BIOTECH Biochemistry Notes PDF

I believe that as the methods of structural chemistry are further applied to physiological problems, it will be found that the signiﬁcance of the hydrogen bond for physiology is greater than that of any other single structural feature. 2 —Linus Pauling, The Nature of the Chemical Bond, 1939 Water 2.1 Weak Interactions in Aqueous Systems 43 perature and favor the extreme ordering of molecules that is typical of crystalline water (ice). Polar biomole- 2.2 Ionization of Water,Weak Acids, and Weak Bases 54 cules dissolve readily in water because they can replace 2.3 Buffering against pH Changes in Biological Systems 59 water-water interactions with more energetically favor- able water-solute interactions. In contrast, nonpolar 2.4 Water as a Reactant 65 biomolecules interfere with water-water interactions but are unable to form water-solute interactions— 2.5 The Fitness of the Aqueous Environment for Living consequently, nonpolar molecules are poorly soluble in Organisms 65 water. In aqueous solutions, nonpolar molecules tend to cluster together. Hydrogen bonds and ionic, hydropho- bic (Greek, “water-fearing”), and van der Waals interac- W ater is the most abundant substance in living systems, making up 70% or more of the weight of most organisms. The ﬁrst living organisms on Earth doubtless arose in an aqueous environment, and the course of evolution has been shaped by the proper- tions are individually weak, but collectively they have a very significant influence on the three-dimensional structures of proteins, nucleic acids, polysaccharides, and membrane lipids. ties of the aqueous medium in which life began. This chapter begins with descriptions of the physical Hydrogen Bonding Gives Water Its Unusual Properties and chemical properties of water, to which all aspects of Water has a higher melting point, boiling point, and cell structure and function are adapted. The attractive heat of vaporization than most other common solvents forces between water molecules and the slight tendency of (Table 2–1). These unusual properties are a conse- water to ionize are of crucial importance to the structure quence of attractions between adjacent water molecules and function of biomolecules. We review the topic of ion- that give liquid water great internal cohesion. A look at ization in terms of equilibrium constants, pH, and titration the electron structure of the H2O molecule reveals the curves, and consider how aqueous solutions of weak acids cause of these intermolecular attractions. or bases and their salts act as buffers against pH changes in Each hydrogen atom of a water molecule shares an biological systems. The water molecule and its ionization electron pair with the central oxygen atom. The geome- products, H and OH, profoundly inﬂuence the structure, try of the molecule is dictated by the shapes of the outer self-assembly, and properties of all cellular components, in- electron orbitals of the oxygen atom, which are similar cluding proteins, nucleic acids, and lipids. The noncovalent to the sp3 bonding orbitals of carbon (see Fig. 1–14). interactions responsible for the strength and speciﬁcity of These orbitals describe a rough tetrahedron, with a hy- “recognition” among biomolecules are decisively inﬂu- drogen atom at each of two corners and unshared elec- enced by the solvent properties of water, including its abil- tron pairs at the other two corners (Fig. 2–1a). The ity to form hydrogen bonds with itself and with solutes. H—O—H bond angle is 104.5, slightly less than the 109.5 of a perfect tetrahedron because of crowding by the nonbonding orbitals of the oxygen atom. 2.1 Weak Interactions in Aqueous Systems The oxygen nucleus attracts electrons more strongly Hydrogen bonds between water molecules provide the than does the hydrogen nucleus (a proton); that is, oxy- cohesive forces that make water a liquid at room tem- gen is more electronegative. This means that the shared 43 44 Water TABLE 2–1 Melting Point, Boiling Point, and Heat of Vaporization of Some Common Solvents Melting point (C) Boiling point (C) Heat of vaporization (J/g)* Water 0 100 2,260 Methanol (CH3OH) 98 65 1,100 Ethanol (CH3CH2OH) 117 78 854 Propanol (CH3CH2CH2OH) 127 97 687 Butanol (CH3(CH2)2CH2OH) 90 117 590 Acetone (CH3COCH3) 95 56 523 Hexane (CH3(CH2)4CH3) 98 69 423 Benzene (C6H6) 6 80 394 Butane (CH3(CH2)2CH3) 135 0.5 381 Chloroform (CHCl3) 63 61 247 *The heat energy required to convert 1.0 g of a liquid at its boiling point and at atmospheric pressure into its gaseous state at the same temperature. It is a direct measure of the energy required to overcome attractive forces between molecules in the liquid phase. electrons are more often in the vicinity of the oxygen quired to break a bond) of about 23 kJ/mol, compared atom than of the hydrogen. The result of this unequal with 470 kJ/mol for the covalent O—H bond in water or electron sharing is two electric dipoles in the water mol- 348 kJ/mol for a covalent C—C bond. The hydrogen bond ecule, one along each of the H—O bonds; each hydrogen is about 10% covalent, due to overlaps in the bonding or- bears a partial positive charge (), and the oxygen bitals, and about 90% electrostatic. At room tempera- atom bears a partial negative charge equal in magnitude ture, the thermal energy of an aqueous solution (the to the sum of the two partial positives (2). As a result, kinetic energy of motion of the individual atoms and there is an electrostatic attraction between the oxygen molecules) is of the same order of magnitude as that re- atom of one water molecule and the hydrogen of another quired to break hydrogen bonds. When water is heated, (Fig. 2–1b), called a hydrogen bond. Throughout this the increase in temperature reﬂects the faster motion of book, we represent hydrogen bonds with three parallel individual water molecules. At any given time, most of blue lines, as in Figure 2–1b. the molecules in liquid water are hydrogen bonded, but Hydrogen bonds are relatively weak. Those in liquid the lifetime of each hydrogen bond is just 1 to 20 pi- water have a bond dissociation energy (the energy re- coseconds (1 ps 1012 s); when one hydrogen bond breaks, another hydrogen bond forms, with the same partner or a new one, within 0.1 ps. The apt phrase “ﬂick- 104.5 ering clusters” has been applied to the short-lived groups of water molecules interlinked by hydrogen bonds in liquid water. The sum of all the hydrogen bonds between H2O molecules confers great internal cohesion on liquid Hydrogen bond water. Extended networks of