《生物化学原理》英文版 chapter 2 WATER

8885dc02_47-747/25/0310:05 AM Page49mac76mac76:385 Chapter 2 Water water or 348 kJ/mol for a covalent C-C bond. The hy- drogen bond is about 10% covalent, due to overlaps in the bonding orbitals, and about 90% electrostatic. At 03 room temperature, the thermal energy of an aqueous solution(the kinetic energy of motion of the individual atoms and molecules) is of the same order of magnitude as that required to break hydrogen bonds. When water is heated, the increase in temperature reflects the faster motion of individual water molecules. At any given time, most of the molecules in liquid water are engaged in hy- drogen bonding, but the lifetime of each hydrogen bond is just l to 20 picoseconds(l ps=10-S); upon break- age of one hydrogen bond, another hydrogen bond forms, with the same partner or a new one, within 0. 1 ps The apt phrase"flickering clusters"has been applied to the short-lived groups of water molecules interlinked by hydrogen bonds in liquid water. The sum of all the hy- drogen bonds between H,O molecules confers great in- ternal cohesion on liquid water. Extended networks of hydrogen-bonded water molecules also form bridges be tween solutes (proteins and nucleic acids, for example) that allow the larger molecules to interact with each FIGURE 2-2 Hydrogen bonding in ice. In ice, each water molecule other over distances of several nanometers without physically touching forms the maximum of four hydrogen bonds, creating a regular crys- tal lattice. By contrast, in liquid water at room temperature and at- The nearly tetrahedral arrangement of the orbitals mospheric pressure, each water molecule hydrogen-bonds with molecule to form hydrogen bonds with as many as four less dense than liquid water, and thus ice floats on liquid water i about the oxygen atom(Fig. 2-la) allows each water erage of 3.4 other water molecules. This crystal lattice of ice make neighboring water molecules. In liquid water at room temperature and atmospheric pressure, however, water molecules are disorganized and in continuous motion, and breaking bonds, and AS the change in randomness so that each molecule forms hydrogen bonds with an av- Because AH is positive for melting and evaporation, it erage of only 3.4 other molecules In ice, on the other is clearly the increase in entropy(As)that makes AG hand, each water molecule is fixed in space and forms negative and drives these transformations. hydrogen bonds with a full complement of four other water molecules to yield a regular lattice structure (fig Water Forms Hydrogen Bonds with Polar Solutes 2-2). Breaking a sufficient proportion of hydrogen bonds to destabilize the crystal lattice of ice requires Hydrogen bonds are not unique to water. They readily much thermal energy, which accounts for the relatively form between an electronegative atom(the hydrogen high melting point of water (Table 2-1). When ice melts acceptor, usually oxygen or nitrogen with a lone pair of or water evaporates, heat is taken up by the system electrons) and a hydrogen atom covalently bonded to another electronegative atom(the hydrogen donor) in H2O( solid)→→H2 Aliquid 1=+5.9 k/mol the same or another molecule(Fig. 2-3). Hydrogen atoms covalently bonded to carbon atoms do not par H2 O(liquid)→→H2O(gas) H=+44.0 k/mol ticipate in hydrogen bonding, because carbon is only During melting or evaporation, the entropy of the aqueous system increases as more highly ordered arrays of water molecules relax into the less orderly hydrogen bonded arrays in liquid water or the wholly disordered Hydrogen gaseous state At room temperature, both the melting of ice and the evaporation of water occur spontaneously HHH HH the tendency of the water molecules to associate through ydrogen bonds is outweighed by the energetic push toward randomness. Recall that the free-energy change (AG) must have a negative value for a process to occur 2-3 Common hydrogen bonds in biological systems. The spontaneously: AG= AH-TAS, where AG represents hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor the driving force, AH the enthalpy change from making is another electronegative atom

