon atoms of the fatty acid chains, the fats are called saturated, monounsaturated, or polyunsaturated (Figure 2.18). In animal fats the carbons of the fatty acid chains are all bonded by single covalent bonds, meaning that all carbons are bonded to the maximal number of hydrogens. Therefore, animal fats are saturated, closely packed together, and solid at room temperature. On the other hand, plant lipids are oils. They have some double bonds between the carbons, causing bends in the shape of the molecule. These oils are unsaturated fats and are liquid at room temperature.

FIGURE 2.18 Saturated and unsaturated fats.
A, Saturated fat with no double bonds within the carbon chain. B, Polyunsaturated fat with several double bonds within the carbon chain.

Fats and oils are essential forms of energy storage. Animals convert excess sugars into fats if the glycogen storage capacity is reached. Although some seeds and fruits store energy as oil, most plants store excess sugars as starch. Fats can store over twice the amount of energy (9.3 kcal/g) than do carbohydrates (3.79 kcal/g).

HEALTHCARE APPLICATION
Selected Lipid Storage Diseases

Disease	Cause	Symptoms	Treatment
Gaucher disease (three common clinical subtypes)	Glucocerebrosidase deficiency	Enlarged spleen and liver, liver malfunction, skeletal disorders, neurological complications, lymph node swelling, distended abdomen, low platelet count, yellow spots in eyes	Enzyme replacement; bone marrow transplant for nonneurological manifestations; splenectomy may be required; blood transfusion for anemia; no effective treatment for severe brain damage
Niemann-Pick disease (four categories)	Accumulation of fat and cholesterol in cells of the liver, spleen, bone marrow, lungs, and sometimes brain; inherited in an autosomal recessive pattern	Enlarged spleen and liver, cherry red spot in the eye, neurological disorders, decline of motor skills	No cure, supportive treatment of symptoms
Fabry disease	α-Galactosidase-A deficiency; buildup of fatty material in the autonomic nervous system, eyes, kidneys, and cardiovascular system	Burning pain in arms and legs, clouding of vision, impaired circulation, heart enlargement, progressive kidney impairment, gastrointestinal difficulties, fever	Drug treatment for pain (phenytoin, carbamazepine); kidney transplant or dialysis; enzyme replacement
Farber’s disease	Ceramidase deficiency; accumulation of fatty material in joints, tissues, and central nervous system	Impaired mental ability; liver, heart, and kidneys may be affected; vomiting, arthritis, swollen lymph nodes, swollen joints, joint contractures	No specific treatment; corticosteroids to relieve pain

Phospholipids consist of glycerol, two fatty acid chains, and a phosphate group at one end (Figure 2.19). These molecules are composed of polar (hydrophilic) heads and nonpolar (hydrophobic) tails. The fatty acid groups are hydrophobic whereas the phosphate group is hydrophilic. Phospholipids therefore have a water-soluble polar head and a nonpolar hydrophobic tail. Cell membranes, for example, are composed of a phospholipid bilayer with the tails facing toward each other and the heads facing outward, making their surfaces hydrophilic and their interior hydrophobic. This arrangement is the basis for their biological barrier properties.

FIGURE 2.19 Phospholipid.
A, Phospholipid bilayer as seen in plasma membranes. B, The phospholipid molecule is composed of a polar, hydrophilic head and a nonpolar, hydrophobic tail.

LIFE APPLICATION

Hand Washing

Fatty acids are the main component of soap, an important emulsifying agent. An emulsifying agent is a substance that is soluble in both oil and water and enables both to mix. The tail portions of fatty acids are soluble in oil and their head portions are soluble in water. The heads emulsify the oil or oily dirt and wash it away.

Hand washing with soap and water is an essential hygiene practice to prevent the spread of disease. Dr. Ignaz Semmelweis first demonstrated in 1847 that routine hand washing can prevent the spread of hand-borne diseases (see Chapter 1, Scope of Microbiology). This was a landmark achievement in healthcare settings and for public health in general. Healthcare specialists now state that hand washing is the single most effective way to prevent the transmission of disease.

