The Chemistry of Microbiology

         OF MICROBIOLOGY THE CheMISTRY Learning or reviewing some basic concepts of chemistry will enable you to understand more fully the variety of interactions between microorganisms and their environment—which includes you. If you plan a career in health care, you will find microbial chemistry involved in the diagnosis of disease, the response of the immune system, the growth and identification of pathogenic microorganisms in the laboratory, and the function and selection of antimicrobial drugs. Understanding the fundamentals of chemistry will even help you preserve your own health. In this chapter, we study atoms, which are the basic units of chemistry, and we consider how atoms react with one another to form chemical bonds and molecules. Then we examine the three major categories of chemical reactions. The chapter concludes with a look at the molecules of greatest importance to life: water, acids, bases, lipids, carbohydrates, proteins, nucleic acids, and ATP. Electron shells Nucleus Proton (p+) Neutron (n0) Electron (e–) ▲ FIGURE 2.1 An example of a Bohr model of atomic structure. This drawing is not to scale; for the electrons to be shown in scale with the greatly magnified nucleus, the electrons would have to occupy orbits located many miles from the nucleus. Put another way, the volume of an entire atom is about 100 trillion times the volume of its nucleus. Atoms Matter is defined as anything that takes up space and has mass.1 The smallest chemical units of matter are atoms. Atoms are extremely small, and only the very largest of them can be seen using the most powerful microscopes. Therefore, scientists have developed various models to conceptualize and illustrate the structure of atoms. Atomic Structure LE A R N I N G O U TC O M E S 2.2 Draw and label an atom, showing the parts of the nucleus and orbiting electrons. 2.3 Define element, atomic number, atomic mass, and dalton, and explain how these terms relate to one another. In 1913, the Danish physicist Niels H. D. Bohr (1885–1962) proposed a simple model in which negatively charged subatomic particles called electrons orbit a centrally located nucleus like planets in a miniature solar system (FIGURE 2.1). A nucleus is composed of both uncharged neutrons and positively charged protons. (The only exception is the nucleus of a normal hydrogen atom, which is composed of only a single proton and no neutrons.) An element is matter that is composed of a single type of atom. For example, gold is an element because it consists of only gold atoms. In contrast, the ink in your pen is not an element because it is composed of many different kinds of atoms. Elements differ from one another in their atomic number, which is the number of protons in their nuclei. For example, the atomic numbers of hydrogen, carbon, and oxygen are 1, 6, Isotopes LE A R N I N G O UTC O M E 2.4 List at least four ways that radioactive isotopes are useful. Every atom of an element has the same number of protons, but atoms of a given element can differ in the number of neutrons in their nuclei. Atoms that differ in this way are called isotopes. For example, there are three naturally occurring isotopes of carbon, each having six protons and six electrons (FIGURE 2.2). Over 95% of carbon atoms also have six neutrons. Because these carbon atoms have six protons and six neutrons, the atomic mass of this isotope is about 12 daltons, and it is known as carbon-12, symbolized as 12C. Atoms of carbon-13 (13C) have seven neutrons per nucleus, and 14C atoms each have eight neutrons. ­ relate to one another. and 8, respectively, because all hydrogen nuclei contain a single proton, all carbon nuclei have six protons, and all oxygen nuclei have eight protons. The atomic mass of an atom (sometimes called its atomic weight) is the sum of the masses of its protons, neutrons, and electrons. Protons and neutrons each have a mass of approximately 1 atomic mass unit,2 which is also called a dalton.3 An electron is much less massive, with a mass of about 0.00054 dalton. Electrons are often ignored in discussions of atomic mass because their contribution to the overall mass is negligible. Therefore, the sum of the number of protons and neutrons approximates the atomic mass of an atom. There are 93 naturally occurring elements known;4 however, organisms typically utilize only about 20 elements, each of which has its own symbol that is derived from its English or Latin name (TABLE 2.1). ­ 2.1 Define matter and atom, and explain how these terms ­ LE A R N I N G O U TC O M E atomic mass unit (dalton) is 1/597,728,630,000,000,000,000,000, or 1.673 * 10 -24, grams. for John Dalton, the British chemist who helped develop atomic theory around 1800. 4For many years, scientists thought that there were only 92 naturally occurring elements, but natural plutonium was discovered in Africa in 1997. 2An 1Mass and weight are sometimes confused. Mass is the quantity of material in something, whereas weight is the effect of gravity on mass. Even though an astronaut is weightless in space, his mass is the same in space as on Earth. M02_BAUM7206_05_SE_C02.indd 28 3Named 12/1/16 3:25 PM       A toms TABLE 2.1 Common Elements of Life Symbol ydrogen H H Element Atomic Number Atomic Massa (daltons) 1 1 Biological Significance Component of organic molecules and water; h + released by acids B 5 11 Carbon C 6 12 Backbone of organic molecules Nitrogen N 7 14 Component of amino acids, proteins, and nucleic acids Oxygen O 8 16 Component of many organic molecules and water; Oh - released by bases; necessary for aerobic metabolism Sodium (Natrium) Na 11 23 Principal cation outside cells Magnesium Mg 12 24 Component of many energy-transferring enzymes Silicon Si 14 28 Component of cell wall of diatoms Phosphorus P 15 31 Component of nucleic acids and TP Sulfur S 16 32 Component of proteins Chlorine Cl 17 35 Principal anion outside cells Potassium (Kalium) K 19 39 Principal cation inside cells; essential for nerve impulses Calcium Ca 20 40 Utilized in many intercellular signaling processes; essential for muscular contraction Manganese Mn 25 54 Component of some enzymes; acts as intracellular antioxidant; used in photosynthesis Iron (Ferrum) Fe 26 56 Component of energy-transferring proteins; transports oxygen in the blood of many animals Cobalt Co 27 59 Component of vitamin B 12 Copper (Cuprum) Cu 29 64 Component of some enzymes; used in photosynthesis Zinc Zn 30 65 Component of some enzymes Molybdenum Mo 42 96 Component of some enzymes I 53 127 Component of many brown and red algae Iodine aRounded ssential for plant growth A E Boron to nearest whole number. Unlike carbon-12 and carbon-13, the nucleus of carbon-14 is unstable because of the ratio of its protons and neutrons. Unstable atomic nuclei release energy and subatomic particles such as neutrons, protons, and electrons in a process called radioactive decay. Atoms that undergo radioactive decay are radioactive isotopes. Radioactive decay and radioactive isotopes play important roles in microbiological research, medical diagnosis, the treatment of disease, and the complete destruction of Additional neutron Proton Neutron (a) Carbon-12 6 Protons 6 Neutrons (b) Carbon-13 6 Protons 7 Neutrons (c) Carbon-14 6 Protons 8 Neutrons E ­ ▲ FIGURE 2.2 Nuclei of the three naturally occurring isotopes of carbon. ach isotope also has six electrons, which are not shown. What are the atomic number and atomic mass of each of these isotopes? Figure 2.2 The atomic number of all three is 6; their atomic masses are 12, 13, and 14, respectively. M02_BAUM7206_05_SE_C02.indd 29 contaminating microbes (sterilization) of medical equipment and chemicals. Electron Configurations Although the nuclei of atoms determine their identities, it is electrons that determine an atom’s chemical behavior. Nuclei of different atoms almost never come close enough together to interact.5 Typically, only the electrons of atoms interact. Thus, because all of the isotopes of carbon (for example) have the same number of electrons, all these isotopes behave the same way in chemical reactions, even though their nuclei are different. Scientists know that electrons do not really orbit the nucleus in a two-dimensional circle, as indicated by a Bohr model; instead, they speed around a nucleus 100 quadrillion times per second in three-dimensional electron shells or clouds that assume unique shapes dependent on the energy of the electrons (FIGURE 2.3a). More accurately put, an electron shell depicts the probable locations of electrons at a given time; nevertheless, it is simpler and more convenient to draw electron shells as circles (FIGURE 2.3b). 5Except during nuclear reactions, such as occur in nuclear power plants. 12/1/16 3:25 PM e 2    The Chemistry of Microbiology r Each electron shell can hold only a certain maximum number of electrons. For example, the first shell (the one nearest the nucleus) can accommodate a maximum of two electrons, and the second shell can hold no more than eight electrons. Atoms of hydrogen and helium have one and two electrons, respectively; thus, these two elements have only a single electron shell. A lithium atom, which has three electrons, has two shells. Atoms with more than 10 electrons require more shells. The third shell holds up to eight electrons when it is the outermost shell, though its capacity increases to 18 when the fourth shell contains two electrons. Heavier atoms have even more shells, but these atoms do not play significant roles in the processes of life. Electrons in the outermost shell of atoms are called valence electrons. FIGURE 2.4 depicts the electron configurations of atoms of some elements important to microbial life. Notice that except for helium, atoms of all elements in a given column of the periodic table of elements have the same number of valence electrons. Helium is placed in the far right-hand column with the other inert gases because its outer shell is full, though it has two rather than eight valence electrons. Valence electrons are critical for interactions between atoms. Next, we consider these interactions, which are called chemical bonds. ­ pt       Cha + + First shell = Second shell (a) Electron shells of neon: three-dimensional view + First shell = Second shell (b) Electron shells of neon: two-dimensional view ▲ FIGURE 2.3 Electron configurations. (a) Three-dimensional model of the electron shells of neon. In this model, the first shell is a small sphere, whereas the second shell consists of a larger sphere plus three pairs of ellipses that extend from the nucleus at right angles. Larger shells (not shown) are even more complex. (b) Two-dimensional model (Bohr diagram) of the electron shells of neon. How many shells does a sodium atom need to hold its 11 electrons? TELL ME WHY Electrons zip around the nucleus at about 5 million miles per hour. Why don’t they fly off? Figure 2.3 Three shells: two electrons in the first shell, eight in the second, and a single electron in the third. H He Li Be B Na Mg Hydrogen (H) Periodic table of elements C N O F Ne Al Si P S Cl Ar Helium (He) K Ca 2p 1p Lithium (Li) Beryllium (Be) Boron (B) Carbon (C) Nitrogen (N) Oxygen (O) Fluorine (F) Neon (Ne) 3p 4p 5p 6p 7p 8p 9p 10p Sodium (Na) Magnesium (Mg) Aluminum (Al) Silicon (Si) Phosphorus (P) Sulfur (S) Chlorine (Cl) Argon (Ar) 11p 12p 13p 14p 15p 16p 17p 18p Potassium (K) Calcium (Ca) ▲ FIGURE 2.4 Bohr diagrams of the first 20 elements and their places 20p within the chart known as the periodic table of the elements. Note that the number of valence electrons increases from left to right in each row and that every element in a column has the same number of valence electrons (with the exception of helium). eavier atoms have been omitted because most heavy elements are less important to living organisms. Neutrons are not shown because they have little effect on chemistry. H 19p M02_BAUM7206_05_SE_C02.indd 30 12/1/16 3:25 PM       Chemical Bonds LE A R N I N G O U TC O M E S 2.5 Describe the configuration of electrons in a stable atom, and explain how valence electrons form chemical bonds. 