water or 348 kJ/mol for a covalent COC bond. The hy￾drogen bond is about 10% covalent, due to overlaps in the bonding orbitals, and about 90% electrostatic. At room temperature, the thermal energy of an aqueous solution (the kinetic energy of motion of the individual atoms and molecules) is of the same order of magnitude as that required to break hydrogen bonds. When water is heated, the increase in temperature reflects the faster motion of individual water molecules. At any given time, most of the molecules in liquid water are engaged in hy￾drogen bonding, but the lifetime of each hydrogen bond is just 1 to 20 picoseconds (1 ps  1012 s); upon break￾age of one hydrogen bond, another hydrogen bond forms, with the same partner or a new one, within 0.1 ps. The apt phrase “flickering clusters” has been applied to the short-lived groups of water molecules interlinked by hydrogen bonds in liquid water. The sum of all the hy￾drogen bonds between H2O molecules confers great in￾ternal cohesion on liquid water. Extended networks of hydrogen-bonded water molecules also form bridges be￾tween solutes (proteins and nucleic acids, for example) that allow the larger molecules to interact with each other over distances of several nanometers without physically touching. The nearly tetrahedral arrangement of the orbitals about the oxygen atom (Fig. 2–1a) allows each water molecule to form hydrogen bonds with as many as four neighboring water molecules. In liquid water at room temperature and atmospheric pressure, however, water molecules are disorganized and in continuous motion, so that each molecule forms hydrogen bonds with an av￾erage of only 3.4 other molecules. In ice, on the other hand, each water molecule is fixed in space and forms hydrogen bonds with a full complement of four other water molecules to yield a regular lattice structure (Fig. 2–2). Breaking a sufficient proportion of hydrogen bonds to destabilize the crystal lattice of ice requires much thermal energy, which accounts for the relatively high melting point of water (Table 2–1). When ice melts or water evaporates, heat is taken up by the system: H2O(solid) 88n H2O(liquid) H 
5.9 kJ/mol H2O(liquid) 88n H2O(gas) H 
44.0 kJ/mol During melting or evaporation, the entropy of the aqueous system increases as more highly ordered arrays of water molecules relax into the less orderly hydrogen￾bonded arrays in liquid water or the wholly disordered gaseous state. At room temperature, both the melting of ice and the evaporation of water occur spontaneously; the tendency of the water molecules to associate through hydrogen bonds is outweighed by the energetic push toward randomness. Recall that the free-energy change (G) must have a negative value for a process to occur spontaneously: G  H T S, where G represents the driving force, H the enthalpy change from making and breaking bonds, and S the change in randomness. Because H is positive for melting and evaporation, it is clearly the increase in entropy (S) that makes G negative and drives these transformations. Water Forms Hydrogen Bonds with Polar Solutes Hydrogen bonds are not unique to water. They readily form between an electronegative atom (the hydrogen acceptor, usually oxygen or nitrogen with a lone pair of electrons) and a hydrogen atom covalently bonded to another electronegative atom (the hydrogen donor) in the same or another molecule (Fig. 2–3). Hydrogen atoms covalently bonded to carbon atoms do not par￾ticipate in hydrogen bonding, because carbon is only Chapter 2 Water 49 FIGURE 2–2 Hydrogen bonding in ice. In ice, each water molecule forms the maximum of four hydrogen bonds, creating a regular crys￾tal lattice. By contrast, in liquid water at room temperature and at￾mospheric pressure, each water molecule hydrogen-bonds with an av￾erage of 3.4 other water molecules. This crystal lattice of ice makes it less dense than liquid water, and thus ice floats on liquid water. Hydrogen Hydrogen donor acceptor H O O P C D G DO O J H N OO O J D H N NO O DD H O OO O H O N P C G DO O DD H N OO O FIGURE 2–3 Common hydrogen bonds in biological systems. The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom. 8885d_c02_47-74 7/25/03 10:05 AM Page 49 mac76 mac76:385_reb:

8885dc02_47-747/25/0310:05 AM Page50mac76mac76:385 Part I Structure and Catalysis lightly more electronegative than hydrogen and thus the C-H bond is only very weakly polar. The distinc tion explains why butanol(CH3(CH2)2.OH) has a rel- atively high boiling point of 117C, whereas butane R|o|H=O Weaker (CH3(CH2)2CH3) has a boiling point of only -0. C Bu hydrogen bond tanol has a polar hydroxyl group and thus can form in termolecular hydrogen bonds. Uncharged but polar bio- molecules such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds be FIGURE 2-5 Directionality of the hydrogen bond. The attraction be tween the hydroxyl groups or carbonyl oxygen of the tween the partial electric charges(see Fig. 2-1)is greatest when the three atoms involved (in this case O, H, and O) lie in a straight line sugar and the polar water molecules. Alcohols, alde When the hydrogen-bonded moieties are structurally constrained (as hydes, ketones, and compounds containing N-H bonds when they are parts of a single protein molecule, for example),this all form hydrogen bonds with water molecules(Fig 2-4) ideal geometry may not be possible and the resulting hydrogen bond and tend to be soluble in water Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic inter- action, which occurs when the hydrogen atom and the ing two hydrogen-bonded molecules or groups in a spe- two atoms that share it are in a straight line-that is, cific geometric arrangement. As we shall see later, this when the acceptor atom is in line with the covalent bond property of hydrogen bonds confers very precise three between the donor atom and H(Fig. 2-5). Hydrogen dimensional structures on protein and nucleic acid bonds are thus highly directional and capable of hold bond Water Interacts Electrostatically Between the Between the Between peptid with Charged Solutes hydroxyl group carbonyl group Water is a polar solvent. It readily dissolves most bio- and water and water molecules, which are generally charged or polar com- pounds (Table 2-2); compounds that dissolve easily in water are hydrophilie (Greek, "water-loving"). In con- trast, nonpolar solvents such as chloroform and benzene H are poor solvents for polar biomolecules but easily dis O H solve those that are hydrophobic--nonpolar molecules H such as lipids and waxes Water dissolves salts such as Nacl by hydrating and labilizing the Na and cl ions, weakening the elec- trostatic interactions between them and thus counter acting their tendency to associate in a crystalline lattice Between (Fig 2-6). The same factors apply to charged biomole- lementary bases of dna cules, compounds with functional groups such as ion- ized carboxylic acids (Co0), protonated amines (NH3), and phosphate esters or anhydrides. Water readily dissolves such compounds by replacing solute- solute hydrogen bonds with solute-water hydrogen bonds, thus screening the electrostatic interactions be- tween solute molecules Water is especially effective in screening the elec- trostatic interactions between dissolved ions because it H has a high dielectric constant, a physical property re- flecting the number of dipoles in a solvent. The strength Adenine or force(F), of ionic interactions in a solution depends upon the magnitude of the charges (@), the distance between the charged groups (r), and the dielectric con- stant(8) of the solvent in which the interactions occur FIGURE 2-4 Some biologically important hydrogen bonds

slightly more electronegative than hydrogen and thus the COH bond is only very weakly polar. The distinc￾tion explains why butanol (CH3(CH2)2CH2OH) has a rel￾atively high boiling point of 117 C, whereas butane (CH3(CH2)2CH3) has a boiling point of only 0.5 C. Bu￾tanol has a polar hydroxyl group and thus can form in￾termolecular hydrogen bonds. Uncharged but polar bio￾molecules such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds be￾tween the hydroxyl groups or carbonyl oxygen of the sugar and the polar water molecules. Alcohols, alde￾hydes, ketones, and compounds containing NOH bonds all form hydrogen bonds with water molecules (Fig. 2–4) and tend to be soluble in water. Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic inter￾action, which occurs when the hydrogen atom and the two atoms that share it are in a straight line—that is, when the acceptor atom is in line with the covalent bond between the donor atom and H (Fig. 2–5). Hydrogen bonds are thus highly directional and capable of hold￾ing two hydrogen-bonded molecules or groups in a spe￾cific geometric arrangement. As we shall see later, this property of hydrogen bonds confers very precise three￾dimensional structures on protein and nucleic acid molecules, which have many intramolecular hydrogen bonds. Water Interacts Electrostatically with Charged Solutes Water is a polar solvent. It readily dissolves most bio￾molecules, which are generally charged or polar com￾pounds (Table 2–2); compounds that dissolve easily in water are hydrophilic (Greek, “water-loving”). In con￾trast, nonpolar solvents such as chloroform and benzene are poor solvents for polar biomolecules but easily dis￾solve those that are hydrophobic—nonpolar molecules such as lipids and waxes. Water dissolves salts such as NaCl by hydrating and stabilizing the Na