Steroids (Figure 2.20) have a carbon skeleton consisting of four fused rings with various functional groups attached. Hundreds of different steroids have been found in plants, animals, and fungi. Categories are as follows:

FIGURE 2.20 Steroids.
Various steroid molecules are shown, all with four fused carbon rings but differing in their functional groups.

• Anabolic steroids interact with androgen receptors to increase muscle and bone mass. Besides the naturally occurring anabolic steroids, synthetic ones exist that are used (sometimes illegally) by athletes in an attempt to enhance their performance.

• Sex steroids are responsible for the secondary sex characteristics of males and females. They include androgens, estrogens, and progesterones.

• Mineralocorticoids help maintain blood volume, electrolyte balance, and osmolarity by controlling the renal secretion of electrolytes.

• Glucocorticoids play a role in many aspects of metabolism and in immune function. They are often prescribed to reduce inflammatory conditions.

• Phytosterols are a group of steroid alcohols that naturally occur in plants such as yeasts and fungi. Plants contain a wide range of phytosterols that are a structural component in their cell membrane and serve the same function as cholesterol does in animal cells. Ergosterol, a phytosterol, is also referred to as provitamin D₂. It is a biological precursor that is converted by ultraviolet irradiation into ergocalciferol, or vitamin D₂.

MEDICAL HIGHLIGHTS

Phytosterol and Ergosterol

Phytosterol is used as a food additive because it lowers cholesterol by reducing cholesterol absorption in the intestines. It may act in cancer prevention as well. Small amounts of phytosterols naturally occur in vegetable oils, especially in soybean oil.

Ergosterol inhibitors are used as antifungal compounds (see Chapter 22, Antimicrobial Drugs) because they disrupt the membrane functions in fungal membranes and affect the membranes of higher eukaryotic cells.

Cholesterol (C₂₇H₄₅OH) is a sterol, a combination of a steroid and an alcohol. Cholesterol plays a major role in many biochemical processes such as in the metabolism of fat-soluble vitamins (vitamins A, D, E, and K) and acts as a precursor for vitamin D. Cholesterol also is a precursor for the synthesis of steroid hormones and is an important component of cell membranes.

Prostaglandins consist of a fatty acid and a cyclic hydrocarbon group. A variety of prostaglandins have been identified and all of them play different roles in a variety of tissues. Prostaglandins are called local hormones because they act as chemical messengers but do not move to other sites. They play a major role in a variety of body and cell regulatory processes, including involvement in defense mechanisms such as blood clotting and inflammation.

MEDICAL HIGHLIGHTS

Prostaglandins

During tissue damage white blood cells are transported to the site to minimize tissue destruction and infection. This invasion of white blood cells results in the release of prostaglandin, which activates the inflammatory response, causing pain and fever (see Chapter 20, The Immune System). One type of prostaglandin, thromboxane, stimulates the platelets to form a platelet clot during the blood-clotting process. Other prostaglandins stimulate uterine contraction during induction of labor. In allergic reactions prostaglandins are responsible for vasodilatation, increased vascular permeability, bronchoconstriction, and increased sensitivity to pain. Antiinflammatory drugs such as aspirin and corticosteroids act by preventing the actions of prostaglandins (see Chapter 21, Pharmacology, and Chapter 22, Antimicrobial Drugs).