hydrogen-bonded water H 0.177 nm molecules also form bridges between solutes (proteins and nucleic acids, for example) that allow the larger molecules to interact with each other over distances of Covalent bond O 0.0965 nm several nanometers without physically touching. The nearly tetrahedral arrangement of the orbitals H about the oxygen atom (Fig. 2–1a) allows each water molecule to form hydrogen bonds with as many as four (a) (b) neighboring water molecules. In liquid water at room FIGURE 2–1 Structure of the water molecule. (a) The dipolar nature of temperature and atmospheric pressure, however, water the H2O molecule is shown in a ball-and-stick model; the dashed lines molecules are disorganized and in continuous motion, so represent the nonbonding orbitals. There is a nearly tetrahedral that each molecule forms hydrogen bonds with an aver- arrangement of the outer-shell electron pairs around the oxygen atom; age of only 3.4 other molecules. In ice, on the other hand, the two hydrogen atoms have localized partial positive charges () each water molecule is ﬁxed in space and forms hydro- and the oxygen atom has a partial negative charge (). (b) Two H2O gen bonds with a full complement of four other water molecules joined by a hydrogen bond (designated here, and through- molecules to yield a regular lattice structure (Fig. 2–2). out this book, by three blue lines) between the oxygen atom of the up- Breaking a sufﬁcient proportion of hydrogen bonds to per molecule and a hydrogen atom of the lower one. Hydrogen bonds destabilize the crystal lattice of ice requires much thermal are longer and weaker than covalent O—H bonds. energy, which accounts for the relatively high melting 2.1 Weak Interactions in Aqueous Systems 45 A A C E CJ Hydrogen E J J D D D D D J D acceptor O N O O O N Hydrogen H H H H H H O O O O O O O O O donor O O O N N N D G D G D G FIGURE 2–3 Common hydrogen bonds in biological systems. The hy- drogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom. to carbon atoms do not participate in hydrogen bond- ing, because carbon is only slightly more electronega- tive than hydrogen and thus the C—H bond is only very weakly polar. The distinction explains why butanol (CH3(CH2)2CH2OH) has a relatively high boiling point of 117 C, whereas butane (CH3(CH2)2CH3) has a boil- ing point of only 0.5 C. Butanol has a polar hydroxyl group and thus can form intermolecular hydrogen bonds. Uncharged but polar biomolecules such as sug- ars dissolve readily in water because of the stabilizing FIGURE 2–2 Hydrogen bonding in ice. In ice, each water molecule effect of hydrogen bonds between the hydroxyl groups forms four hydrogen bonds, the maximum possible for a water mole- or carbonyl oxygen of the sugar and the polar water cule, creating a regular crystal lattice. By contrast, in liquid water at molecules. Alcohols, aldehydes, ketones, and com- room temperature and atmospheric pressure, each water molecule pounds containing N—H bonds all form hydrogen hydrogen-bonds with an average of 3.4 other water molecules. This bonds with water molecules (Fig. 2–4) and tend to be crystal lattice structure makes ice less dense than liquid water, and thus soluble in water. ice ﬂoats on liquid water. Between the Between the Between peptide point of water (Table 2–1). When ice melts or water hydroxyl group carbonyl group groups in evaporates, heat is taken up by the system: of an alcohol of a ketone polypeptides and water and water H2O 1solid2 S H2O 1liquid2 ¢H 5.9 kJ/mol R R R2 G G A C H2O (liquid) S H2O (gas) ¢H 44.0 kJ/mol O H G D HH A 1ECJ NO C R J During melting or evaporation, the entropy of the H O O aqueous system increases as more highly ordered arrays of water molecules relax into the less orderly hydrogen- E OH H A H H H A bonded arrays in liquid water or into the wholly disor- EO HENH H C C dered gaseous state. At room temperature, both the A B melting of ice and the evaporation of water occur spon- R O taneously; the tendency of the water molecules to asso- ciate through hydrogen bonds is outweighed by the Between complementary energetic push toward randomness. Recall that the free- bases of DNA energy change (G) must have a negative value for a process to occur spontaneously: G H T S, H A where G represents the driving force, H the enthalpy R H E N ECH3 C change from making and breaking bonds, and S the N C A A Thymine change in randomness. Because H is positive for melt- ing and evaporation, it is clearly the increase in entropy KCH EC N O N O (S) that makes G negative and drives these changes. A H H A Water Forms Hydrogen Bonds with Polar Solutes H N H E N E NH C C Hydrogen bonds are not unique to water. They readily B A Adenine N C form between an electronegative atom (the hydrogen H K H C N acceptor, usually oxygen or nitrogen) and a hydrogen i l NOCH atom covalently bonded to another electronegative E R atom (the hydrogen donor) in the same or another mol- ecule (Fig. 2–3). Hydrogen atoms covalently bonded FIGURE 2–4 Some biologically important hydrogen bonds. 46 Water R trast, nonpolar solvents such as chloroform and benzene A R O A are poor solvents for polar biomolecules but easily dis- A O solve those that are hydrophobic—nonpolar molecules H Strong A H Weaker such as lipids and waxes. hydrogen bond G KO G KO hydrogen bond Water dissolves salts such as NaCl by hydrating and OP OP D D stabilizing the Na and Cl ions, weakening the electro- static interactions between them and thus counteract- FIGURE 2–5 Directionality of the hydrogen bond. The attraction be- ing their tendency to associate in a crystalline lattice tween the partial electric charges (see Fig. 2–1) is greatest when the three (Fig. 2–6). The same factors apply to charged biomole- atoms involved in the bond (in this case O, H, and O) lie in a straight line. cules, compounds with functional groups such as When the hydrogen-bonded moieties are structurally constrained (when ionized carboxylic acids (—COO), protonated amines they are parts of a single protein molecule, for example), this ideal geom- (—NH 3 ), and phosphate esters or anhydrides. Water etry may not be possible and the resulting hydrogen bond is weaker. readily dissolves such compounds by replacing solute- solute hydrogen bonds with solute-water hydrogen Hydrogen bonds are strongest when the bonded bonds, thus screening the electrostatic interactions be- molecules are oriented to maximize electrostatic inter- tween solute molecules. action, which occurs when the hydrogen atom and the Water is effective in screening the electrostatic