2.6 Contrast molecules and compounds. Outer electron shells are stable when they contain eight electrons (except for the first electron shell, which is stable with only two electrons, because that is its maximum number). When atoms’ outer shells are not filled with eight electrons, they either have room for more electrons or have “extra” electrons, depending on whether it is easier for them to gain electrons or lose electrons. For example, an oxygen atom, with six electrons in its outer shell, has two “unfilled spaces” (see Figure 2.4) because it requires less energy for the oxygen atom to gain two electrons than to lose six electrons. A calcium atom, by contrast, has two “extra” electrons in its outer (fourth) shell because it requires less energy to lose these two electrons than to gain six new ones. When a calcium atom loses two electrons, its third shell, which is then its outer shell, is full and stable with eight electrons. As previously noted, an atom’s outermost electrons are called valence electrons, and thus the outermost shell of an atom is the valence shell. An atom’s valence,6 defined as its combining capacity, is considered to be positive if its valence shell has extra electrons to give up and to be negative if its valence shell has spaces to fill. Thus, a calcium atom, with two electrons in its valence shell, has a valence of +2, whereas an oxygen atom, with two spaces to fill in its valence shell, has a valence of -2. Atoms combine with one another by either sharing or transferring valence electrons in such a way as to fill their valence shells. Such interactions between atoms are called chemical bonds. Two or more atoms held together by chemical bonds form a molecule. A molecule that contains atoms of more than one element is a compound. Two hydrogen atoms bonded together form a hydrogen molecule, which is not a compound because only one element is involved. However, two hydrogen atoms bonded to an oxygen atom form a molecule of water (H2O), which is a compound. In this section, we discuss the three principal types of chemical bonds: nonpolar covalent bonds, polar covalent bonds, and ionic bonds. We also consider hydrogen bonds, which are weak forces that act with polar covalent bonds to give certain large chemicals their characteristic three-dimensional shapes. Nonpolar Covalent Bonds LE A R N I N G O U TC O M E S 2.7 Contrast nonpolar covalent, polar covalent, and ionic bonds. 2.8 Define organic compound. A covalent7 bond is the sharing of a pair of electrons between two atoms. Consider, for example, what happens when two 6From Latin valentia, meaning “strength.” 7 From Latin co, meaning “with” or “together,” and valentia, meaning “strength.” M02_BAUM7206_05_SE_C02.indd 31 hydrogen atoms approach one another. Each hydrogen atom consists of a single proton orbited by a single electron. Because the valence shell of each hydrogen atom requires two electrons in order to be full, each atom shares its single electron with the other, forming a hydrogen molecule in which both atoms have full shells (FIGURE 2.5a). Similarly, two oxygen atoms can share electrons, but they must share two pairs of electrons for their valence shells to be full (FIGURE 2.5b). Because two pairs of electrons are involved, oxygen atoms form two covalent bonds, or a double covalent bond, with one another. The attraction of an atom for electrons is called its electronegativity. The more electronegative an atom, the greater the pull its nucleus exerts on electrons. Note in FIGURE 2.6, which displays the electronegativities of atoms of several elements, that electronegativities tend to increase from left to right in the chart. The reason is that elements toward the right of the chart have more protons and thus exert a greater pull on electrons. Electronegativities of elements decrease from top to bottom in the chart because the distance between the nucleus and the valence shell increases as elements get larger. Atoms with equal or nearly equal electronegativities, such as two hydrogen atoms or a hydrogen and a carbon, share electrons equally or nearly equally. In chemistry and physics, “poles” are opposed forces, such as north and south magnetic poles or positive and negative terminals of a battery. In the case of atoms with similar electronegativities, the shared electrons tend to spend an equal amount of time around each nucleus of the pair, and no poles exist; therefore, the bond between them is a nonpolar covalent bond. (The covalent bonds illustrated in Figure 2.5a–c are nonpolar; formaldehyde, shown in Figure 2.5d, is polar.) A hydrogen molecule can be symbolized a number of ways: ­ Chemical Bonds H ¬ H H : H H2 In the first symbol, the dash represents the chemical bond between the atoms. In the second symbol, the dots represent the electron pair of the covalent bond. These two symbols are known as structural formulas. In the third symbol, known as a molecular formula, the subscript “2” indicates the number of hydrogen atoms that are bonded, not the number of shared electrons. Each of these symbols indicates the same thing—two hydrogen atoms are sharing a pair of electrons. Many atoms need more than one electron to fill their valence shell. For instance, a carbon atom has four valence electrons and needs to gain four more if it is to have eight in its valence shell. (FIGURE 2.5c) illustrates a carbon atom sharing electrons with four hydrogen atoms. As before, a line in the structural formula represents a covalent bond formed from the sharing of two electrons. Two covalent bonds are formed between an oxygen atom and a carbon atom in formaldehyde (FIGURE 2.5d). This fact is represented by a double line, which indicates that the carbon atom shares four electrons with the oxygen atom. Carbon atoms are critical to life. Because a carbon atom has four electrons in its valence shell, it has equal tendencies to either lose four electrons or gain four electrons. Either event produces a full outer shell. The result is that carbon atoms 12/1/16 3:25 PM (a) e 2    The Chemistry of Microbiology r pt       Cha + 1p Hydrogen (H) (b) 1p Hydrogen (H) + 8p 1p Oxygen (O) 1p Hydrogen molecule (H2) 8p 8p Oxygen (O) 4 + 1p Oxygen molecule (O2) 6p H H or O O 8p 1p (c) or 1p 6p 1p H 1p H Hydrogen (H) Methane (CH4) Carbon (C) or C H H 1p (d) + 2 6p 1p + 8p 6p 8p H 1p C Carbon (C) Hydrogen (H) Formaldehyde (CH2O) Oxygen (O) or O H E H ▲ FIGURE 2.5 Four molecules formed by covalent bonds. (a) ydrogen. ach hydrogen A atom needs another electron to have a full valence shell. The two atoms share their electrons, forming a covalent bond. (b) Oxygen. Oxygen atoms have six electrons in their valence shells; thus, they need two electrons each. When they share with each other, two covalent bonds are formed. Note that the valence electrons of oxygen atoms are in the second shell. (c) methane molecule, which has four single covalent bonds. (d) Formaldehyde. The carbon atom forms a double bond with the oxygen atom and single bonds with two hydrogen atoms. Which of these molecules are also compounds? Why? Figure 2.5 Methane and formaldehyde molecules are also compounds because they are composed of more than one element. tend to share electrons and form four covalent bonds with one another and with many other types of atoms. Each carbon atom, in effect, acts as a four-way intersection where different components of a molecule can attach. One result of this feature is that carbon atoms can form very large chains that constitute the “backbone” of many biologically important molecules. Carbon chains can be branched or unbranched, and some even close back on themselves to form rings. Compounds that contain carbon and hydrogen atoms are called organic compounds. Among the many biologically important organic compounds are proteins and carbohydrates, which are discussed later in the chapter. M02_BAUM7206_05_SE_C02.indd 32 Polar Covalent Bonds LE A R N I N G O UTC O M E 2.9 Explain the relationship between electronegativity and the polarity of a covalent bond. If two covalently bound atoms have significantly different electronegativities, their electrons will not be shared equally. Instead, the electron pair will spend more time orbiting the nucleus of the atom with greater electronegativity. This type of bond, in which there is unequal sharing of electrons, is a polar 12/1/16 3:25 PM       Chemical Bonds I II III IV V VI However, it is more convenient to simply draw molecules as though all the atoms were in one plane; for example, H ¬ O ¬ H. Inert gases VII H 2.1 He 0.0 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne 0.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar 0.0 K 0.8 Ca 1.0 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Kr 0.0 Ionic Bonds LE A R N I N G O UTC O M E S 2.10 Define ionization using the terms cation and anion. 2.11 Contrast the formation of an ionic bond with that of a covalent bond. ▲ FIGURE 2.6 Electronegativity values of selected ­ elements. The values are expressed according to the Pauling scale, named for the Nobel Prize–winning chemist Linus Pauling, who based the scale on bond energies. Pauling chose to compare the electronegativity of each element to that of fluorine, to which he assigned a value of 4.0. covalent bond. An example of a molecule with polar covalent bonds is water (FIGURE 2.7a). Because oxygen is more electronegative than hydrogen, the electrons spend more time near the oxygen nucleus than near the hydrogen nuclei, and thus the oxygen atom acquires a transient (partial) negative charge (symbolized as d - ). Each of the hydrogen nuclei has a corresponding transient positive charge (d + ). The covalent bond between an oxygen atom and a hydrogen atom is called polar because the atoms have opposite partial electrical charges. Polar covalent bonds can form between many different elements. Generally, molecules with polar covalent bonds are water soluble, and nonpolar molecules are not. The most important polar covalent bonds for life are those that involve hydrogen because they allow hydrogen bonding, which we discuss shortly. Both nonpolar and polar covalent bonds form angles between atoms such that the distances between electron orbits are maximized. The bond angle for water is shown in FIGURE 2.7b. Consider what happens when two atoms with vastly different electronegativities—for example, sodium, with one electron in its valence shell and an electronegativity of 0.9, and chlorine, with seven electrons in its valence shell and an electronegativity of 3.0—come together (FIGURE 2.8). Chlorine has such a higher electronegativity that it very strongly attracts sodium’s single valence electron, and the result is that the sodium loses that electron to chlorine 1 . Now that the chlorine atom has one more electron than it has protons, it has a full negative charge, and the sodium atom, which has lost an electron, now has a full positive charge 2 . An atom or group of atoms that has either a full negative charge or a full positive charge is called an ion. Positively charged ions are called cations, whereas negatively charged ions are called anions. Because of their opposite charges, cations and anions attract each other and form what is termed an ionic bond 3 . They form 1 Electron lost Na Cl Sodium atom Chlorine atom 2 Attraction of opposite charges + Na f– – Sodium ion (Na+) Cl Chlorine ion (Cl–) 8p f+ (a) 1p 1p f+ H f+ Of – 1055 H 3 Formation of an ionic bond f+ (b) ▲ FIGURE 2.7 Polar covalent bonding in a water molecule. Na +– Cl A A (a) Bohr model of a water molecule, which has two polar covalent bonds. When the electronegativities of two atoms are significantly different, the shared electrons of covalent bonds spend more time around the more electronegative atom, giving it a transient negative charge (d - ). Its partner has a transient positive charge (d + ). (b) The bond angle in a water molecule. toms maximize the distances between electron orbitals in polar and nonpolar covalent bonds. M02_BAUM7206_05_SE_C02.indd 33 Sodium chloride (NaCl) ▲ FIGURE 2.8 The interaction of sodium and chlorine to form an ionic bond. 12/1/16 3:25 PM e 2    The Chemistry of Microbiology r pt       Cha In nature, chemical bonds range from nonpolar bonds to polar bonds to ionic bonds. The important thing to remember is that electrons are shared between atoms in covalent bonds and transferred from one atom to another in ionic bonds. Hydrogen Bonds LE A R N I N G O UTC O M E 2.12 Describe hydrogen bonds, and discuss their importance in living organisms. As we have seen, hydrogen atoms bind to oxygen atoms by means of polar covalent bonds, resulting in transient positive charges on the hydrogen atoms. Hydrogen atoms form polar covalent bonds with atoms of other elements as well. The electrical attraction between a partially charged hydrogen atom and a full or partial negative charge on either a different region of the same molecule or another molecule is called a hydrogen bond (FIGURE 2.10). Hydrogen bonds can be likened to weak ionic bonds in that they arise from the attraction of positive and negative charges. Notice also that although they are a consequence of polar covalent bonds between hydrogen atoms and other, more electronegative atoms, hydrogen bonds themselves are not covalent bonds—they do not involve the sharing of electrons. As we have seen, covalent bonds are essential for life because they strongly link atoms together to form molecules. Hydrogen bonds, though weaker than covalent bonds, are also