and Cl ions, weakening the elec￾trostatic interactions between them and thus counter￾acting their tendency to associate in a crystalline lattice (Fig. 2–6). The same factors apply to charged biomole￾cules, compounds with functional groups such as ion￾ized carboxylic acids (OCOO), protonated amines (ONH3
), and phosphate esters or anhydrides. Water readily dissolves such compounds by replacing solute￾solute hydrogen bonds with solute-water hydrogen bonds, thus screening the electrostatic interactions be￾tween solute molecules. Water is especially effective in screening the elec￾trostatic interactions between dissolved ions because it has a high dielectric constant, a physical property re￾flecting the number of dipoles in a solvent. The strength, or force (F), of ionic interactions in a solution depends upon the magnitude of the charges (Q), the distance between the charged groups (r), and the dielectric con￾stant () of the solvent in which the interactions occur: F   Q  1 r Q 2 2  50 Part I Structure and Catalysis Between the hydroxyl group of an alcohol and water Between the carbonyl group of a ketone and water Between peptide groups in polypeptides Between complementary bases of DNA O H A O G R H H H O G R1 O R O E A H O B HN A H H B O H N C C EC A R H C HN A H EC A NO H A A N HR EC H NCECH3 HC K HN H ENN ENH N ER OCH D R2 A C KO N B C C A C l A Thymine Adenine B C i A H H H E H FIGURE 2–4 Some biologically important hydrogen bonds. Strong hydrogen bond Weaker hydrogen bond P KO H A O A R P KO H A O A R G D G O D O FIGURE 2–5 Directionality of the hydrogen bond. The attraction be￾tween the partial electric charges (see Fig. 2–1) is greatest when the three atoms involved (in this case O, H, and O) lie in a straight line. When the hydrogen-bonded moieties are structurally constrained (as when they are parts of a single protein molecule, for example), this ideal geometry may not be possible and the resulting hydrogen bond is weaker. 8885d_c02_47-74 7/25/03 10:05 AM Page 50 mac76 mac76:385_reb:

8885dc02_0517/25/0311:52 AM Page51mac76mac76:385reb: Chapter 2 Water TABLE 2-2 Some Examples of Polar, Nonpolar, and Amphipathic Biomolecules(Shown as lonic Forms at pH 7) Polar Glucose CHO CH3(CH2)-CH-CH-(CH2)-CH2 O CH3(CH2)--CH-CH-(CHa)7-CH Glycine+NH3-CH2-COO NH3 CHo-CH-CO0- OoC—CH2-CH-CO0 Phosphatidylcholine Lactate CH3-CH-CO0 CH3(CH2)15CH2-C-O-CHa CH3(CHg)sCH, 0-CH N(CH3)3 CH2-0-P--0-CH2-CH2 HOCH2-CH-CH2OH Polar groups Nonpolar groups For water at 25C, 8(which is dimensionless) is 78.5 and for the very nonpolar solvent benzene, a is 4.6. Thus Entropy Increases as Crystalline Substances Dissolve ionic interactions are much stronger in less polar As a salt such as Nacl dissolves, the Na and Cl ions ronments. The dependence on r- is such that ionic leaving the crystal lattice acquire far greater freedom of tractions or repulsions operate only over short motion(Fig. 2-6). The resulting increase in entropy ances--in the range of 10 to 40 nm(depending on the (randomness) of the system is largely responsible for electrolyte concentration) when the solvent is water. the ease of dissolving salts such as Nacl in water. In Hydrate Na+ion 8a ob oo d Na of the water molecules q=)8 E 2-6 Water as solvent Water dissolves many crystalline salts charges are partially neutralized, and the electrostatic attractions nec- by hydrating their component ions. The NaCl crystal lattice is disrupted essary for lattice formation are weakened as water molecules cluster about the Cl and Na ions the ion

For water at 25
C,
(which is dimensionless) is 78.5, and for the very nonpolar solvent benzene,