in- two atoms that share it are in a straight line—that is, teractions between dissolved ions because it has a high when the acceptor atom is in line with the covalent bond dielectric constant, a physical property that reﬂects the between the donor atom and H (Fig. 2–5), putting the number of dipoles in a solvent. The strength, or force positive charge of the hydrogen ion directly between the (F ), of ionic interactions in a solution depends on the two partial negative charges. Hydrogen bonds are thus magnitude of the charges (Q), the distance between the highly directional and capable of holding two hydrogen- charged groups (r), and the dielectric constant ( , bonded molecules or groups in a specific geometric which is dimensionless) of the solvent in which the in- arrangement. As we shall see later, this property of teractions occur: hydrogen bonds confers very precise three-dimensional Q1Q2 structures on protein and nucleic acid molecules, which F er2 have many intramolecular hydrogen bonds. For water at 25 C, is 78.5, and for the very nonpolar solvent benzene, is 4.6. Thus, ionic interactions be- Water Interacts Electrostatically with Charged Solutes tween dissolved ions are much stronger in less polar en- Water is a polar solvent. It readily dissolves most bio- vironments. The dependence on r 2 is such that ionic molecules, which are generally charged or polar com- attractions or repulsions operate only over short pounds (Table 2–2); compounds that dissolve easily in distances—in the range of 10 to 40 nm (depending on water are hydrophilic (Greek, “water-loving”). In con- the electrolyte concentration) when the solvent is water. TABLE 2–2 Some Examples of Polar, Nonpolar, and Amphipathic Biomolecules (Shown as Ionic Forms at pH 7) Polar Nonpolar O Glucose CH2OH Typical wax CH3(CH2)7 CH CH (CH2)6 CH2 C O H OH O H CH3 (CH2)7 CH CH (CH2)7 CH2 OH H H HO H OH Amphipathic NH GNH3 Glycine 3 CH2 COO Phenylalanine CH2 CH COOJ Aspartate NH 3 OOC CH2 CH COO Phosphatidylcholine O Lactate CH3 CH COO CH3(CH2)15CH2 C O CH2 OH CH3(CH2)15CH2 C O CH O GN(CH3)3 O CH2 O P O CH2 CH2 Glycerol OH OJ HOCH2 CH CH2OH Polar groups Nonpolar groups 2.1 Weak Interactions in Aqueous Systems 47 FIGURE 2–6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The Hydrated NaCl crystal lattice is disrupted as water molecules H2O Cl– Cl– ion cluster about the Cl and Na ions. The ionic charges Cl– Na+ – are partially neutralized, and the electrostatic attrac- – tions necessary for lattice formation are weakened. + – Note the nonrandom – – orientation of the + + water molecules – – – – + + Na+ Hydrated – – Na+ ion – – Entropy Increases as Crystalline Substances Dissolve crease in entropy when they enter solution combine to make them very poorly soluble in water (Table 2–3). As a salt such as NaCl dissolves, the Na and Cl ions leaving the crystal lattice acquire far greater freedom of Some organisms have water-soluble “carrier proteins” motion (Fig. 2–6). The resulting increase in entropy (hemoglobin and myoglobin, for example) that facili- (randomness) of the system is largely responsible for tate the transport of O2. Carbon dioxide forms carbonic the ease of dissolving salts such as NaCl in water. In acid (H2CO3) in aqueous solution and is transported as thermodynamic terms, formation of the solution occurs the HCO 3 (bicarbonate) ion, either free—bicarbonate with a favorable free-energy change: G H T S, is very soluble in water (100 g/L at 25 C)—or bound where H has a small positive value and T S a large to hemoglobin. Three other gases, NH3, NO, and H2S, positive value; thus G is negative. also have biological roles in some organisms; these gases are polar, dissolve readily in water, and ionize in Nonpolar Gases Are Poorly Soluble in Water aqueous solution. The molecules of the biologically important gases CO2, Nonpolar Compounds Force Energetically Unfavorable O2, and N2 are nonpolar. In O2 and N2, electrons are Changes in the Structure of Water shared equally by both atoms. In CO2, each CUO bond is polar, but the two dipoles are oppositely directed and When water is mixed with benzene or hexane, two cancel each other (Table 2–3). The movement of mole- phases form; neither liquid is soluble in the other. Non- cules from the disordered gas phase into aqueous solu- polar compounds such as benzene and hexane are tion constrains their motion and the motion of water hydrophobic—they are unable to undergo energeti- molecules and therefore represents a decrease in en- cally favorable interactions with water molecules, and tropy. The nonpolar nature of these gases and the de- they interfere with the hydrogen bonding among water TABLE 2–3 Solubilities of Some Gases in Water Solubility Gas Structure* Polarity in water (g/L)† Nitrogen NqN Nonpolar 0.018 (40 C) Oxygen OPO Nonpolar 0.035 (50 C) Carbon dioxide Nonpolar 0.97 (45 C) OPCP O Ammonia H H Polar 900 (10 C) A H G D N Hydrogen sulﬁde H G D H Polar 1,860 (40 C) S *The arrows represent electric dipoles; there is a partial negative charge () at the head of the arrow, a partial positive charge (; not shown here) at the tail. † Note that polar molecules dissolve far better even at low temperatures than do nonpolar molecules at relatively high temperatures. 48 Water molecules. All molecules or ions in aqueous solution in- Amphipathic compounds contain regions that are terfere with the hydrogen bonding of some water mole- polar (or charged) and regions that are nonpolar cules in their immediate vicinity, but polar or charged (Table 2–2). When an amphipathic compound is mixed solutes (such as NaCl) compensate for lost water-water with water, the polar, hydrophilic region interacts favor- hydrogen bonds by forming new solute-water interac- ably with the solvent and tends to dissolve, but the non- tions. The net change in enthalpy (H) for dissolving polar, hydrophobic region tends to avoid contact with these solutes is generally small. Hydrophobic solutes, the water (Fig. 2–7a). The nonpolar regions of the however, offer no such compensation, and their addition molecules cluster together to present the smallest to water may therefore result in a small gain of enthalpy; hydrophobic area to the aqueous solvent, and the polar the breaking of hydrogen bonds between water mole- regions are arranged to maximize their interaction with cules takes up energy from the system, requiring the the solvent (Fig. 2–7b). These stable structures of am- input of energy from the surroundings. In addition to phipathic compounds in water, called micelles, may requiring this input