essential. The cumulative effect of numerous hydrogen bonds is to stabilize the three-dimensional shapes of large molecules. For example, the familiar double-helix shape of DNA is due in part to the stabilizing effects of thousands of hydrogen bonds holding the molecule together. Hydrogen bonds play a role in the recognition of target cells by pathogens and also play a role in maintaining the exact shape of enzymes, antibodies, and intercellular chemical messengers, which is critical for their correct function. Further, because hydrogen bonds are weak, ­ crystalline compounds composed of metallic and nonmetallic ions. These compounds are known as salts. Examples include sodium chloride (NaCl), also known as table salt, and potassium chloride (KCl, sodium-free table salt). Ionic bonds differ from covalent bonds in that ions do not share electrons. Instead, the bond is formed from the attraction of opposite electrical charges. The polar bonds of water molecules interfere with the ionic bonds of salts, causing dissociation (also called ionization) (FIGURE 2.9). This occurs as the partial negative charge on the oxygen atom of water attracts cations, and the partial positive charge on hydrogen atoms attracts anions. The presence of polar bonds interferes with the attraction between the cation and anion. When cations and anions dissociate from one another and become surrounded by water molecules (are hydrated), they are called electrolytes because they can conduct electricity through the solution. Electrolytes are critical for life because they stabilize a variety of compounds, act as electron carriers, and allow electrical gradients to exist within cells. (We examine these functions of electrolytes in later chapters.) Na+ Cl– Na+ Na+ When water surrounds ions in salt crystal, individual Na+ and Cl– ions break away. H H O f– f– f – Na+ f– HO H f+ f– f+ – + f f Cl– f+ Hydrogen bond H f+ f H C + H Hydrated chlorine ion H ▲ FIGURE 2.9 Dissociation of NaCl in water. When water surrounds the ions in a NaCl crystal, the partial charges on water molecules are attracted to charged ions, and the water molecules hydrate the ions by surrounding them. The partial negative charges (d - ) on oxygen atoms are attracted to cations (in this case, the sodium ions), and the partial positive charges (d + ) on hydrogen atoms are attracted to anions (the chlorine ions). Because the ions no longer attract one another, the salt crystal dissolves. ydrated ions are called electrolytes. M02_BAUM7206_05_SE_C02.indd 34 f– C C N N Hydrated sodium ion N H C f+ H f+ H f– f+ O H Cytosine f– O N C C N C C C N H N H N H Guanine ▲ FIGURE 2.10 Hydrogen bonds. The transient positive charge (symbolized d + ) on a hydrogen atom is attracted to a transient negative charge (d - ) on another atom. Such attraction is a hydrogen bond. ydrogen bonds can hold together portions of the same molecule or hold two different molecules together. In this case, three hydrogen bonds are holding molecules of cytosine and guanine together. H Na+ 12/1/16 3:25 PM       Chemical Reactions TABLE 2.2 Characteristics of Chemical Bonds Description Relative Strength Nonpolar covalent bond Pair of electrons is nearly equally shared between two atoms Strong H Ionic bond ydrogen bond E Polar covalent bond E Type of Bond lectrons spend more time around the more electronegative of two atoms Strong lectrons are stripped from a cation by an anion Weaker than covalent in aqueous environments Partial positive charges on hydrogen atoms are attracted to full and partial negative charges on other molecules or other regions of the same molecule they can be overcome when necessary. For example, the two complementary halves of a DNA molecule are held together primarily by hydrogen bonds, and they can be separated for DNA replication and other processes (see Chapter 7, Figure 7.6). TABLE 2.2 summarizes some characteristics of chemical bonds. Weaker than ionic Synthesis Reactions LE A R N I N G O UTC O M E S 2.14 Give an example of a synthesis reaction that involves the formation of a water molecule. 2.15 Contrast endothermic and exothermic chemical reactions. TELL ME WHY Chlorine and potassium atoms form ionic bonds, carbon atoms form nonpolar covalent bonds with nitrogen atoms, and oxygen forms polar covalent bonds with phosphorus. Explain why these bonds are the types they are. Chemical Reactions LE A R N I N G O U TC O M E 2.13 Describe three general types of chemical reactions found in living things. You are already familiar with many consequences of chemical reactions: you add yeast to bread dough, and it rises; enzymes in your laundry detergent remove grass stains; and gasoline burned in your car releases energy to speed you on your way. What exactly is happening in these reactions? What is the precise definition of chemical reaction? We have discussed how bonds are formed via the sharing of electrons or the attraction of positive and negative charges. Scientists define chemical reactions as the making or breaking of such chemical bonds. All chemical reactions begin with reactants—the atoms, ions, or molecules that exist at the beginning of a reaction. Similarly, all chemical reactions result in products— the atoms, ions, or molecules left after the reaction is complete. Biochemistry involves the chemical reactions of living things. Reactants and products may have very different physical and chemical characteristics. For example, hydrogen and oxygen are gases and have very different properties from water, which is composed of hydrogen and oxygen atoms. However, the numbers and types of atoms never change in a chemical reaction; atoms are neither destroyed nor created, only rearranged. Now let’s turn our attention to three general categories of biochemical reactions (reactions that occur in organisms): synthesis, decomposition, and exchange reactions. M02_BAUM7206_05_SE_C02.indd 35 Synthesis reactions involve the formation of larger, more complex molecules. Synthesis reactions can be expressed symbolically as Reactant + Reactant S Product(s) The arrow indicates the direction of the reaction and the formation of new chemical bonds. For example, algae make their own glucose (sugar) using the following reaction: 6 H2O + 6 CO2 S C6H12O6 + 6 O2 The reaction is read, “Six molecules of water plus six molecules of carbon dioxide yield one molecule of glucose and six molecules of oxygen.” Notice that the total number and kind of atoms are the same on both sides of the reaction. A common synthesis reaction in biochemistry is a dehydration synthesis, in which two smaller molecules are joined together by a covalent bond and a water molecule is also formed (FIGURE 2.11a). The word dehydration in the name of this type of reaction refers to the fact that one of the products is a water molecule formed when a hydrogen ion (H + ) from one reactant combines with a hydroxyl ion (OH - ) from another reactant. Synthesis reactions require energy to break bonds in the reactants and to form new bonds to make products. Reactions that require energy are said to be endothermic 8 reactions because they trap energy within new molecular bonds. An energy supply for fueling synthesis reactions is one common requirement of all living things (Chapter 6). Taken together, all of the synthesis reactions in an organism are called anabolism. Decomposition Reactions LE A R N I N G O UTC O M E 2.16 Give an example of a decomposition reaction that involves breaking the bonds of a water molecule. 8From Greek endon, meaning “within,” and thermos, meaning “heat” (energy). 12/1/16 3:25 PM e 2    The Chemistry of Microbiology r pt       Cha H OH Small molecule H OH + Small molecule H Energy OH + + Larger molecule H2O H2O (a) Dehydration synthesis H OH + Large molecule + H2O H2O H Energy OH Small molecule H + OH Small molecule (b) Hydrolysis ▲ FIGURE 2.11 Two types of chemical reactions in living things. (a) Dehydration synthesis. In this energyH H H requiring reaction, a hydroxyl ion (Oh - ) removed from one reactant and a hydrogen ion (h + ) removed from another reactant combine to form hydrogen hydroxide ( O ), which is water. (b) ydrolysis, an energy-yielding reaction that is the reverse of a dehydration synthesis reaction. What are the scientific words meaning “energy-requiring” and “energy-releasing”? Figure 2.11 Endothermic means “energy-requiring,” and exothermic means “energy-releasing.” Decomposition reactions are the reverse of synthesis reactions in that they break bonds within larger molecules to form smaller atoms, ions, and molecules. These reactions release energy and are therefore exothermic9 reactions. In general, decomposition reactions can be represented by the following formula: As the name suggests, atoms are moved from one molecule to another. In general, these reactions can be represented as either Reactant + Product S Product(s) AB + CD S AD + BC An example of a biologically important decomposition reaction is the aerobic decomposition of glucose to form carbon dioxide and water: C6H12O6 + 6 O2 S 6 H2O + 6 CO2 Note that this reaction is exactly the reverse of the synthesis reaction in algae that we examined previously. Synthesis and decomposition reactions are often reversible in living things. A common type of decomposition reaction in biochemistry is hydrolysis,10 the reverse of dehydration synthesis (FIGURE 2.11b). In hydrolytic reactions, a covalent bond in a large molecule is broken, and the ionic components of water (H + and OH - ) are added to the products. Collectively, all of the decomposition reactions in an organism are called catabolism. A + BC S AB + C or An important exchange reaction within organisms is the phosphorylation of glucose: C6H12O6 + A Glucose P P P Adenosine triphosphate C6H11O6 Glucose phosphate P +A P P Adenosine diphosphate The sum of all of the chemical reactions in an organism, including catabolic, anabolic, and exchange reactions, is called metabolism. (Chapter 5 examines metabolism in more detail.) TELL ME WHY Why are decomposition reactions exothermic, that is, energy releasing? Exchange Reactions Water, Acids, Bases, and Salts LE A R N I N G O U TC O M E As previously noted, living things depend on organic compounds, those that contain carbon and hydrogen atoms. Living things also require a variety of inorganic chemicals, which typically lack carbon. Such inorganic substances include water, oxygen molecules, metal ions, and many acids, bases, and salts. In this section, we examine the characteristics of some of these inorganic substances. 2.17 Compare exchange reactions to synthesis and decompo- sition reactions. Exchange reactions (also called transfer reactions) have features similar to both synthesis and decomposition reactions. For instance, they involve breaking and forming covalent bonds, and they involve both endothermic and exothermic steps. + H+ Water LE A R N I N G O UTC O M E 9From Greek exo, meaning “outside,” and thermos, meaning “heat” (energy). Greek hydor, meaning “water,” and lysis, meaning “loosening.” 