Nucleic Acids

The monomers of nucleic acids are nucleotides, containing the elements carbon, hydrogen, oxygen, nitrogen, and phosphorus. Main functions of nucleic acids include the following: storage of genetic information (DNA), directing protein synthesis (RNA), and energy transfers (ATP and NAD). Nucleotides are composed of three units: a pentose sugar, a phosphate, and nitrogen base (Figure 2.21). The nucleotide structure can be broken down into two main functional parts: a sugar–phosphate backbone and the base. Individual nucleic acids are named according to the sugar they contain, a ribose (RNA) or a deoxyribose (DNA). Five possible nitrogen bases are subdivided into two main groups: purines with adenine (A) and guanine (G) that have a distinctive two-ring structure, and pyrimidines with cytosine (C), thymine (T), and uracil (U) with a single-ringed structure (see Figure 2.21). The bases in a nucleic acid polymer can form hydrogen bonds with the neighboring bases by a process called complementary base pairing. In DNA molecules, adenine always pairs with thymine and guanine with cytosine. In RNA molecules, thymine is replaced with uracil.

FIGURE 2.21 DNA molecule.
Deoxyribonucleic acid (DNA) is composed of nucleotides, each of which consists of a phosphate, a pentose sugar (deoxyribose), and a nitrogen base (shaded area). The nitrogen bases of nucleic acids are either **(A)** purines or **(B)** pyrimidines.

Deoxyribonucleic acid (DNA) is a nucleic acid with a double helix structure containing the sugar deoxyribose and 10 bases per turn (see Figure 2.21). DNA polymers can be thousands of bases long. DNA contains the genetic code and therefore serves for information storage. DNA is responsible for inherited characteristics, growth, and cell reproduction. It is present in both prokaryotic and eukaryotic cells as well as in a group of viruses. DNA contains the information necessary for protein synthesis. The language of DNA, the genetic code, consists of four letters that represent the nitrogen bases, C, G, A, and T, used in three-letter “words” called codons to indicate the 20 naturally occurring amino acids. The combination of codons can create an infinite variety of “sentences.” Each codon represents a specific amino acid and the combination of amino acids results in different polypeptides (see Table 3.6 in Chapter 3, Cell Structure and Function).

Chromosomes are the microscopic structures that carry DNA within the nucleus of cells. One chromosome represents a single molecule of DNA. In bacteria chromosomes with two types of DNA are present: the chromosomal DNA located in the nucleoid area and plasmids, which are simple circles of DNA floating feely in the organism. Plasmids are capable of autonomous replication and therefore can replicate independently of the chromosomal DNA. In eukaryotes the chromosomes are highly complex structures and DNA molecules are linear rather than circular.

Ribonucleic acid (RNA) is similar to DNA but is a single-stranded molecule, its sugar is ribose, and uracil replaces thymine. RNA is specialized for the synthesis of proteins. Three different types of RNA are necessary for the process of protein synthesis: ribosomal RNA (rRNA), messenger RNA (mRNA), and transfer RNA (tRNA). Ribosomes composed of rRNA are made in the nucleus and transported into the cytoplasm, where they attach to rough endoplasmic reticulum (rER) or remain free in the cytoplasm as polyribosomes. Both polyribosomes and ribosomes on the rER (see Chapter 3, Cell Structure and Function) are sites of protein synthesis. Messenger RNA contains genetic information that encodes the sequence of amino acids in proteins. Base triplets, referred to as codons, direct the amino acid sequence in a polypeptide chain. The third type of RNA, tRNA, contains a triplet of nitrogen bases called the anticodon. Anticodons contain complementary bases to the codons on mRNA, and are necessary for the synthesis of polypeptides (see Protein Synthesis in Chapter 3).

Adenosine triphosphate (ATP) is the energy molecule of cells. When energy is released during catabolism it is captured in the high-energy bonds of ATP (Figure 2-22). In transferring energy to other molecules, ATP loses one or two of its phosphate groups, resulting in adenosine diphosphate (ADP) or adenosine monophosphate (AMP). This is an exergonic reaction. Both ADP and AMP can be converted back to ATP by photosynthesis or through chemical energy during anabolism. Photosynthetic microorganisms use sunlight as an energy source during their anabolism. Microorganisms that need nutrient molecules for ATP production are called chemotrophs and the ones utilizing sunlight for energy are called phototrophs (see Chapter 6, Bacteria and Archaea).