of energy, dissolving hydrophobic contain hundreds or thousands of molecules. The forces compounds in water produces a measurable decrease in entropy. Water molecules in the immediate vicinity of a nonpolar solute are constrained in their possible orien- tations as they form a highly ordered cagelike shell around each solute molecule. These water molecules Dispersion of are not as highly oriented as those in clathrates, crys- lipids in H2O talline compounds of nonpolar solutes and water, but Each lipid the effect is the same in both cases: the ordering of water molecule forces molecules reduces entropy. The number of ordered surrounding H2O molecules to become water molecules, and therefore the magnitude of the en- highly ordered. tropy decrease, is proportional to the surface area of the hydrophobic solute enclosed within the cage of water molecules. The free-energy change for dissolving a non- polar solute in water is thus unfavorable: G H T S, where H has a positive value, S has a negative value, and G is positive. Hydrophilic O – “head group” O H C H O H C H Clusters of lipid molecules Only lipid portions at the edge of the cluster force the ordering of water. Fewer H2O molecules are ordered, and Hydrophobic entropy is increased. alkyl group “Flickering clusters” of H2O molecules in bulk phase Highly ordered H2O molecules form Micelles “cages” around the hydrophobic alkyl chains All hydrophobic groups are (a) sequestered from water; ordered FIGURE 2–7 Amphipathic compounds in aqueous solution. (a) Long- shell of H2O chain fatty acids have very hydrophobic alkyl chains, each of which is molecules is surrounded by a layer of highly ordered water molecules. (b) By clus- minimized, and entropy is further tering together in micelles, the fatty acid molecules expose the small- increased. est possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle. (b) 2.1 Weak Interactions in Aqueous Systems 49 that hold the nonpolar regions of the molecules together portant determinants of structure in biological mem- are called hydrophobic interactions. The strength of branes. Hydrophobic interactions between nonpolar hydrophobic interactions is not due to any intrinsic at- amino acids also stabilize the three-dimensional struc- traction between nonpolar moieties. Rather, it results tures of proteins. from the system’s achieving greatest thermodynamic Hydrogen bonding between water and polar solutes stability by minimizing the number of ordered water also causes an ordering of water molecules, but the en- molecules required to surround hydrophobic portions of ergetic effect is less significant than with nonpolar the solute molecules. solutes. Part of the driving force for binding of a polar Many biomolecules are amphipathic; proteins, pig- substrate (reactant) to the complementary polar sur- ments, certain vitamins, and the sterols and phospho- face of an enzyme is the entropy increase as the enzyme lipids of membranes all have both polar and nonpolar displaces ordered water from the substrate, and as the surface regions. Structures composed of these mole- substrate displaces ordered water from the enzyme sur- cules are stabilized by hydrophobic interactions among face (Fig. 2–8). the nonpolar regions. Hydrophobic interactions among lipids, and between lipids and proteins, are the most im- van der Waals Interactions Are Weak Interatomic Attractions Ordered water interacting with When two uncharged atoms are brought very close to- substrate and enzyme gether, their surrounding electron clouds inﬂuence each other. Random variations in the positions of the elec- trons around one nucleus may create a transient electric dipole, which induces a transient, opposite electric di- Substrate pole in the nearby atom. The two dipoles weakly attract each other, bringing the two nuclei closer. These weak attractions are called van der Waals interactions (also known as London forces). As the two nuclei Enzyme draw closer together, their electron clouds begin to repel each other. At the point where the net attraction is maximal, the nuclei are said to be in van der Waals con- tact. Each atom has a characteristic van der Waals radius, a measure of how close that atom will allow another to approach (Table 2–4). In the “space-ﬁlling” molecular models shown throughout this book, the atoms are depicted in sizes proportional to their van der Waals radii. van der Waals Radii and Covalent TABLE 2–4 Disordered water (Single-Bond) Radii of Some Elements displaced by enzyme-substrate van der Waals Covalent radius for interaction Element radius (nm) single bond (nm) H 0.11 0.030 O 0.15 0.066 N 0.15 0.070 C 0.17 0.077 S 0.18 0.104 P 0.19 0.110 I 0.21 0.133 Enzyme-substrate interaction stabilized by hydrogen-bonding, ionic, and hydrophobic interactions Sources: For van der Waals radii, Chauvin, R. (1992) Explicit periodic trend of van der Waals radii. J. Phys. Chem. 96, 9194–9197. For covalent radii, Pauling, L. (1960) FIGURE 2–8 Release of ordered water favors formation of an enzyme- Nature of the Chemical Bond, 3rd edn, Cornell University Press, Ithaca, NY. Note: van der Waals radii describe the space-ﬁlling dimensions of atoms. When two substrate complex. While separate, both enzyme and substrate force atoms are joined covalently, the atomic radii at the point of bonding are less than the neighboring water molecules into an ordered shell. Binding of sub- van der Waals radii, because the joined atoms are pulled together by the shared elec- strate to enzyme releases some of the ordered water, and the resulting tron pair. The distance between nuclei in a van der Waals interaction or a covalent bond is about equal to the sum of the van der Waals or covalent radii, respectively, for the two increase in entropy provides a thermodynamic push toward formation atoms. Thus the length of a carbon-carbon single bond is about 0.077 nm 0.077 nm of the enzyme-substrate complex (see p. 192). 