10From M02_BAUM7206_05_SE_C02.indd 36 2.18 Describe five qualities of water that make it vital to life. 12/1/16 3:25 PM f+ H O f – O H H f+ f – H f+ H f– O H f+ f– f+ O Hydrogen bonds H f+ f– O H f+ H       A Water, cids, Bases, and Salts H f+ f+ f+ (a) (b) Aquarius remigis, a water strider 6 mm ▲ FIGURE 2.12 The cohesiveness of liquid water. (a) Water molecules are cohesive because hydrogen bonds cause them to stick to one another. (b) One result of cohesiveness in water is surface tension, which can be strong enough to support the weight of insects known as water striders. Water is the most abundant substance in organisms, constituting 50% to 99% of their mass. Most of the special characteristics that make water vital result from the fact that a water molecule has two polar covalent bonds, which allow hydrogen bonding between water molecules and their neighbors. Among the special properties of water are the following: • Water molecules are cohesive; that is, they tend to stick to one another through hydrogen bonding (FIGURE 2.12). This property generates many special characteristics of water, including surface tension, which allows water to form a thin layer on the surface of cells. This aqueous layer is necessary for the transport of dissolved materials into and out of a cell. • Water is an excellent solvent; that is, it dissolves salts and other electrically charged molecules because it is attracted to both positive and negative charges (see Figure 2.9). • Water remains a liquid across a wider range of temperatures than other molecules of its size. This is critical because living things require water in liquid form. • Water can absorb significant amounts of heat energy without itself changing temperature. Further, when heated water molecules eventually evaporate, they take much of this absorbed energy with them. These properties moderate temperature fluctuations that would otherwise damage organisms. • Water molecules participate in many chemical reactions within cells both as reactants in hydrolysis and as products of dehydration synthesis. As we have seen, the polar bonds of water molecules dissociate salts into their component cations and anions. A similar process occurs with substances known as acids and bases. An acid is a substance that dissociates into one or more hydrogen ions (H + ) and one or more anions (FIGURE 2.13a). Acids can be inorganic molecules, such as hydrochloric acid (HCl) and sulfuric acid (H2SO4), or organic molecules, such as amino acids and nucleic acids. Familiar organic acids are found in lemon juice, black coffee, and tea. Of course, the anions of organic acids contain carbon, whereas those of inorganic acids do not. A base is a molecule that binds with H + when dissolved in water. Some bases dissociate into cations and hydroxyl ions (OH - ) (FIGURE 2.13b), which then combine with hydrogen ions to form water molecules: H + + OH - S H2O H+ Cation+ Anion– Acids and Bases – OH LE A R N I N G O U TC O M E buffers. M02_BAUM7206_05_SE_C02.indd 37 (a) Acid (b) Base ▲ FIGURE 2.13 Acids and bases. (a) cids dissociate in water into hydrogen ions and anions. (b) Many bases dissociate into hydroxyl ions and cations. A 2.19 Contrast acids, bases, and salts, and explain the role of 12/1/16 3:25 PM e 2    The Chemistry of Microbiology r pt       Cha released; therefore, pH buffers are often added to them. One common buffer used in microbiological media is KH2PO4 (potassium dihydrogen phosphate), which exists as either a weak ­ Other bases, such as household ammonia (NH3), directly accept hydrogen ions and become compound ions such as NH4+ (ammonium). Another common household base is baking soda (sodium bicarbonate, NaHCO3). Metabolism requires a relatively constant balance of acids and bases because hydrogen ions and hydroxyl ions are involved in many chemical reactions. Further, many complex molecules such as proteins lose their functional shapes when acidity changes. If the concentration of either hydrogen ions or hydroxyl ions deviates too far from normal, metabolism ceases. The concentration of hydrogen ions in a solution is expressed using a logarithmic pH scale (FIGURE 2.14). The term pH comes from potential hydrogen, which is the negative of the logarithm of the concentration of hydrogen ions. In this logarithmic scale, it is important to notice that acidity increases as pH values decrease and that each decrease by a whole number in pH indicates a 10-fold increase in acidity (hydrogen ion concentration). For example, a glass of grapefruit juice, which has a pH of 3.0, contains 10 times as many hydrogen ions as the same volume of tomato juice, which has a pH of 4.0. Similarly, tomato juice is 1000 times more acidic than pure water, which has a pH of 7.0 (neutral). Water is neutral because it dissociates into one hydrogen cation and one hydroxyl anion: H2O S H + OH Increasing concentration of H+ Extremely Acidic - Increasing concentration of OH– - pH Alkaline (basic) substances have pH values greater than 7.0. They reduce the number of free hydrogen ions by combining with them. For bases that produce hydroxyl ions, the concentration of hydroxyl ions is inversely related to the concentration of hydrogen ions. Organisms can tolerate only a certain, relatively narrow pH range. Fluctuations outside an organism’s preferred range inhibit its metabolism and may even be fatal. Most organisms contain natural buffers—substances, such as proteins, that prevent drastic changes in internal pH by removing excess hydrogen and hydroxyl ions. In a laboratory culture, the metabolic activity of microorganisms can change the pH of microbial growth solutions as nutrients are taken up and wastes are Extremely Basic 0 Battery acid 1 Hydrochloric acid 2 Lemon juice 3 Beer, vinegar 4 Wine, tomatoes 5 Black coffee 6 Urine, milk 7 Pure water 8 Seawater 9 Baking soda 10 Milk of magnesia 11 Household ammonia 12 Household bleach 13 Oven cleaner 14 Sodium hydroxide ▲ FIGURE 2.14 The pH scale. Values below 7 are acidic; values above 7 are basic. E A E E B N FICI L MICROB S Architecture-Preserving Bacteria A A A M02_BAUM7206_05_SE_C02.indd 38 With the help of Myxococcus xanthus, a bacterium commonly found in soil, a team of researchers led by mineralogist Carlos Rodríguez-Navarro of the University of Granada may have found a way to protect the stone of structures like the lhambra. In many natural environments, bacteria instigate the formation of calcite crystals like the ones in limestone. In tests conducted using samples of the limestone commonly used in historic Spanish buildings, M. xanthus formed calcite crystals that lined the stone’s pores rather than plugging them. The crystals formed by the bacteria are even more durable than the original stone, offering the potential for long-term protection. A ­ A The lhambra, a Moorish palace constructed of limestone and marble beginning in the 9th century, was built to last. But not even stone lasts forever. Wind and rain wear away the surface. cid rain reacts with the calcite crystals in limestone and marble. s years pass, stone slowly crumbles. Those who would preserve the lhambra and other historic structures face a dilemma. The microscopic pores that riddle limestone and marble make these materials particularly susceptible to weathering and decay. Sealing the stone’s pores can reduce weathering but can also lock in moisture that speeds the stone’s decay. 12/1/16 3:26 PM       Organic Macromolecules ­ ­ acid or a weak base, depending on the pH of its environment. Under acidic conditions, KH2PO4 is a base that combines with H + , neutralizing the acidic environment; in alkaline conditions, however, KH2PO4 acts as an acid, releasing hydrogen ions. Microorganisms differ in their ability to tolerate various ranges of pH. Many grow best when the pH is between 6.5 and 8.5. Photosynthetic bacteria known as cyanobacteria grow well in more basic solutions. Fungi generally tolerate acidic environments better than most prokaryotes, though acid-loving prokaryotes, called acidophiles, require acidic conditions. Some bacteria are tolerant of acid. One such bacterium is Propionibacterium acnes (prō-pē-on-ibak-tēr´ē-ŭm ak´nēz). This bacterium can cause acne in the skin, which normally has a pH of about 4.0. Another is Helicobacter pylori (hel´ĭ-kō-bak´ter pī´lō-rē)11 a curved bacterium that has been shown to cause ulcers in the stomach, where pH can fall as low as 1.5 when acid is being actively secreted. Clinical Case Study: Raw Oysters and Antacids: A Deadly Mix? focuses on how the use of antacids may increase the survival rates of certain disease-causing bacteria in the stomach. Microorganisms can change the pH of their environment by utilizing acids and bases and by producing acidic or basic wastes. For example, fermentative microorganisms form organic acids from the decomposition of sugar, and the bacterium Thiobacillus (thī-ō-bă-sil´ŭs) can reduce the pH of its environment to 0.0. Some mining companies rely on acid produced by this bacterium to dissolve uranium and copper from low-grade ore. Scientists measure pH with a pH meter or with test papers impregnated with chemicals (such as litmus or phenol red) that change color in response to pH. In a microbiological laboratory, changes in color of such pH indicators incorporated into microbial growth media are commonly used to distinguish among bacterial genera. Mg(OH)2 Magnesium hydroxide + 2 HCl MgCl2 + Hydrochloric Magnesium acid chloride (salt) 2H2O Water Cations and anions of salts are electrolytes. A cell uses electrolytes to create electrical differences between its inside and outside, to transfer electrons from one location to another, and as important components of many enzymes. Certain organisms also use salts such as calcium carbonate (CaCO3) to provide structure and support for their cells. TELL ME WHY Why does the neutralization of an acid by a base often produce water? Organic Macromolecules Inorganic molecules play important roles in an organism’s metabolism; however, water excluded, they compose only about 1.5% of its mass. Inorganic molecules are typically too small and too simple to constitute an organism’s basic structures or to perform the complicated chemical reactions required of life. Raw Oysters and Antacids: A Deadly Mix? A infections die. V. vulnificus especially affects the immunocompromised and people with long-term liver disease. Researchers have discovered that taking antacids may make people more susceptible to becoming ill from V. vulnificus. They found that antacids in a simulated gastric environment significantly increased the survival rate of V. vulnificus. 1. Why are patients who take antacids at greater risk for infections with V. vulnificus? 2. Will antacids raise or lower the p H ­ The highly acidic environment of the stomach kills most bacteria before they cause disease. One bacterium that can slightly tolerate conditions as it passes through the stomach is Vibrio vulnificus—a bacterium commonly ingested by eating raw tainted oysters. The bacterium cannot be seen, tasted, or smelled in food or water. V. vulnificus is an emerging pathogen and a growing cause of food poisoning in the United States: it triggers vomiting, diarrhea, and abdominal pain. The pathogen can also infect the bloodstream, causing life-threatening illness characterized by fever, chills, skin lesions, and deadly loss of blood pressure. bout 50% of patients with bloodstream 11The As we have seen, a salt is a compound that dissociates in water into cations and anions other than H + and OH - . Acids and hydroxyl-yielding bases neutralize each other during exchange reactions that produce water and salt. For instance, milk of magnesia (magnesium hydroxide) is an antacid used to neutralize excess stomach acid. The chemical reaction is of the stomach? 3. Other than refraining from antacids, what can people do to reduce their risk of infection? Reference: dapted from MMWR 45:621–624, 1996. A CLINICAL CASE STUDY Salts name pylori refers to the pylorus, a region of the stomach. M02_BAUM7206_05_SE_C02.indd 39 12/1/16 3:26 PM e 2    The Chemistry of Microbiology r pt       Cha These functions are fulfilled by organic molecules, which are generally larger and much more complex. Functional Groups LE A R N I N G O U TC O M E 2.20 Define functional group as it relates to organic chemistry. As we have seen, organic molecules contain carbon and hydrogen atoms, and each carbon atom can form four covalent bonds with other atoms (see Figure 2.5c and d). Carbon atoms that are linked together in branched chains, unbranched chains, and rings provide the basic frameworks of organic molecules. Atoms of other elements are bound to these carbon frameworks to form an unlimited number of compounds. Besides carbon and hydrogen, the most common elements in organic compounds are oxygen, nitrogen, phosphorus, and sulfur. Other elements, such as iron, copper, molybdenum, manganese, zinc, and iodine, are important in some proteins. Atoms often appear in certain common arrangements called functional groups that are common in organic molecules. For example, ¬ NH2, the amino functional group, is found in all amino acids, and ¬ OH, the hydroxyl functional group,12 is common to all alcohols. When a class of organic molecules is discussed, the letter R (for residue) designates atoms in the compound that vary from one molecule to another. The symbol R ¬ OH, therefore, represents the general formula for an alcohol. TABLE 2.3 describes some common functional groups of organic molecules. There is a great variety of organic compounds, but all organisms use certain basic types. These molecules—known as macromolecules because they are very large—are lipids, carbohydrates, proteins, and nucleic acids. TABLE 2.3 F  unctional Groups of Organic Molecules and Some Classes of Compounds in Which They Are Found Name H OH Class of Compounds ydroxyl A Structure lcohol O CH2 R¿ C R¿ Terminal carbonyl—a carbon atom (in R group) on only one side O R C C Carboxyl C ldehyde O mino acid Protein H Fatty acid A H R Carbohydrate H O R Ketone A R Polysaccharide Internal carbonyl—a carbon atom (in R group) on each side O mino acid Disaccharide A CH2 ther mino A R E A Monosaccharide mino acid Protein NH2 H C R CH2 O Wax Sulfhydryl SH mino acid Protein Organic phosphate O P Phospholipid Nucleotide O A CH2 Fat R¿ OH R ster A R E O TP OH 12Note that the hydroxyl functional group is not the same thing as the hydroxyl ion, because the former is covalently bonded to a carbon atom and the latter is unbound. M02_BAUM7206_05_SE_C02.indd 40 12/1/16 3:26 PM       Organic Macromolecules Lipids triglycerides because they contain three fatty acid molecules linked to glycerol. The three fatty acids in a fat molecule may be identical or different from one another, but each usually has 12–20 carbon atoms. An important difference among fatty acids is the presence and location of double bonds between carbon atoms. When carbon atoms are linked solely by single bonds, every carbon atom, with the exception of the terminal ones, is covalently linked to two hydrogen atoms. Such a fatty acid is saturated with hydrogen and is thus termed a saturated fatty acid (FIGURE 2.15b). In contrast, unsaturated fatty acids contain one double bond between adjacent carbon atoms (monounsaturated fatty acid) or more than one double bond between carbon atoms (polyunsatu­ rated fatty acid). Triglycerides with at least one polyunsaturated fatty acid are polyunsaturated fats. Saturated fats (composed solely of saturated fatty acids), like those found in animals, are usually solid at room temperature because their fatty acids can be packed closely together. Unsaturated fatty acids, by contrast, are bent at every double bond, and so unsaturated fats cannot be packed tightly; they remain liquid at room temperature. Most fats in plants are unsaturated or polyunsaturated. TABLE 2.4 on p. 42 compares the structures and melting points of four common fatty acids. Fats contain an abundance of energy stored in their carbon-carbon covalent bonds. Indeed, a major role of fats in organisms is to store energy. Fats can be catabolized to provide energy for movement, synthesis, and transport (Chapter 5). LE A R N I N G O U TC O M E S 2.21 Describe the structure of a molecule of fat (triglyceride), and compare it to that of phospholipids, waxes, and steroids. 2.22 Distinguish among saturated, unsaturated, and polyunsaturated fatty acids. Lipids are a diverse group of organic macromolecules not composed of regular subunits. They have one common trait— they are hydrophobic;13 that is, they are insoluble in water. Lipids have little or no affinity for water because they are composed almost entirely of carbon and hydrogen atoms linked by nonpolar covalent bonds. Because these bonds are nonpolar, they have no attraction to the polar bonds of water molecules. To look at it another way, the polar water molecules are attracted to each other and exclude the nonpolar lipid molecules. There are four major groups of lipids in cells: fats (triglycerides), phospholipids, waxes, and steroids. Fats H H H O C OH HO C 3 H2O O H C OH HO C OH HO H C O O C O C Dehydration synthesis O H ­ Organisms make fats via dehydration synthesis reactions that form esters between an alcohol named glycerol and three fatty acids, which are long hydrocarbon chains capped by a carboxyl functional group (FIGURE 2.15a). Fats are also called C H O C O H H C Ester bond C O C H Glycerol + 3 fatty acids Fat (triglyceride) (a) CH2 O C CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH3 CH2 HO Saturated fatty acid CH2 O C CH2 CH2 CH2 CH2 HO CH2 CH2 CH CH CH2 Monounsaturated fatty acid CH2 CH2 (b) CH2 CH2 CH2 CH3 Figure 2.15 Stearic acid, palmitic acid, and oleic acid. 13From CH2 ▲ FIGURE 2.15 Fats (triglycerides). (a) Fats are made in dehydration synthesis reactions that form ester bonds between a glycerol molecule and three fatty acids. (b) Saturated fatty acids have only single bonds between their carbon atoms, whereas unsaturated fatty acids have double bonds between carbon atoms. Scientists often use abbreviated diagrams of fatty acids in which each angle represents a carbon atom and most hydrogen atoms are not shown, as seen in part (a) of this figure. According to Table 2.4, which fatty acids are shown in Figure 2.15a? Greek hydor, meaning “water,” and phobos, meaning “fear.” M02_BAUM7206_05_SE_C02.indd 41 12/1/16 3:26 PM TABLE 2.4 e 2    The Chemistry of Microbiology r pt       Cha Common Fatty Acids in Fats and Cell Membranes Numbers of Carbon Atoms: Double Bonds Type of Fatty Acid 16:0 Saturated Structure and Formula CH3(CH2)14COOH CH3 CH2 CH2 CH2 CH2 CH2 CH2 CH2 Palmitic acid 63°C Stearic acid 70°C Oleic acid 16°C Linoleic acid - 5°C O OH Hydrocarbon chain Saturated Melting Point C CH2 CH2 CH2 CH2 CH2 CH2 CH2 18:0 Common Name Ch3(Ch2)16COOh CH3 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 O C OH 18:1 Monounsaturated Ch3(CH2)7CH “ CH(CH2)7COOH Double bond CH2 CH2 CH2 CH CH2 CH2 CH2 CH3 18:2 Polyunsaturated CH CH2 CH2 CH2 CH2 O C OH CH2 Carboxyl group CH2 CH2 CH2 Ch3(Ch2)4(Ch “ ChCh2)2(Ch2)6COOh CH2 CH2 CH2 CH CH3 CH2 CH2 CH2 CH2 CH CH CH2 CH2 CH2 CH2 O C OH CH2 CH Phospholipids Waxes Waxes contain one long-chain fatty acid linked covalently to a long-chain alcohol by an ester bond. Waxes do not have hydrophilic heads; thus, they are completely water insoluble. Certain microorganisms, such as Mycobacterium tuberculosis (mī´kō-baktēr´ē-ŭm too-ber-kyū-lō´sis) are surrounded by a waxy wall, making them resistant to drying. Some marine microbes use waxes instead of fats as energy storage molecules. ­ Phospholipids are similar to fats, but they contain only two fatty acid chains instead of three. In phospholipids, the third carbon atom of glycerol is linked to a phosphate (PO4) functional group instead of a fatty acid (FIGURE 2.16a). Like fats, different phospholipids contain different fatty acids. Small organic groups linked to the phosphate group provide additional variety among phospholipid molecules. The fatty acid “tail” portion of a phospholipid molecule is nonpolar and thus hydrophobic, whereas the phospholipid “head” is polar and thus hydrophilic.14 As a result, phospholipids placed in a watery environment will always self-assemble into forms that keep the fatty acid tails away from water. One way t hey do t h i s i s to for m a pho sphol ipid bi layer (FIGURE 2.16c). The fatty acid tails, which are hydrophobic, congregate in the water-free interior of bilayers. The polar phosphate heads, which are hydrophilic, orient toward the water on either side of the bilayer. Phospholipid bilayers make up the membranes surrounding cells as well as the internal membranes of plant, fungal, and animal cells. 14From Greek philos, meaning “love.” M02_BAUM7206_05_SE_C02.indd 42 Steroids A final group of lipids are steroids. Steroids consist of four rings (each containing five or six carbon atoms) that are fused to one another and attached to various side chains and functional groups (FIGURE 2.17a). Steroids play many roles in human metabolism. Some act as hormones; another steroid, cholesterol, is perhaps familiar to you as a less desirable component of food. However, cholesterol is also an essential part of the phospholipid 12/1/16 3:26 PM       Organic Macromolecules Organic group Cell O Phosphate group Glycerol H2C O P O O– H O C CH2 O O C C Unsaturated fatty acid Hydrophilic (polar) head O (b) Symbol Double bond Saturated fatty acid (c) Phospholipid bilayer Hydrophobic (nonpolar) tails Unsaturated fatty acid (a) Phospholipid A ▲ FIGURE 2.16 Phospholipids. (a) phospholipid is composed of a hydrophilic (polar) “head,” which consists of glycerol and a phosphate group, and two hydrophobic (nonpolar) fatty acid “tails.” (b) The symbol used to represent phospholipids. (c) In water, phospholipids can self-assemble into spherical bilayers. Phospholipids containing unsaturated fatty acids do not pack together as tightly as those containing saturated fatty acids. bilayer membrane surrounding an animal cell. Cells of fungi, plants, and one group of bacteria (mycoplasmas) have similar sterol molecules in their membranes. Sterols, which are steroids with an ¬ OH functional group, interfere with the tight packing of the fatty acid chains of phospholipids (FIGURE 2.17b). This keeps the membranes fluid and flexible at low temperatures. Without steroids such as cholesterol, the membranes of cells would become stiff and inflexible in the cold. Carbohydrates, proteins, and nucleic acid macromolecules are composed of simpler subunits known as monomers,15 which are basic building blocks. The monomers of these macromolecules are joined together to form chains of monomers called polymers.16 Some macromolecular polymers are composed of hundreds of thousands of monomers. Carbohydrates LE A R N I N G O U TC O M E 2.23 Discuss the roles carbohydrates play in living systems. Carbohydrates are organic molecules composed solely of atoms of carbon, hydrogen, and oxygen. Most carbohydrate 15From 16From Greek mono, meaning “one,” and meris, meaning “part.” Greek poly, meaning “many,” and meris, meaning “part.” M02_BAUM7206_05_SE_C02.indd 43 compounds contain an equal number of oxygen and carbon atoms and twice as many hydrogen atoms as carbon atoms, so the general formula for a carbohydrate is (CH2 O)n , where n indicates the number of CH2O units. Carbohydrates play many important roles in organisms. Large carbohydrates, such as starch and glycogen, are used for the long-term storage of chemical energy, and smaller carbohydrate molecules serve as a ready energy source, form part of the backbones of DNA and RNA, and may be converted into amino acids. Additionally, polymers of carbohydrate form the cell walls of most fungi, plants, algae, and prokaryotes and are involved in intercellular interactions between animal cells. Monosaccharides The simplest carbohydrates are monosaccharides17—simple sugars (FIGURE 2.18). The general names for the classes of monosaccharides are formed from a prefix indicating the number of carbon atoms and from the suffix -ose. For example, pentoses are sugars with five carbon atoms, and hexoses are sugars with six carbon atoms. Pentoses and hexoses are particularly important in cellular metabolism. For example, deoxyribose, which is the sugar component of DNA, is a pentose. Glucose is a hexose and the primary energy molecule of cells, and fructose is a hexose found in fruit. Chemists assign numbers to the carbon atoms. 17From Greek sakcharon, meaning “sugar.” 12/1/16 3:26 PM r e 2    The Chemistry of Microbiology pt       Cha C C C C C C C C C C C C C C C C C (a) H 1C O H 2C OH HO 3C H 4 H 4C OH HO H 5C OH H 6 CH2OH 5 O H OH 3 H H 2 H 1 4 OH HO H OH c configuration 6 CH2OH 6 CH2OH 5 O H OH H 3 2 OH 1 H H OH d configuration (a) Glucose H H HO HO O 1C H O 2C N C 3C H H 4C OH H 5C OH H CH3 4 OH Acetyl group 6 CH2OH 5 O OH H 3 2 1 H NH H C O CH3 6 CH2OH Phospholipids OH (b) N-acetylglucosamine A ▲ FIGURE 2.18 Monosaccharides (simple sugars). lthough A Cholesterol A ­ simple sugars may exist as either linear molecules (at left) or rings (at right), energy dynamics in the watery cytoplasm of cells generally favor ring forms. (a) Glucose, a hexose, is the primary energy source for cellular metabolism and an important monomer in many larger carbohydrates. Chemists number the carbon atoms as shown. lpha and beta ring configurations differ in the location of oxygen bound to carbon 1. (b) N-acetylglucosamine (N G), a monomer in bacterial cell walls. ▲ FIGURE 2.17 Steroids. (a) Steroids are lipids characterized by four “fused” rings. (b) The steroid cholesterol, which has a short chain of atoms extending from the fourth ring, functions in animal and protozoan cell membranes to prevent packing of phospholipids, thereby keeping the membranes fluid at low temperatures. Monosaccharides may exist as linear molecules, but because of energy dynamics, each usually takes a cyclic (ring) form. In some cases, more than one cyclic structure may exist. For example, glucose can assume an alpha (a) configuration or a beta (β) configuration (see FIGURE 2.18a). As we will see, these configurations play important roles in the formation of different polymers. Disaccharides When two monosaccharide molecules are linked together via dehydration synthesis, the result is a disaccharide. For example, the linkage of two hexoses, glucose and fructose, forms sucrose (table sugar) in a dehydration reaction (FIGURE 2.19a). Other disaccharides include maltose (malt sugar) and lactose (milk sugar). Disaccharides can be broken M02_BAUM7206_05_SE_C02.indd 44 down via hydrolysis into their constituent monosaccharides (FIGURE 2.19b). Polysaccharides Polysaccharides are polymers composed of tens, hundreds, or thousands of monosaccharides that have been covalently linked in dehydration synthesis reactions. Even polysaccharides that contain only glucose monomers can be quite diverse because they can differ according to their monosaccharide monomer configurations (either alpha or beta) and their shapes (either branched or unbranched). Cellulose, the main constituent of the cell walls of plants and some green algae, is a long unbranched molecule that contains only b@monomers of glucose linked between carbons 1 and 4 of alternating monomers; such bonds are termed b@1,4 bonds (FIGURE 2.20a). Amylose, a starch storage compound in plants, has only a@1,4 bonds and is unbranched (FIGURE 2.20b); glycogen, a storage molecule formed in the liver and muscle cells of animals, is a highly branched molecule with both a@1,4 and a@1,6 bonds (FIGURE 2.20c). The cell walls of bacteria are composed of peptidogly­ can, which is made of polysaccharides and amino acids (see Figure 2.15). Polysaccharides may also be linked to lipids to form glycolipids, which can form cell markers such as those involved in the ABO blood typing system in humans. ­ (b) ­ Cell membrane 12/1/16 3:26 PM       Organic Macromolecules CH2OH H HO O H OH H H CH2OH H + HOCH2 OH HO OH H O HO OH Glucose C6H12O6 H H CH2OH Energy H HO O H H OH O H O OH Fructose C6H12O6 + HOCH2 H H OH H HO CH2OH + H2O H Sucrose C12H22O11 Water H 2O + (a) Dehydration synthesis of sucrose CH2OH H HO O H OH H CH2OH HOCH2 H H O OH H OH Sucrose C12H22O11 O H HO CH2OH H + H2O Energy H + Water H2O HO O H OH H H OH Glucose C6H12O6 H OH + HOCH2 HO H OH + O H HO CH2OH H Fructose C6H12O6 (b) Hydrolysis of sucrose ▲ FIGURE 2.19 Disaccharides. (a) Formation of the disaccharide sucrose via dehydration synthesis. (b) Breakdown of sucrose via hydrolysis. LE A R N I N G O U TC O M E S 2.24 Describe five general functions of proteins in organisms. 