is 4.6. Thus, ionic interactions are much stronger in less polar envi￾ronments. The dependence on r2 is such that ionic at￾tractions or repulsions operate only over short dis￾tances—in the range of 10 to 40 nm (depending on the electrolyte concentration) when the solvent is water. Entropy Increases as Crystalline Substances Dissolve As a salt such as NaCl dissolves, the Na and Cl ions leaving the crystal lattice acquire far greater freedom of motion (Fig. 2–6). The resulting increase in entropy (randomness) of the system is largely responsible for the ease of dissolving salts such as NaCl in water. In Chapter 2 Water 51 TABLE 2–2 Some Examples of Polar, Nonpolar, and Amphipathic Biomolecules (Shown as Ionic Forms at pH 7) + Hydrated Na+ ion Note the orientation of the water molecules Hydrated Cl– ion H2O Na+ Cl– + – + – + – – – – + + + + – – – – – – – – – + – – FIGURE 2–6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the Cl and Na ions. The ionic charges are partially neutralized, and the electrostatic attractions nec￾essary for lattice formation are weakened. H HO CH2OH O OH OH OH CH2 NH3 COO CH2 OOC COO H H H H NH3 CH CH OH OH CH3 CH COO HOCH2 CH2OH CH3(CH2)7 CH CH (CH2)6 CH2 C CH3(CH2)7 CH CH (CH2)7 CH2 CH2 CH GNH3 GN(CH3)3 O O COOJ CH3(CH2)15CH2 CH2 O CH2 CH2 O OJ C CH3(CH2)15CH2 O CH O O CH2 P C O O Polar groups Nonpolar groups Polar Glucose Glycine Aspartate Lactate Glycerol Nonpolar Typical wax Amphipathic Phenylalanine Phosphatidylcholine 8885d_c02_051 7/25/03 11:52 AM Page 51 mac76 mac76:385_reb:

8885dc02_47-747/25/0310:05 AM Page52mac76mac76:385 Part I Structure and Catalysis thermodynamic terms, formation of the solution occurs hydrophobic--they are unable to undergo energetically with a favorable free-energy change: AG=AH -TAS, favorable interactions with water molecules, and they where AH has a small positive value and T'As a large interfere with the hydrogen bonding among water mol positive value; thus AG is negative ecules. All molecules or ions in aqueous solution inter fere with the hydrogen bonding of some water mole Nonpolar Gases Are Poorly Soluble in Water cules in their immediate vicinity, but polar or charged solutes(such as Nacl) compensate for lost water-water The molecules of the biologically important gases CO2, hydrogen bonds by forming new solute-water interac O2, and N2 are nonpolar In O2 and N2, electrons are tions. The net change in enthalpy(AH) for dissolving shared equally by both atoms In CO2, each C-0 bond these solutes is generally small Hydrophobic solutes is polar, but the two dipoles are oppositely directed and however, offer no such compensation, and their addi cancel each other(Table 2-3). The movement of mole- tion to water may therefore result in a small gain of en- cules from the disordered gas phase into aqueous solu- thalpy; the breaking of hydrogen bonds between water tion constrains their motion and the motion of water molecules takes up energy from the system. Further molecules and therefore represents a decrease in en- more, dissolving hydrophobic compounds in water pro tropy. The nonpolar nature of these gases and the de duces a measurable decrease in entropy. Water mole- crease in entropy when they enter solution combine to cules in the immediate vicinity of a nonpolar solute are make them very poorly soluble in water (Table 2-3). constrained in their possible orientations as they form Some organisms have water-soluble carrier proteins a highly ordered cagelike shell around each solute mol- (hemoglobin and myoglobin, for example) that facilitate ecule. These water molecules are not as highly oriented nsport Carbon dioxide forms carbonic acid as those in clathrates, crystalline compounds of non (H2CO3) in aqueous solution and is transported as the polar solutes and water, but the effect is the same in HCO3 (bicarbonate)ion, either free--bicarbonate is both cases: the ordering of water molecules reduces en- very soluble in water(-100 g/L at 25C)-or bound to tropy. The number of ordered water molecules, and hemoglobin. Two other gases, NHa and H2s, also have therefore the magnitude of the entropy decrease, is pro- ological roles in some organisms; these gases are po- portional to the surface area of the hydrophobic solute r and dissolve readily in water. enclosed within the cage of water molecules. The free- energy change for dissolving a nonpolar solute in water Nonpolar Compounds Force Energetically Unfavorable is thus unfavorable: AG-AH-TAS, where AH has Changes in the Structure of Water a positive value, AS has a negative value, and AG is positive When water ed with benzene or hexane. two Amphipathic compounds contain regions that are phases form; neither liquid is soluble in the other Non- polar (or charged) and regions that are nonpolar ( table polar compounds such as benzene and hexane are 2-2). When an amphipathic compound is mixed with TABlE 2-3 Solubilities of some gases in Water Solubility in water(g/L) N≡三N 0.018(40°C) Carbon dioxide 0.97(45°C) O=C=0 0(10°C) Hydrogen sulfide 1860(40°C) he arrows represent electric dipoles there is a partial negative charge(8 )at the head of the arrow, a partial positive charge (6: not shown here)at the tail. TNote that polar molecules dissolve far better even at low temperatures than do nonpolar molecules at relatively high temperatures