0.154 nm. 50 Water Weak Interactions Are Crucial to Macromolecular one or more ionic interactions, as well as hydrophobic Structure and Function and van der Waals interactions. The formation of each of these weak bonds contributes to a net decrease in the The noncovalent interactions we have described— free energy of the system. We can calculate the stability hydrogen bonds and ionic, hydrophobic, and van der Waals of a noncovalent interaction, such as that of a small mol- interactions (Table 2–5)—are much weaker than covalent ecule hydrogen-bonded to its macromolecular partner, bonds. An input of about 350 kJ of energy is required to from the binding energy. Stability, as measured by the break a mole (6 1023) of C—C single bonds, and about equilibrium constant (see below) of the binding reac- 410 kJ to break a mole of C—H bonds, but as little as 4 kJ tion, varies exponentially with binding energy. The dis- is sufﬁcient to disrupt a mole of typical van der Waals in- sociation of two biomolecules (such as an enzyme and teractions. Hydrophobic interactions are also much weaker its bound substrate) that are associated noncovalently than covalent bonds, although they are substantially through multiple weak interactions requires all these in- strengthened by a highly polar solvent (a concentrated salt teractions to be disrupted at the same time. Because the solution, for example). Ionic interactions and hydrogen interactions ﬂuctuate randomly, such simultaneous dis- bonds are variable in strength, depending on the polarity of ruptions are very unlikely. The molecular stability be- the solvent and the alignment of the hydrogen-bonded stowed by 5 or 20 weak interactions is therefore much atoms, but they are always signiﬁcantly weaker than cova- greater than would be expected intuitively from a lent bonds. In aqueous solvent at 25 C, the available ther- simple summation of small binding energies. mal energy can be of the same order of magnitude as the Macromolecules such as proteins, DNA, and RNA strength of these weak interactions, and the interaction be- contain so many sites of potential hydrogen bonding or tween solute and solvent (water) molecules is nearly ionic, van der Waals, or hydrophobic interactions that the as favorable as solute-solute interactions. Consequently, cumulative effect of the many small binding forces can hydrogen bonds and ionic, hydrophobic, and van der Waals be enormous. For macromolecules, the most stable (that interactions are continually forming and breaking. is, the native) structure is usually that in which weak Although these four types of interactions are indi- interactions are maximized. The folding of a single vidually weak relative to covalent bonds, the cumulative polypeptide or polynucleotide chain into its three- effect of many such interactions can be very signiﬁcant. dimensional shape is determined by this principle. The For example, the noncovalent binding of an enzyme to binding of an antigen to a speciﬁc antibody depends on its substrate may involve several hydrogen bonds and the cumulative effects of many weak interactions. As noted earlier, the energy released when an enzyme binds noncovalently to its substrate is the main source of the enzyme’s catalytic power. The binding of a hormone or a Four Types of Noncovalent (“Weak”) neurotransmitter to its cellular receptor protein is the TABLE 2–5 Interactions among Biomolecules in result of multiple weak interactions. One consequence of Aqueous Solvent the large size of enzymes and receptors (relative to their Hydrogen bonds substrates or ligands) is that their extensive surfaces G provide many opportunities for weak interactions. At the Between neutral groups D CP O HO OO molecular level, the complementarity between interact- ing biomolecules reﬂects the complementarity and weak G interactions between polar, charged, and hydrophobic Between peptide bonds D CP G O HON groups on the surfaces of the molecules. D When the structure of a protein such as hemoglobin (Fig. 2–9) is determined by x-ray crystallography Ionic interactions O Attraction B (see Box 4–5, p. 132), water molecules are often found ONH3 O O CO to be bound so tightly that they are part of the crystal Repulsion structure; the same is true for water in crystals of RNA ONH3 H3N O or DNA. These bound water molecules, which can also be detected in aqueous solutions by nuclear magnetic resonance, have distinctly different properties from those of the “bulk” water of the solvent. They are, for ex- water Hydrophobic CH3 CH3 ample, not osmotically active (see below). For many G D proteins, tightly bound water molecules are essential to interactions CH A their function. In a reaction central to the process of CH2 photosynthesis, for example, light drives protons across A CH2 a biological membrane as electrons ﬂow through a series A of electron-carrying proteins (see Fig. 19–60). One of van der Waals interactions Any two atoms in these proteins, cytochrome f, has a chain of ﬁve bound close proximity water molecules (Fig. 2–10) that may provide a path for protons to move through the membrane by a process 2.1 Weak Interactions in Aqueous Systems 51 Solutes Affect the Colligative Properties of Aqueous Solutions Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic pressure. These are called colligative properties (colligative meaning “tied together”), because the ef- fect of solutes on all four properties has the same ba- (a) (b) sis: the concentration of water is lower in solutions FIGURE 2–9 Water binding in hemoglobin. (PDB ID 1A3N) The crystal than in pure water. The effect of solute concentration structure of hemoglobin, shown (a) with bound water molecules (red on the colligative properties of water is independent spheres) and (b) without the water molecules. The water molecules are of the chemical properties of the solute; it depends so ﬁrmly bound to the protein that they affect the x-ray diffraction pat- only on the number of solute particles (molecules, tern as though they were ﬁxed parts of the crystal. The two subunits ions) in a given amount of water. A compound such of hemoglobin are shown in gray, the two subunits in blue. Each sub- as NaCl, which dissociates in solution, has an effect unit has a bound heme group (red stick structure), visible only in the on osmotic pressure, for example, that is twice that subunits in this view. The structure and function of hemoglobin are of an equal number of moles of a nondissociating discussed in detail in Chapter 5. solute such as glucose. Water molecules tend to move from a region of higher water concentration to one of lower water con- known as “proton hopping” (described below). Another centration, in accordance with the tendency in nature such light-driven proton pump, bacteriorhodopsin, al- for a system to become disordered. When two different most certainly uses a chain of precisely oriented bound aqueous solutions are separated by a semipermeable water molecules in the transmembrane movement of membrane (one that allows the passage of water but not protons (see Fig. 19–67). solute molecules), water molecules diffusing from the region of higher water concentration to the region of lower water concentration produce osmotic pressure (Fig. 2–11). This pressure, , measured as the force Val60 Pro231 Gln59 H H O – O Heme N H propionate water O Force () H Pure Nonpermeant resists osmosis O H water solute dissolved N O Asn232 in water Asn168 H HN Piston O H Arg156 O h 153 Asn H HN O NH2 N H H Gln158 O N N Ala27 Fe H N H (a) (b) (c) HO C C Semipermeable H membrane O FIGURE 2–11 Osmosis and the measurement of osmotic pressure. FIGURE 2–10 Water chain in cytochrome f. Water is bound in a proton (a) The initial state. The tube contains an aqueous solution, the beaker channel of the membrane protein cytochrome f, which is part of contains pure water, and the semipermeable membrane allows the pas- the energy-trapping machinery of photosynthesis in chloroplasts (see sage of water but not solute. Water ﬂows from the beaker into the tube Fig. 19–64). Five water molecules are hydrogen-bonded to each other to equalize its concentration across the membrane. (b) The ﬁnal state. and to functional groups of the protein: the peptide backbone atoms of Water has moved into the solution of the nonpermeant compound, di- valine, proline, arginine, and alanine residues, and the side chains of luting it and raising the column of water within the tube. At equilibrium, three asparagine and two glutamine residues. The protein has a bound the force of gravity operating on the solution in the tube exactly balances heme (see Fig. 5–1), its iron ion facilitating electron ﬂow during pho- the tendency of water to move into the tube, where its concentration is tosynthesis. Electron flow is coupled to the movement of protons lower. (c) Osmotic pressure () is measured as the force that must be ap- across the membrane, which probably involves “proton hopping” (see plied to return the solution in the tube to the level of that in the beaker. Fig. 2–13) through this chain of bound water molecules. This force is proportional to the height, h, of the column in (b). 54 Water SUMMARY 2.1 Weak Interactions in reaction. We therefore turn now to a brief discussion of the ionization of water and of weak acids and bases dis- Aqueous Systems solved in water. The very different electronegativities of H and O make water a highly polar molecule, capable of Pure Water Is Slightly Ionized forming hydrogen bonds with itself and with solutes. Hydrogen bonds are ﬂeeting, primarily electrostatic, Water molecules have a slight tendency to undergo re- and weaker than covalent bonds. Water is a good versible ionization to yield a hydrogen ion (a proton) solvent for polar (hydrophilic) solutes, with which it and a hydroxide ion, giving the equilibrium forms hydrogen bonds, and for charged solutes, H2O Δ H OH (2–1) with which it interacts electrostatically. Although we commonly show the dissociation product Nonpolar (hydrophobic) compounds dissolve of water as H, free protons do not exist in solution; hy- poorly in water; they cannot hydrogen-bond with drogen ions formed in water are immediately hydrated the solvent, and their presence forces an to hydronium ions (H3O). Hydrogen bonding be- energetically unfavorable ordering of water tween water molecules makes the hydration of dissoci- molecules at their hydrophobic surfaces. To ating protons virtually instantaneous: minimize the surface exposed to water, nonpolar ÷ compounds such as lipids form aggregates H O H O H O H OH (micelles) in which the hydrophobic moieties are H H H sequestered in the interior, associating through hydrophobic interactions, and only the more polar The ionization of water can be measured by its elec- moieties interact with water. trical conductivity; pure water carries electrical current as H3O migrates toward the cathode and OH toward Weak, noncovalent interactions, in large numbers, the anode. The movement of hydronium and hydroxide decisively inﬂuence the folding of macromolecules ions in the electric ﬁeld is extremely fast compared with such as proteins and nucleic acids. The most stable that of other ions such as Na, K, and Cl. This high macromolecular conformations are those in which ionic mobility results from the kind of “proton hopping” hydrogen bonding is maximized within the molecule shown in Figure 2–13. No individual proton moves very and between the molecule and the solvent, and in far through the bulk solution, but a series of proton hops which hydrophobic moieties cluster in the interior of the molecule away from the aqueous solvent. The physical properties of aqueous solutions are Hydronium ion gives up a proton strongly inﬂuenced by the concentrations of H H Proton hop solutes. When two aqueous compartments are O+ separated by a semipermeable membrane (such as H the plasma membrane separating a cell from its O H H surroundings), water moves across that membrane H O H to equalize the osmolarity in the two compart- H O H ments. This tendency for water to move across a semipermeable membrane is the osmotic pressure. H O H 2.2 Ionization of Water,Weak Acids, and O H Weak Bases H H Although many of the solvent properties of water can be O explained in terms of the uncharged H2O molecule, the H small degree of ionization of water to hydrogen ions (H) and hydroxide ions (OH) must also be taken into O H account. Like all reversible reactions, the ionization of H water can be described by an equilibrium constant. Water accepts proton and When weak acids are dissolved in water, they contribute becomes a hydronium ion H by ionizing; weak bases consume H by becoming FIGURE 2–13 Proton hopping. Short “hops” of protons between a series protonated. These processes are also governed by equi- of hydrogen-bonded water molecules result in an extremely rapid net librium constants. The total hydrogen ion concentration movement of a proton over a long distance. As a hydronium ion (upper from all sources is experimentally measurable and is ex- left) gives up a proton, a water molecule some distance away (lower right) pressed as the pH of the solution. To predict the state of acquires one, becoming a hydronium ion. Proton hopping is much faster ionization of solutes in water, we must take into account than true diffusion and explains the remarkably high ionic mobility of H the relevant equilibrium constants for each ionization ions compared with other monovalent cations such as Na and K. 2.2 Ionization of Water, Weak Acids, and Weak Bases 55 between hydrogen-bonded water molecules causes the weight: (1,000 g/L)/(18.015 g/mol)—and is essentially net movement of a proton over a long distance in a re- constant in relation to the very low concentrations of H markably short time. As a result of the high ionic mo- and OH, namely, 1 107 M. Accordingly, we can sub- bility of H (and of OH, which also moves rapidly by stitute 55.5 M in the equilibrium constant expression proton hopping, but in the opposite direction), acid- (Eqn 2–3) to yield base reactions in aqueous solutions are exceptionally fast. As noted above, proton hopping very likely also [H ][OH ] Keq plays a role in biological proton-transfer reactions [55.5 M] (Fig. 2–10; see also Fig. 19–67). Because reversible ionization is crucial to the role of On rearranging, this becomes water in cellular function, we must have a means of ex- (55.5 M) (Keq) [H ] [OH ] Kw (2–4) pressing the extent of ionization of water in quantitative terms. A brief review of some properties of reversible where Kw designates the product (55.5 M)(Keq), the ion chemical reactions shows how this can be done. product of water at 25 C. The position of equilibrium of any chemical reaction The value for Keq, determined by electrical- is given by its equilibrium constant, Keq (sometimes conductivity measurements of pure water, is 1.8 expressed simply as K ). For the generalized reaction 1016 M at 25 C. Substituting this value for Keq in AB Δ CD (2–2) Equation 2–4 gives the value of the ion product of water: an equilibrium constant can be deﬁned in terms of the Kw [H ][OH ] (55.5 M) (1.8 1016 M) concentrations of reactants (A and B) and products (C and D) at equilibrium: 1.0 1014 M2 [C]eq[D]eq Thus the product [H][OH] in aqueous solutions at Keq 25 C always equals 1 1014 M2. When there are ex- [A]eq[B]eq actly equal concentrations of H and OH, as in pure Strictly speaking, the concentration terms should be the water, the solution is said to be at neutral pH. At this activities, or effective concentrations in nonideal solu- pH, the concentration of H and OH can be calculated tions, of each species. Except in very accurate work, from the ion product of water as follows: however, the equilibrium constant may be approximated by measuring the concentrations at equilibrium. For rea- Kw [H][OH] [H]2 [OH]2 sons beyond the scope of this discussion, equilibrium con- stants are dimensionless. Nonetheless, we have generally Solving for [H] gives retained the concentration units (M) in the equilibrium ex- pressions used in this book to remind you that molarity is [H] 2Kw 21 10 14 M2 the unit of concentration used in calculating Keq. [H] [OH ] 107 M The equilibrium constant is ﬁxed and characteristic for any given chemical reaction at a speciﬁed tempera- As the ion product of water is constant, whenever [H] ture. It deﬁnes the composition of the ﬁnal equilibrium is greater than 1 107 M, [OH] must be less than 1 mixture, regardless of the starting amounts of reactants 107 M, and vice versa. When [H] is very high, as in a and products. Conversely, we can calculate the equilib- solution of hydrochloric acid, [OH] must be very low. rium constant for a given reaction at a given tempera- From the ion product of water we can calculate [H] if ture if the equilibrium concentrations of all its reactants we know [OH], and vice versa. and products are known. As we showed in Chapter 1 (p. 24), the standard free-energy change (G) is di- rectly related to ln Keq. WORKED EXAMPLE 2–3 Calculation of [H] What is the concentration of H in a solution of 0.1 M The Ionization of Water Is Expressed by NaOH? an Equilibrium Constant Solution: We begin with the equation for the ion product The degree of ionization of water at equilibrium (Eqn 2–1) of water: is small; at 25 C only about two of every 109 molecules in Kw [H][OH] pure water are ionized at any instant. The equilibrium constant for the reversible ionization of water is With [OH] 0.1 M, solving for [H] gives [H ][OH ] Keq (2–3) Kw 1 1014 M2 1014 M2 [H2O] [H] [OH ] 0.1 M 101 M In pure water at 25 C, the concentration of water is 1013 M 55.5 M—grams of H2O in 1 L divided by its gram molecular JWCL281_c11_359-385.qxd 6/3/10 10:35 AM Page 359 Sugars and Polysaccharides CHAPTER 11 1 Monosaccharides weight. Polysaccharides such as starch in plants and glyco- A. Classification gen in animals serve as important nutritional reservoirs. B. Configurations and Conformations The elucidation of the structures and functions of carbo- C. Sugar Derivatives hydrates has lagged well behind those of proteins and nu- 2 Polysaccharides cleic acids. This can be attributed to several factors. Carbo- A. Carbohydrate Analysis hydrate compounds are often heterogeneous, both in size B. Disaccharides and in composition, which greatly complicates their physi- C. Structural Polysaccharides: Cellulose and Chitin cal and chemical characterization. They are not subject to D. Storage Polysaccharides: Starch and Glycogen the types of genetic analysis that have been invaluable in E. Glycosaminoglycans the study of proteins and nucleic acids because saccharide 3 Glycoproteins sequences are not genetically specified but are built up A. Proteoglycans through the sequential actions of specific enzymes (Section B. Bacterial Cell Walls 23-3B). Furthermore, it has been difficult to establish as- C. Glycoprotein Structure and Function says for the biological activities of polysaccharides because D. Glycomics of their largely passive roles. Nevertheless, it is abundantly clear that carbohydrates are essential elements in many, if not most, biological processes. Carbohydrates or saccharides (Greek: sakcharon, sugar) In this chapter, we explore the structures, chemistry, are essential components of all living organisms and are, in and, to a limited extent, the functions of carbohydrates, fact, the most abundant class of biological molecules. The alone and in association with proteins. Glycolipid struc- name carbohydrate, which literally means “carbon hy- tures are considered in Section 12-1D. The biosynthesis of drate,” stems from their chemical composition, which is complex carbohydrates is discussed in Section 23-3. roughly (C ⴢ H2O)n, where n ⱖ 3. The basic units of carbo- hydrates are known as monosaccharides. Many of these compounds are synthesized from simpler substances in a 1 MONOSACCHARIDES process named gluconeogenesis (Section 23-1). Others (and ultimately nearly all biological molecules) are the Monosaccharides or simple sugars are aldehyde or ketone products of photosynthesis (Section 24-3), the light-powered derivatives of straight-chain polyhydroxy alcohols contain- combination of CO2 and H2O through which plants and ing at least three carbon atoms. Such substances, for exam- certain bacteria form “carbon hydrates.” The metabolic ple, D-glucose and D-ribulose, cannot be hydrolyzed to breakdown of monosaccharides (Chapters 17 and 21) form simpler saccharides. provides much of the energy used to power biological O H processes. Monosaccharides are also principal components 1 C of nucleic acids (Section 5-1A), as well as important ele- 2 1 ments of complex lipids (Section 12-1D). H C OH CH2OH Oligosaccharides consist of a few covalently linked HO 3 C H 2 C O monosaccharide units. They are often associated with pro- 4 3 teins (glycoproteins) and lipids (glycolipids) in which they H C OH H C OH have both structural and regulatory functions (glycopro- 5 4 H C OH H C OH teins and glycolipids are collectively called glycoconju- 6 5 gates). Polysaccharides consist of many covalently linked CH2OH CH2OH monosaccharide units and have molecular masses ranging D-Glucose D-Ribulose well into the millions of daltons. They have indispensable structural functions in all types of organisms but are most In this section, the structures of the monosaccharides and conspicuous in plants because cellulose, their principal some of their biologically important derivatives are dis- structural material, comprises up to 80% of their dry cussed. 359 JWCL281_c11_359-385.qxd 6/3/10 10:35 AM Page 360 360 Chapter 11. Sugars and Polysaccharides A. Classification Fischer elucidated these configurations for the aldohexoses in 1896.According to the Fischer convention (Section 4-2B), D Monosaccharides are classified according to the chemical sugars have the same absolute configuration at the asymmet- nature of their carbonyl group and the number of their C ric center farthest removed from their carbonyl group as does atoms. If the carbonyl group is an aldehyde, as in glucose, D-glyceraldehyde. The L sugars, in accordance with this con- the sugar is an aldose. If the carbonyl group is a ketone, as vention, are mirror images of their D counterparts, as is in ribulose, the sugar is a ketose. The smallest monosaccha- shown below in Fischer projection for glucose. rides, those with three carbon atoms, are trioses. Those with four, five, six, seven, etc., C atoms are, respectively, tet- O H O H roses, pentoses, hexoses, heptoses, etc. These terms may be C C combined so that, for example, glucose is an aldohexose, whereas ribulose is a ketopentose. H C OH HO C H Examination of D-glucose’s molecular formula indicates HO C H H C OH that all but two of its six C atoms¬C1 and C6¬are chiral H C OH HO C H centers, so that D-glucose is one of 24 ⫽ 16 stereoisomers that comprise all possible aldohexoses. In general, n-carbon H C OH HO C H aldoses have 2n⫺2 stereoisomers. The stereochemistry and CH2OH CH2OH names of the D-aldoses are presented in Fig. 11-1. Emil D-Glucose L-Glucose CHO HCOH Aldotriose CH2OH D-Glyceraldehyde CHO CHO HCOH HOCH Aldotetroses HCOH HCOH CH2OH CH2OH D-Erythrose D-Threose CHO CHO CHO CHO HCOH HOCH HCOH HOCH HCOH HCOH HOCH HOCH Aldopentoses HCOH HCOH HCOH HCOH CH2OH CH2OH CH2OH CH2OH D-Ribose (Rib) D-Arabinose (Ara) D-Xylose (Xyl) D-Lyxose (Lyx) CHO CHO CHO CHO CHO CHO CHO CHO HCOH HOCH HCOH HOCH HCOH HOCH HCOH HOCH HCOH HCOH HOCH HOCH HCOH HCOH HOCH HOCH Aldohexoses HCOH HCOH HCOH HCOH HOCH HOCH HOCH HOCH HCOH HCOH HCOH HCOH HCOH HCOH HCOH HCOH CH2OH CH2OH CH2OH CH2OH CH2OH CH2OH CH2OH CH2OH D-Allose D-Altrose D-Glucose D-Mannose D-Gulose D-Idose D-Galactose D-Talose (Glc) (Man) (Gal) Figure 11-1 The stereochemical relationships, shown in distinguishes the members of each pair. The L- counterparts of Fischer projection, among the D-aldoses with three to six carbon these 15 sugars are their mirror images. The biologically most atoms. The arrows indicate stereochemical relationships (not common aldoses are boxed. biosynthetic reactions). The configuration about C2 (red) JWCL281_c11_359-385.qxd 6/3/10 10:35 AM Page 361 Section 11-1. Monosaccharides 361 Sugars that differ only by the configuration about one C function at C2 are the most common form (Fig. 11-2). Note atom are known as epimers of one another. Thus D-glucose that some of these ketoses are named by the insertion of and D-mannose are epimers with respect to C2, whereas -ul- before the suffix -ose in the name of the corresponding D-glucose and D-galactose are epimers with respect to C4 aldose; thus D-xylulose is the ketose corresponding to the (Fig. 11-1). However, D-mannose and D-galactose are not aldose D-xylose. Dihydroxyacetone, D-fructose, D-ribulose, epimers of each other because they differ in configuration and D-xylulose are the biologically most prominent about two of their C atoms. ketoses. D-Glucose is the only aldose that commonly occurs in na- ture as a monosaccharide. However, it and several other B. Configurations and Conformations monosaccharides including D-glyceraldehyde, D-ribose, D-mannose, and D-galactose are important components of Alcohols react with the carbonyl groups of aldehydes and larger biological molecules. L Sugars are biologically much ketones to form hemiacetals and hemiketals, respectively less abundant than D sugars. (Fig. 11-3). The hydroxyl and either the aldehyde or the ke- The position of their carbonyl group gives ketoses one tone functions of monosaccharides can likewise react in- less asymmetric center than their isomeric aldoses (e.g., tramolecularly to form cyclic hemiacetals and hemiketals compare D-fructose and D-glucose). n-Carbon ketoses (Fig. 11-4). The configurations of the substituents to each therefore have 2n⫺3 stereoisomers. Those with their ketone carbon atom of these sugar rings are conveniently repre- sented by their Haworth projection formulas. A sugar with a six-membered ring is known as a pyra- nose in analogy with pyran, the simplest compound con- taining such a ring. Similarly, sugars with five-membered CH2OH rings are designated furanoses in analogy with furan. C O O O CH2OH Dihydroxyacetone Pyran Furan CH2OH The cyclic forms of glucose and fructose with six- and five- C O membered rings are therefore known as glucopyranose and fructofuranose, respectively. HCOH CH2OH a. Cyclic Sugars Have Two Anomeric Forms D-Erythrulose The Greek letters preceding the names in Fig. 11-4 still need to be explained. The cyclization of a monosaccharide renders the former carbonyl carbon asymmetric. The re- CH2OH CH2OH sulting pair of diastereomers are known as anomers and the hemiacetal or hemiketal carbon is referred to as the C O C O anomeric carbon. In the ␣ anomer, the OH substituent to HCOH HOCH the anomeric carbon is on the

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