2.25 Sketch and label four levels of protein structure. The most complex organic compounds are proteins, which are composed mostly of carbon, hydrogen, oxygen, nitrogen, and sulfur. Proteins perform many functions in cells, including the following: • Structure. Proteins are structural components found in cell walls, in membranes, and within cells themselves. Proteins are also the primary structural material of hair, nails, the outer cells of skin, muscle, and flagella and cilia (the last two act to move microorganisms through their environment). • Enzymatic catalysis. Catalysts are chemicals that enhance the speed or likelihood of a chemical reaction. Protein catalysts in cells are called enzymes. • Regulation. Some proteins regulate cell function by stimulating or hindering either the action of other proteins or the expression of genes. Hormones are examples of regulatory proteins. • Transportation. Certain proteins act as channels and “pumps” that move substances into or out of cells. • Defense and offense. Antibodies and complement are examples of proteins that defend your body against pathogens. A protein’s function is dependent on its shape, which is determined by the molecular structures of its constituent parts and by the bonds within the molecule. M02_BAUM7206_05_SE_C02.indd 45 Amino Acids Proteins are polymers composed of monomers called amino acids. Amino acids contain a basic amino group ( ¬ NH2), a hydrogen atom, and an acidic carboxyl group ( ¬ COOH). All attach to the same carbon atom, which is known as the a@carbon (FIGURE 2.21). A fourth bond attaches the a@carbon to a side group ( ¬ R) that varies among different amino acids. The side group may be a single hydrogen atom, various chains, or various complex ring structures. Hundreds of amino acids are possible, but most organisms use only 21 amino acids in synthesizing proteins.18 The different side groups affect the way amino acids interact with one another within a given protein as well as how a protein interacts with other molecules. A change in an amino acid’s side group may seriously interfere with a protein’s normal function. Because amino acids contain both an acidic carboxyl group and a basic amino group, they have both positive and negative charges and are easily soluble in water. Aqueous solutions of organic molecules such as amino acids and simple sugars bend light rays passing through the solution. Molecules known as d forms19 bend light rays clockwise; other molecules bend light rays counterclockwise and are known as l forms.20 Many organic molecules exist as both d and l forms that are stereoisomers of one another; that is, they have the same atoms and functional groups but are mirror images of each other (FIGURE 2.22). Amino acids in proteins are almost always l forms—except for glycine, which does not have a stereoisomer. ­ Proteins 18Although 20 amino acids are commonly listed, the genes of almost all organisms code for 21 amino acids; the 21st is selenocysteine. 19From Latin dexter, meaning “on the right.” 20From Latin laevus, meaning “on the left.” 12/1/16 3:26 PM r e 2    The Chemistry of Microbiology pt       Cha d-1,4 bonds 6 CH2OH 5 O CH2OH OH O O 4 O 1 O 3 Hydrogen bond O O 2 OH CH2 OH OH OH OH OH CH2 CH2 O O O O CH2OH OH 5 OH c-1,4 bonds CH2OH 6 CH2OH O O O O (a) Cellulose Hydrogen bonds CH2OH CH2OH O O O 1 4 O O 3 O O O 2 OH (b) Amylose (unbranched) OH 6 CH OH 2 5 O O c-1,6 bond 1 O O 3 OH c-1,4 bond CH2OH 4 O 2 OH OH CH2OH CH2OH O O O O CH2 CH2OH O O OH (c) Glycogen OH OH c-1,4 bond O O c-1,6 bond O OH OH CH2 CH2 O O O O OH O OH A ▲ FIGURE 2.20 Polysaccharides. ll three polysaccharides shown here are composed solely of glucose but differ in the configuration of the A glucose monomers and the amount of branching. (a) Cellulose, the major structural material in plants, is unbranched and contains only b@1,4 bonds. (b) mylose is an unbranched plant starch with only a@1,4 bonds. (c) Glycogen, a highly branched storage molecule in animals, is composed of glucose monomers linked by a@1,4 or a@1,6 bonds. Interestingly, organisms almost always use d sugars in metabolism and polysaccharides. Rare stereoisomers—d amino acids and l sugars—do exist in some bacterial cell walls and in some antibiotics. Peptide Bonds Cells link amino acids together in chains that somewhat resemble beads on a necklace. By a dehydration synthesis reaction, a covalent bond is formed between the carbon of the carboxyl group of M02_BAUM7206_05_SE_C02.indd 46 one amino acid and the nitrogen of the amino group of the next amino acid in the chain (FIGURE 2.23). Cells follow the organism’s genetic instructions to link amino acids together in precise sequences. (Chapter 7 examines this process in more detail.) Scientists refer to covalent bonds between amino acids by a special name: peptide21 bonds. A molecule composed of two peptone, the name given to short chains of amino acids resulting from the partial digestion of protein. 21From 12/1/16 3:26 PM       Organic Macromolecules O C Side c-carbon group R H N H C C H2N OH H Amino group H O C CH2 COOH H2N H Carboxyl group COOH H2N SH Se CH2 CH2 CH2 CH2 C H Glycine (Gly) (a) C COOH H Phenylalanine (Phe) NH2 H2N C COOH H Cysteine (Cys) Selenocysteine (SeCys) H2N C CH3 COOH H Glutamine (Gln) H2N C COOH H Alanine (Ala) (b) ▲ FIGURE 2.21 Amino acids. (a) The basic structure of an amino acid. The central a@carbon is attached to an amino group, a hydrogen atom, a carboxyl group, and a side group ( ¬ R group) that varies among amino acids. (b) Some selected amino acids, with their side groups highlighted. Note that each amino acid has a distinctive abbreviation. structure is critical to understanding certain specific chemical reactions, the action of antibiotics, and specific defenses against pathogens. Every protein has at least three levels of structure, and some proteins have four levels. Mirror H H H N H H C C H H H C C H O O H L-Alanine (Ala) Left H O C C H O H H N H H D-Alanine (Ala) Right ­ ▲ FIGURE 2.22 Stereoisomers, molecules that are mirror images of one another. When dissolved in water, d isomers bend light ­ clockwise, and l forms bend light counterclockwise. Just as a right-handed glove does not fit a left hand, so a d stereoisomer cannot be substituted for an l stereoisomer in metabolic reactions. amino acids linked together by a single peptide bond is called a d ip ept ide; longer c h a i n s of a m i no ac id s a r e c a l le d polypeptides. Protein Structure Proteins are unbranched polypeptides composed of hundreds to thousands of amino acids linked together in specific patterns as determined by genes. The structure of a protein molecule is directly related to its function; therefore, understanding protein M02_BAUM7206_05_SE_C02.indd 47 • Primary structure. The primary structure of a protein is its sequence of amino acids (FIGURE 2.24a). Cells use many different types of amino acids in proteins, though not every protein contains all types. The primary structures of proteins vary widely in length and amino acid sequence. A change in a single amino acid in the sequence can drastically affect a protein’s overall structure and function, though this is not always the case. For instance, a change in a particular amino acid in the primary structure of a sheep brain protein, called cellular prion (pre′on) protein, may result in a disease called scrapie. The altered protein has spread into cows, causing mad cow disease, and spread from cows into humans, causing variant Creutzfeldt­Jakob (kroytsfelt-yah-kŭp) disease.22 • Secondary structure. Ionic bonds, hydrogen bonds, and hydrophobic and hydrophilic characteristics cause many polypeptide chains to fold into either coils called a@helices or accordion-like structures called b@pleated sheets (FIGURE 2.24b). Proteins are typically composed of both a@helices and b@pleated sheets linked by short sequences of amino acids that do not show such secondary structure. Because of its primary structure, the protein that causes variant Creutzfeldt-Jakob disease has b@pleated sheets in locations where the normal protein has a@helices (see Figure 13.22). • Tertiary structure. Polypeptides further fold into complex three-dimensional shapes that are not repetitive like a@helices and b@pleated sheets (FIGURE 2.24c) but are uniquely designed to accomplish the function of the protein. Scientists are only beginning to understand the interactions that determine tertiary structure, but it is clear that covalent bonds between ¬ R groups of amino acids, 22Named for the two German neurobiologists who first described the disease. 12/1/16 3:26 PM H H2N r O C R1 e 2    The Chemistry of Microbiology pt       Cha + C OH Carboxyl group H2O R2 H N H C COOH H Amino group Amino acid 1 Dehydration synthesis H2N H O C C R1 ◀ FIGURE 2.23 The linkage of amino acids by peptide bonds via a dehydration reaction. In this R2 N C H H COOH reaction, removing a hydroxyl group from amino acid 1 and a hydrogen atom from amino acid 2 produces a dipeptide—which is two amino acids linked by a single peptide bond—and a molecule of water. Peptide bond Amino acid 2 Dipeptide H N H C Amino acid 1 C O Peptide bond H N C R2 d-pleated sheet Amino acid 2 O C H O c-helix Hydrogen bond N C R3 Amino acid 3 C H N (b) Secondary structure C R4 R1 O Amino acid 4 d-pleated sheets C (a) Primary structure ▲ FIGURE 2.24 Levels of protein structure. (a) protein’s primary structure c-helix A is the sequence of amino acids in a polypeptide. (b) Secondary structure arises as a result of interactions, such as hydrogen bonding, between regions of the polypeptide. Secondary structure takes two basic shapes: a@helices and b@pleated sheets. (c) The more complex tertiary structure is a three-dimensional shape defined by further hydrogen bonding as well as disulfide bridges between neighboring cysteine amino acid molecules. (d) Proteins that are composed of more than one polypeptide chain have a quaternary structure. M02_BAUM7206_05_SE_C02.indd 48 (c) Tertiary structure (d) Quaternary structure: Two or more polypeptides act together as a single protein. 12/1/16 3:26 PM       Organic Macromolecules Nucleotides and Nucleic Acids hydrogen bonds, ionic bonds, and other molecular interactions are important. For instance, nonpolar (hydrophobic) side chains fold into the interior of molecules, away from the presence of water. • Some proteins form strong covalent bonds between sulfur atoms of cysteine amino acids that are brought into proximity by the folding of the polypeptide. These disulfide bridges are critical in maintaining tertiary structure of many proteins. • Quaternary structure. Some proteins are composed of two or more polypeptide chains linked together by disulfide bridges or other bonds. The overall shape of such a protein may be fibrous (threadlike) or more globular (FIGURE 2.24d). LE A R N I N G O UTC O M E S 2.26 Describe the basic structure of a nucleotide. 2.27 Compare and contrast DNA and RNA. 2.28 Contrast the structures of ATP, ADP, and AMP. The nucleic acids deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) are the vital genetic material of cells and viruses. Moreover, RNA, acting as an enzyme, binds amino acids together to form polypeptides. Both DNA and RNA are unbranched macromolecular polymers that differ primarily in the structures of their monomers, which we discuss next. Organisms may further modify proteins by combining them with other organic or inorganic molecules. For instance, glycoproteins are proteins covalently bound with carbohydrates, lipoproteins are proteins bonded with lipids, and nucleoproteins are proteins bonded with nucleic acids. Because protein shape determines protein function, anything that severely interrupts shape also disrupts function. As we have seen, amino acid substitution can alter shape and function. Additionally, physical and chemical factors, such as heat, changes in pH, and salt concentration, can interfere with hydrogen and ionic bonding between parts within a protein. This in turn can disrupt the three-dimensional structure. This process is called denaturation. Denaturation can be temporary (if the denatured protein is able to return to its original shape again) or permanent. O CH2 O O N H N N Thymine (T) (Used in DNA) O H Pentose sugar N (a) H H O Adenine (A) (Used in DNA and RNA) N N H NH2 H N N Guanine (G) (Used in DNA and RNA) H H OH H OH CH2OH O H H H Deoxyribose (b) M02_BAUM7206_05_SE_C02.indd 49 N H NH2 CH2OH O N H H H H o N ◀ FIGURE 2.25 Nucleotides. (a) The basic structure of nucleotides, each of which is composed of a phosphate, a pentose sugar, and a nitrogenous base. (b) The pentose sugars deoxyribose, which is found in deoxyribonucleic acid (DN ), and ribose, V e T which is found in ribonucleic acid (RN ). (c) The nitrogenous bases, which are either the doubleringed purines adenine or guanine, or the single-ringed pyrimidines thymine, cytosine, or uracil. How does a nucleoside differ from a nucleotide? id Purine or pyrimidine nitrogenous base CH3 utor O– Pyrimidines NH2 A P N A O Purines H Figure 2.25 A nucleoside is composed only of a nitrogenous base and a sugar, whereas a nucleotide has a base, sugar, and phosphate. –O Phosphate group Nucleotides and Nucleosides Each monomer of nucleic acids is a nucleotide and consists of three parts (FIGURE 2.25a): (1) phosphate (PO43- ); (2) a pentose sugar, either deoxyribose or ribose (FIGURE 2.25b); and (3) one of five cyclic (ring-shaped) nitrogenous bases: adenine (A), g uanine (G), cytosine (C), thymine (T), or uracil (U) (FIGURE 2.25c). Adenine and guanine are double-ringed molecules of a class called purines, whereas cytosine, thymine, and uracil have single rings and are pyrimidines. DNA contains A, G, C, and T bases, whereas RNA contains A, G, C, and U bases. As their names suggest, DNA nucleotides contain deoxyribose, and RNA nucleotides contain ribose. The similarly named nucleosides are nucleotides lacking phosphate; that is, a nucleoside is one of the nitrogenous bases attached only to a sugar. O Cytosine (C) (Used in DNA and RNA) H O OH H H OH OH H C N H N H H O Uracil (U) (Used in RNA) Ribose (c) 12/1/16 3:26 PM r pt       e 2    The Chemistry of Microbiology Each nucleotide or nucleoside is also named for the base it contains. Thus, a nucleotide made with ribose, uracil, and phosphate is a uracil RNA nucleotide, which is also called a uracil ribonucleotide. Likewise, a nucleoside composed of adenine and deoxyribose is an adenine DNA nucleoside (or adenine deoxyribonucleoside). To explore the structure and function of nucleotides more fully, scan the QR code with your smart phone or tablet. The atoms of the bases in nucleotides are arranged in such a manner that hydrogen bonds readily form between specific bases of two adjacent nucleic acid chains. Three hydrogen bonds form between an adjacent pair composed of cytosine (C) and guanine (G), whereas two hydrogen bonds form between an adjacent pair composed of adenine (A) and thymine (T) in DNA (FIGURE 2.26b) or between an adjacent pair composed of adenine (A) and uracil (U) in RNA. Hydrogen bonds do not readily form between other combinations of nucleotide bases; for example, adenine does not readily pair with cytosine, guanine, or another adenine nucleotide. In cells, DNA molecules are double stranded. Most DNA viruses (viruses that use DNA as a genome) also use double-stranded DNA, though some single-stranded DNA viruses are known. The two strands of double-stranded DNA are complementary to one another; that is, the specificity of nucleotide base pairing ensures that opposite strands are composed of complementary nucleotides. For instance, if one strand has the sequence AATGCT, then its complement has TTACGA. The two strands are also antiparallel; that is, they run in opposite directions. One strand runs from the 3′ end to Nucleic Acid Structure Nucleic acids, like polysaccharides and proteins, are polymers. They are composed of nucleotides linked by covalent bonds between the phosphate of one nucleotide and the sugar of the next. Polymerization results in a linear spine composed of alternating sugars and phosphates, with bases extending from it rather like the teeth of a comb (FIGURE 2.26a). The two ends of a chain of nucleotides are different. At one end, called the 5′ end23 (five prime end), carbon 5′ of the sugar is attached to a phosphate group. At the other end (the 3′ end), carbon 3′ of the sugar is attached to a hydroxyl group. 5¿ end P Adenine base 5¿ C O 4¿ Deoxyribose Phosphate Three hydrogen bonds 1¿ 3¿ A T G C T A C C 2¿ 3¿ ◀ FIGURE 2.26 General nucleic acid structure. (a) Nucleotides are polymerized to form chains in which the nitrogenous bases extend from a sugar-phosphate backbone like the teeth of a comb. (b) Specific pairs of nitrogenous bases form hydrogen bonds between adjacent nucleotide chains to form the familiar DN double helix. How can you determine that the molecule in (a) is DNA and not RNA? A 5¿ Two hydrogen bonds G Figure 2.26 It is DNA because its nucleotides have deoxyribose sugar and because some of them have thymine bases (not uracil, as in RNA). Cha G A P Guanine base C O P Thymine base C O P C Sugarphosphate backbones Cytosine base O 3¿ end OH (a) 5¿ 3¿ (b) ­ 23Carbon atoms in organic molecules are commonly identified by numbers. In a nucleotide, carbon atoms 1, 2, 3, and so on, belong to the base, and carbon atoms 1′, 2′, 3′, and so on, belong to the sugar. M02_BAUM7206_05_SE_C02.indd 50 12/1/16 3:26 PM       Organic Macromolecules TABLE 2.5 Comparison of Nucleic Acids DNA RNA Sugar Deoxyribose Ribose and G A Purine nucleotides A Characteristic and G U and C Double stranded in cells and in most DN viruses; single stranded in parvoviruses Single stranded in cells and in most RN viruses; double stranded in reoviruses Function Genetic material of all cells and DN viruses Protein synthesis in all cells; genetic material of RN viruses ­ triphosphate (ă-den´ō-sēn trī-fos´fāt) or ATP, which has three phosphate groups (FIGURE 2.27). ATP is the principal, short-term, recyclable energy supply for cells. When the phosphate bonds of ATP are broken, a significant amount of energy is released; in fact, more energy is released from phosphate bonds than is released from most other covalent bonds. For this reason, the phosphate-phosphate bonds of ATP are known as high­energy bonds, and to show these specialized bonds, ATP can be symbolized as Nucleic Acid Function M02_BAUM7206_05_SE_C02.indd 51 P P ­ Energy is released when ATP is converted to ADP and when phosphate is removed from ADP to form AMP, though the latter reaction is not as common in cells. Energy released from the phosphate bonds of ATP is used for important lifesustaining activities, such as synthesis reactions, locomotion, and transportation of substances into and out of cells. Cells also use ATP as a structural molecule in the formation of coenzymes. Coenzymes such as flavin adenine dinucleotide, NH2 N N N N Adenine O O H 2C O O P O O– P O– O O P O– O– HO OH Ribose Adenosine (nucleoside) Adenosine monophosphate (AMP) Adenosine diphosphate ADP Adenosine triphosphate ATP ▲ FIGURE 2.27 ATP. denosine triphosphate ( TP), the main short-term, recyclable energy supply for cells. nergy is stored in high-energy bonds between the phosphate groups. What is the relationship between AMP and adenine ribonucleotide? E Phosphate in nucleotides and other molecules is a highly reactive functional group and can form covalent bonds with other phosphate groups to make diphosphate and triphosphate molecules. Such molecules made from ribose nucleotides are important in many metabolic reactions. The names of these molecules indicate the nucleotide base and the number of phosphate groups they contain. Thus, cells make adenosine monophosphate (AMP) from the nitrogenous base adenine, ribose sugar, and one phosphate group; adenosine diphosphate (ADP), which has two phosphate groups; and adenosine P A ATP (Adenosine Triphosphate) A Figure 2.27 AMP and adenine ribonucleotide are two names for the same thing. DNA is the genetic material of all organisms and of DNA viruses; it carries instructions for the synthesis of RNA molecules and proteins. By controlling the synthesis of enzymes and regulatory proteins, DNA controls the synthesis of all other molecules in an organism. Genetic instructions are carried in the sequence of nucleotides that make up the nucleic acid. Even though only four kinds of bases are found in DNA (A, T, G, and C), they can be sequenced in distinctive patterns that create genetic diversity and code for an infinite number of proteins, just as an alphabet of only four letters could spell a very large number of words. Cells replicate their DNA molecules and pass copies to their descendants, ensuring that each has the instructions necessary for life. Several kinds of ribonucleic acids, such as messenger RNA, transfer RNA, and ribosomal RNA, play roles in the formation of proteins, including catalyzing the synthesis of proteins. RNA molecules also function in place of DNA as the genome of RNA viruses. TABLE 2.5 compares and contrasts RNA and DNA. (We examine the synthesis and function of DNA and RNA in detail in Chapter 7.) A ­ the 5′ end, whereas its complement runs in the opposite direction, from its 5′ end to its 3′ end. Though hydrogen bonds are relatively weak bonds, thousands of them exist at normal temperatures, forming a stable, double-stranded DNA molecule that looks much like a ladder: the two deoxyribose-phosphate chains are the side rails, and base pairs form the rungs. Hydrogen bonding also twists the phosphate- deoxyribose backbones into a helix. Thus, typical DNA is a double helix. A A A T and C Number of strands A Pyrimidine nucleotides 12/1/16 3:26 PM nicotinamide nucleotide, and coenzyme A function in many metabolic reactions (as discussed in Chapter 5). A cell’s supply of ATP is limited; therefore, an important part of cellular metabolism is to replenish ATP stores. (Chapter 5 discusses the important ATP-generating reactions.) Can Spicy Food Cause Ulcers? Why do the cell membranes of microbes living in Arctic water likely contain more unsaturated fatty acids than do membranes of microbes living in hot springs? a short time, amona is feeling better—and back to enjoying her spicy food. 1. Antacids are sometimes taken to treat ulcers. Given what you’ve learned about the pH environment that H. pylori prefers, explain how antacids might help relieve the symptoms of an ulcer. 2. What do you think would happen if ­ ­ Ramona consumed a large amount of milk? Would you expect her to feel temporary relief from her ulcer, or would you expect her to feel worse? A Check your answers to Micro in the Clinic Follow-Up questions in the MasteringMicrobiology Study rea R H H T H H R T he pain and nausea continue for a month before amona finally decides to see a doctor. It turns out that she does, indeed, have an ulcer—but eating spicy food had nothing to do with it. er doctor explains that most ulcers are caused by bacteria called Helicobacter pylori, which thrive in low-p , highly acidic environments, such as the stomach. H. pylori usually lives in the stomach without causing any problems, but occasionally it causes inflammation in the stomach’s mucous lining, resulting in an ulcer. he doctor notes that it’s a common misconception that spicy food causes ulcers. e adds, however, that alcohol, tobacco, and stress can aggravate ulcers and slow the healing process. e prescribes antibiotics to amona to kill the bacteria. Within TELL ME WHY R       Chapter 2   The Chemistry of Microbiology a a Ch pter Summ ry Atoms (pp. 28–30) 1. Matter is anything that takes up space and has mass. Its smallest chemical units, atoms, contain negatively charged electrons orbiting a nucleus composed of uncharged neutrons and positively charged protons. 2. An element is matter composed of a single type of atom. 3. The number of protons in the nucleus of an atom is its atomic number. The sum of the masses of its protons, neutrons, and electrons is an atom’s atomic mass, which is estimated by adding the number of neutrons and protons (because electrons have little mass). 4. Isotopes are atoms of an element that differ only in the numbers of neutrons they contain. Chemical Bonds (pp. 31–35) 1. The region of space occupied by electrons is an electron shell. The number of electrons in the outermost shell, or valence shell, of an atom determines the atom’s reactivity. Most valence shells hold a maximum of eight electrons. Sharing or transferring valence electrons to fill a valence shell results in chemical bonds. M02_BAUM7206_05_SE_C02.indd 52 2. A chemical bond results when two atoms share a pair of electrons. The electronegativities of each of the atoms, which is the strength of their attraction for electrons, determines whether the bond between them will be a nonpolar covalent bond (equal sharing of electrons), a polar covalent bond (unequal sharing of electrons), or an ionic bond (giving up of electrons from one atom to another). 3. A molecule that contains atoms of more than one element is a compound. Organic compounds are those that contain carbon and hydrogen atoms. 4. An anion is an atom with an extra electron and thus a negative charge. A cation has lost an electron and thus has a positive charge. Ionic bonding between the two types of ions makes salt. When salts dissolve in water, their ions are called electrolytes. 5. Hydrogen bonds are relatively weak but important chemical bonds. They hold molecules in specific shapes and confer unique properties to water molecules. Chemical Reactions (pp. 35–36) 1. Chemical reactions result from the making or breaking of chemical bonds in a process in which reactants are changed into products. Biochemistry involves chemical reactions of life. 06/03/2018 07:08       Questions for Review 5. Metabolism is the sum of all anabolic, catabolic, and exchange chemical reactions in an organism. Water, Acids, Bases, and Salts (pp. 36–39) 1. Inorganic chemicals typically lack carbon. 2. Water is a vital inorganic compound because of its properties as a solvent, its liquidity, its great capacity to absorb heat, and its participation in chemical reactions. 3. Acids release hydrogen ions. Bases release hydroxyl anions. The relative strength of each is assessed on a logarithmic pH scale, which measures the hydrogen ion concentration in a substance. 4. Buffers are substances that prevent drastic changes in pH. Organic Macromolecules (pp. 39–52) 1. Certain groups of atoms in common arrangements, called functional groups, are found in organic macromolecules. Monomers are simple subunits that can be covalently linked to form chainlike polymers. 5. Waxes contain a long-chain fatty acid covalently linked to a long-chain alcohol. Waxes, which are water insoluble, are components of cell walls and are sometimes used as energy storage molecules. 6. Steroid lipids such as cholesterol help maintain the structural integrity of membranes as temperature fluctuates. 7. Carbohydrates such as monosaccharides, disaccharides, and polysaccharides serve as energy sources, structural molecules, and recognition sites during intercellular interactions. 8. Proteins are structural components of cells, enzymatic catalysts, regulators of various activities, molecules involved in the transportation of substances, and defensive molecules. They are composed of amino acids linked by peptide bonds, and they possess primary, secondary, tertiary, and (sometimes) quaternary structures that affect their function. Denaturation of a protein disrupts its structure and consequently its function. 9. Deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) are unbranched macromolecular polymers of nucleotides, each composed of either deoxyribose or ribose sugar, ionized phosphate, V e and a nitrogenous base. Five different bases exist: T adenine, guanine, cytosine, thymine, and uracil. DNA contains A, G, C, and T nucleotides. RNA uses U nucleotides instead of T nucleotides. Nucleosides are nucleotides lacking phosphate. o 4. Exchange reactions involve exchanging atoms between reactants. 4. Phospholipids contain two fatty acid chains and a phosphate functional group. The phospholipid head is hydrophilic, whereas the fatty acid portion of the molecule is hydrophobic. id 3. Decomposition reactions break larger molecules into smaller molecules and are exothermic reactions because they release energy. Hydrolysis is a decomposition reaction that uses water as one of the reactants. The sum of all decomposition reactions in an organism is called catabolism. If several double bonds exist in the fatty acids of a molecule of fat, it is a polyunsaturated fat. utor 2. Synthesis reactions form larger, more complex molecules. In dehydration synthesis, a molecule of water is removed from the reactants as the larger molecule is formed. Endothermic reactions require energy. Anabolism is the sum of all synthesis reactions in an organism. 2. Lipids, which include fats, phospholipids, waxes, and steroids, are hydrophobic (insoluble in water) macromolecules. 10. The structure of nucleic acids allows for genetic diversity, the correct copying of genes for their passage on to the next generation, and the accurate synthesis of proteins. 3. Fat (triglyceride) molecules are formed from a glycerol and three chainlike fatty acids. Saturated fatty acids contain more hydrogen in their structural formulas than unsaturated fatty acids, which contain double bonds between some carbon atoms. 11. Adenosine triphosphate (ATP), which is related to adenine nucleotide, is the most important short-term energy storage molecule in cells. It is also incorporated into the structure of many coenzymes. E E E QU STIONS FOR R VI W Answers to the Questions for Review (except Short Answer questions) begin on p. A-1. Multiple Choice 1. Which of the following structures have no electrical charge? a. protons c. neutrons b. electrons d. ions 2. The atomic mass of an atom most closely approximates the sum of the masses of all its ________. a. protons c. electrons b. isotopes d. protons and neutrons 3. One isotope of iodine differs from another in ________. a. the number of protons c. the number of neutrons b. the number of electrons d. atomic number M02_BAUM7206_05_SE_C02.indd 53 4. Which of the following is not an organic compound? a. monosaccharide c. water b. formaldehyde d. steroid 5. Which of the following terms most correctly describes the bonds in a molecule of water? a. nonpolar covalent bond c. ionic bond b. polar covalent bond d. hydrogen bond 6. In water, cations and anions of salts dissociate from one another and become surrounded by water molecules. In this state, the ions are also called ________. a. electrically negative c. electrolytes b. ionically bonded d. hydrogen bonds 12/1/16 3:26 PM e 2    The Chemistry of Microbiology iz e al Vi u s 7. Which of the following can be most accurately described as a decomposition reaction? a. C6H12O6 + 6 O2 S 6 H2O + 6 CO2 b. glucose + ATP S glucose phosphate + ADP c. 6 H2O + 6 CO2 S C6H12O6 + 6 O2 d. A + BC S AB + C It r pt       Cha ! 1. Label a portion of the molecule below; label two types of secondary structure; circle the tertiary structure. 8. Which of the following statements about a carbonated cola beverage with a pH of 2.9 is true? c. It has equal amounts of a. It has a relatively hydroxyl and hydrogen high concentration of ions. hydrogen ions. d. Cola is a buffered b. It has a relatively solution. low concentration of hydrogen ions. 9. Proteins are polymers of ________. a. amino acids c. nucleic acids b. fatty acids d. monosaccharides 10. Which of the following are hydrophobic organic molecules? a. proteins c. lipids b. carbohydrates d. nucleic acids 2. Shown is the amino acid tryptophan. Put the letter “C” at the site of every carbon atom. Label the amino group, the carboxyl group, and the side group. Fill in the Blanks 1. The outermost electron shell of an atom is known as the __________________ shell. O 2. The type of chemical bond between atoms with nearly equal electronegativities is called a(n) __________________ bond. OH 3. The principal short-term energy storage molecule in cells is __________________. HN 4. Common long-term storage molecules are __________________, __________________, __________________, and __________________. 5. Groups of atoms such as NH2 or OH that appear in certain common arrangements are called __________________. 6. The reverse of dehydration synthesis is __________________. 7. Reactions that release energy are called __________________ reactions. NH2 Short Answer 1. List three main types of chemical bonds, and give an example of each. 2. Name five properties of water that are vital to life. 8. All chemical reactions begin with reactants and result in new molecules called __________________. 3. Describe the difference(s) among saturated fatty acids, unsaturated fatty acids, and polyunsaturated fatty acids. 9. The __________________ scale is a measure of the concentration of hydrogen ions in a solution. 4. What is the difference between atomic oxygen and molecular oxygen? 10. A nucleic acid containing the base uracil would also contain __________________ sugar. M02_BAUM7206_05_SE_C02.indd 54 5. Explain how the polarity of water molecules makes water an excellent solvent. 12/1/16 3:26 PM       Questions for Review CRITICAL THINKING 1. Anthrax is caused by a bacterium, Bacillus anthracis, that avoids the body’s defenses against disease by synthesizing an outer glycoprotein covering made from d-glutamic acid. This covering is not digestible by white blood cells that normally engulf bacteria. Why is the covering indigestible? 2. Dehydrogenation is a chemical reaction in which a saturated fat is converted to an unsaturated fat. Explain why the name for this reaction is an appropriate one. 3. Two freshmen disagree about an aspect of chemistry. The nursing major insists that H + is the symbol for a hydrogen ion. The physics major insists that H + is the symbol for a proton. How can you help them resolve their disagreement? 4. When an egg white is heated, it changes from liquid to solid. When gelatin is cooled, it changes from liquid to solid. Both gelatin and egg white are proteins. From what you have learned about proteins, why can the gelatin be changed back to liquid but the cooked egg cannot? 5. When amino acids are synthesized in a test tube, d and l forms occur in equal amounts. However, cells use only l forms in their proteins. Occasionally, meteorites are found to contain amino acids. Based on these facts, how could NASA scientists determine whether the amino acids recovered from space are evidence of Earth-like extraterrestrial life rather than the result of nonmetabolic processes? 6. The poison glands of many bees and wasps contain acidic compounds. What common household chemical could be used to neutralize this poison? 7. Neon (atomic mass 10) and argon (atomic mass 18) are inert elements, which means that they very rarely form chemical bonds. M02_BAUM7206_05_SE_C02.indd 55 Give the electron configuration of their atoms, and explain why these elements are inert. 8. An article in the local newspaper about gangrene states that the tissue-destroying toxin, lecithinase, is an “organic compound.” But many people consider “organic” chemicals to mean something is good. Explain the apparent contradiction. 9. The deadly poison hydrogen cyanide has the chemical formula H ¬ C ‚ N. Describe the bonds between carbon and hydrogen and between carbon and nitrogen in terms of the number of electrons involved. 10. Triple covalent bonds are stronger and more difficult to break than single covalent bonds. Explain why by referring to the stability of a valence shell that contains eight electrons. 11. How can hydrogen bonding between water molecules help explain water’s ability to absorb large amounts of energy before evaporating? 12. How can a single molecule of magnesium hydroxide neutralize two molecules of hydrochloric acid? 13. We have seen that it is important that biological membranes remain flexible. Most bacteria lack sterols in their membranes and instead incorporate unsaturated phospholipids in the membranes to resist tight packing and solidification. Examine Table 2.4 on p. 42. Which fatty acid might best protect the membranes of an ice-dwelling bacterium? 14. Why is there no stereoisomer of glycine? 15. A textbook states that only five nucleotide bases are found in cells, but a laboratory worker reports that she has isolated eight different nucleotides. Explain why both are correct. 06/03/2018 09:00 M APP CONC EPT       Chapter 2   The Chemistry of Microbiology ING Using the following terms, fill in the following concept map that describes nucleic acids. RNA Ribosomal RNA Transfer RNA Nitrogenous bases (2) Phosphate Ribonucleotides Ribose DNA Double-stranded Guanine Messenger RNA Adenine Cytosine Deoxyribonucleotides Deoxyribose Nucleic acids two types in cells 1 9 is 3 types is 10 Single-stranded 2 4 is composed of is composed of 3 5 11 contain 3 parts contain 3 parts Adenine M02_BAUM7206_05_SE_C02.indd 56 8 12 6 Phosphate 14 7 13 4 types 4 types Cytosine Uracil 15 Guanine 16 Thymine 06/03/2018 10:55