thermodynamic terms, formation of the solution occurs with a favorable free-energy change: G  H T S, where H has a small positive value and T S a large positive value; thus G is negative. Nonpolar Gases Are Poorly Soluble in Water The molecules of the biologically important gases CO2, O2, and N2 are nonpolar. In O2 and N2, electrons are shared equally by both atoms. In CO2, each CUO bond is polar, but the two dipoles are oppositely directed and cancel each other (Table 2–3). The movement of mole￾cules from the disordered gas phase into aqueous solu￾tion constrains their motion and the motion of water molecules and therefore represents a decrease in en￾tropy. The nonpolar nature of these gases and the de￾crease in entropy when they enter solution combine to make them very poorly soluble in water (Table 2–3). Some organisms have water-soluble carrier proteins (hemoglobin and myoglobin, for example) that facilitate the transport of O2. Carbon dioxide forms carbonic acid (H2CO3) in aqueous solution and is transported as the HCO3 (bicarbonate) ion, either free—bicarbonate is very soluble in water (~100 g/L at 25 C)—or bound to hemoglobin. Two other gases, NH3 and H2S, also have biological roles in some organisms; these gases are po￾lar and dissolve readily in water. Nonpolar Compounds Force Energetically Unfavorable Changes in the Structure of Water When water is mixed with benzene or hexane, two phases form; neither liquid is soluble in the other. Non￾polar compounds such as benzene and hexane are hydrophobic—they are unable to undergo energetically favorable interactions with water molecules, and they interfere with the hydrogen bonding among water mol￾ecules. All molecules or ions in aqueous solution inter￾fere with the hydrogen bonding of some water mole￾cules in their immediate vicinity, but polar or charged solutes (such as NaCl) compensate for lost water-water hydrogen bonds by forming new solute-water interac￾tions. The net change in enthalpy (H) for dissolving these solutes is generally small. Hydrophobic solutes, however, offer no such compensation, and their addi￾tion to water may therefore result in a small gain of en￾thalpy; the breaking of hydrogen bonds between water molecules takes up energy from the system. Further￾more, dissolving hydrophobic compounds in water pro￾duces a measurable decrease in entropy. Water mole￾cules in the immediate vicinity of a nonpolar solute are constrained in their possible orientations as they form a highly ordered cagelike shell around each solute mol￾ecule. These water molecules are not as highly oriented as those in clathrates, crystalline compounds of non￾polar solutes and water, but the effect is the same in both cases: the ordering of water molecules reduces en￾tropy. The number of ordered water molecules, and therefore the magnitude of the entropy decrease, is pro￾portional to the surface area of the hydrophobic solute enclosed within the cage of water molecules. The free￾energy change for dissolving a nonpolar 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. Amphipathic compounds contain regions that are polar (or charged) and regions that are nonpolar (Table 2–2). When an amphipathic compound is mixed with 52 Part I Structure and Catalysis TABLE 2–3 Solubilities of Some Gases in Water Solubility Gas Structure* Polarity in water (g/L)† Nitrogen NmN Nonpolar 0.018 (40 °C) Oxygen OPO Nonpolar 0.035 (50 °C) Carbon dioxide Nonpolar 0.97 (45 °C) Ammonia Polar 900 (10 °C) Hydrogen sulfide Polar 1,860 (40 °C) H G S D H  H GN A H D H  OPCPO   *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. 8885d_c02_47-74 7/25/03 10:05 AM Page 52 mac76 mac76:385_reb:

885c0247-747/25/0310:05 AM Page53mac76mac76:385e Chapter 2 Water Hydro 8° g pids in H2° molecule forces molecules to become Hydrophobic g "Flickering clusters"of H2O Highly ordered HoO molecules form cages"around the hydrophobic alkyl chains (a) .@@ Clusters of lipid molecules FIGURE 2-7 Amphipathic compounds in aqueous solution (a)Long. Only lipid portion chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules. (b)By ° cluster force the ordering of water clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water,and fewer 06. are ordered, and water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle. water, the polar, hydrophilic region interacts favorably Micelles with the solvent and tends to dissolve, but the nono- All hydrophobic lar, hydrophobic region tends to avoid contact with the groups are water(Fig. 2-7a). The nonpolar regions of the mole- sequestered from water: ordered cules cluster together to present the smallest hy- shell of H2O drophobic area to the aqueous solvent, and the polarre- gions are arranged to maximize their interaction with ntropy is further the solvent (Fig. 2-7b). These stable structures of am- phipathic compounds in water, called micelles, may contain hundreds or thousands of molecules. The forces that hold the nonpolar regions of the molecules together are called hydrophobic interactions. The strength of hydrophobic interactions is not due to any intrinsic at- traction between nonpolar moieties. Rather, it results from the system's achieving greatest thermodynamic polar regions. Hydrophobic interactions among lipids stability by minimizing the number of ordered water and between lipids and proteins, are the most impor molecules required to surround hydrophobic portions e tant determinants of structure in biological membranes the solute molecules Hydrophobic interactions between nonpolar amino Many biomolecules are amphipathic, proteins, pig- acids also stabilize the three-dimensional structures of ments, certain vitamins, and the sterols and phospho- proteins lipids of membranes all have polar and nonpolar surface Hydrogen bonding between water and polar solutes regions. Structures composed of these molecules are also causes some ordering of water molecules, but the tabilized by hydrophobic interactions among the non- effect is less significant than with nonpolar solutes. Part

Dispersion of lipids in H2O Clusters of lipid molecules Micelles (b) (a) “Flickering clusters” of H2O molecules in bulk phase Highly ordered H2O molecules form “cages” around the hydrophobic alkyl chains Hydrophilic “head group” O O C H C H H H O Each lipid molecule forces surrounding H2O molecules to become highly ordered. Only lipid portions at the edge of the cluster force the ordering of water. Fewer H2O molecules are ordered, and entropy is increased. All hydrophobic groups are sequestered from water; ordered shell of H2O molecules is minimized, and entropy is further increased. – Hydrophobic alkyl group water, the polar, hydrophilic region interacts favorably with the solvent and tends to dissolve, but the nonpo￾lar, hydrophobic region tends to avoid contact with the water (Fig. 2–7a). The nonpolar regions of the mole￾cules cluster together to present the smallest hy￾drophobic area to the aqueous solvent, and the polar re￾gions are arranged to maximize their interaction with the solvent (Fig. 2–7b). These stable structures of am￾phipathic compounds in water, called micelles, may contain hundreds or thousands of molecules. The forces that hold the nonpolar regions of the molecules together are called hydrophobic interactions. The strength of hydrophobic interactions is not due to any intrinsic at￾traction between nonpolar moieties. Rather, it results from the system’s achieving greatest thermodynamic stability by minimizing the number of ordered water molecules required to surround hydrophobic portions of the solute molecules. Many biomolecules are amphipathic; proteins, pig￾ments, certain vitamins, and the sterols and phospho￾lipids of membranes all have polar and nonpolar surface regions. Structures composed of these molecules are stabilized by hydrophobic interactions among the non￾polar regions. Hydrophobic interactions among lipids, and between lipids and proteins, are the most impor￾tant determinants of structure in biological membranes. Hydrophobic interactions between nonpolar amino acids also stabilize the three-dimensional structures of proteins. Hydrogen bonding between water and polar solutes also causes some ordering of water molecules, but the effect is less significant than with nonpolar solutes. Part Chapter 2 Water 53 FIGURE 2–7 Amphipathic compounds in aqueous solution. (a) Long￾chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules. (b) By clustering together in micelles, the fatty acid molecules expose the smallest 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. 8885d_c02_47-74 7/25/03 10:05 AM Page 53 mac76 mac76:385_reb: