MYLAN EXHIBIT - 1012 (Part 1)
`Mylan Pharmaceuticals, Inc. v. Bausch Health Ireland, Ltd. - IPR2022-00722
`
`

`

`Lehninger Principles of Biochemistry Third Edition
`
`David L. Nelson and Michael M. Cox
`
`Copyright © 2000, 1993, 1982 by Worth Publishers
`
`All rights reserved
`Printed in the United States of America
`
`Library of Congress Cataloging-in-Publication Data
`
`Nelson, David L.
`Lehninger principles of biochemistry / David L. Nelson, Michael M. Cox.— 3rd ed.
`p.cm.
`Includes index.
`ISBN 1-57259-153-6
`
`1, Biochemistry. II. Cox, Michael M._III. Title.I. Nelson, David L. (David Lee), 1942-
`
`
`QD415 .L44 2000
`572—dc21
`
`99-049137
`
`Printing:
`
`5
`
`4
`
`3
`
`2
`
`1
`
`Year:
`
`04
`
`03
`
`02
`
`O01
`
`00
`
`DevelopmentEditor: Morgan Ryan, with Linda Strange and Valerie Neal
`Project Editor: Elizabeth Geller
`Art Director: Barbara Rusin
`
`Design: Paul Lacy
`Production Supervisor: Bernadine Richey
`Layout: York Graphic Services and Paul Lacy
`Photo Editor: Deborah Goodsite
`
`Mlustrations: SusanTilberry (with Alan Landau and Joan Waites), J.B. Woolsey & Associates,
`Laura Pardi Duprey, and York Graphic Services
`Molecular Graphics: Jean-Yves Sgro
`Composition: York Graphic Services
`Printing and Binding: R.R. Donnelley and Sons
`
`Cover (from top to bottom): Cut-away view of GroEL, a protein complex in-
`volved in protein folding; cut-away viewof tobacco mosaic virus, an RNAvirus;
`ribbon modelof a 8-barrel structural domain from UDP N-acetylglucosamine
`acyltransferase; cut-awayviewof the F, subunit of ATP synthase, with bound
`ATP shown as a stick structure; mesh surface imageof the electron-transfer
`protein cytochromec, with its heme group shown asa stick structure.
`
`Cover images created by Jean-Yves Sgro.
`
`Illustration credits begin on p. IC-1 and constitute a continuation of the
`copyright page.
`
`Steenbock Memorial Library
`University of/isconsin-Madisan
`550 BabcockDrive
`Madison, Wi 53706-1293
`
`Worth Publishers
`41 Madison Avenue
`
`NewYork, NY 10010
`
`

`

`
`
`This view of Earth from space shows that most of the
`planet's surface is covered with water. The seas, where
`life probably first arose, are today the habitat of countless
`organisms.
`
`82
`
`Water
`
`Water is the most abundant substanceinliving systems, making up 70% or
`more of the weight of most organisms. The first living organisms doubtless
`arose in an aqueous environment, and the course of evolution has been
`shaped bythe properties of the aqueous mediumin whichlife began.
`This chapter begins with descriptions of the physical and chemical
`properties of water, to which all aspects of cell structure and function are
`adapted. The attractive forces between water molecules and the slight ten-
`dencyof waterto ionize are of crucial importanceto the structure and func-
`tion of biomolecules. We will review the topic of ionization in terms of equi-
`librium constants, pH, and titration curves, and consider how aqueous
`solutions of weak acids or bases and their salts act as buffers against pH
`changes in biological systems. The water molecule andits ionization prod-
`ucts, H* and OH, profoundly influence the structure, self-assembly, and
`propertiesofall cellular components, including proteins, nucleic acids, and
`lipids. The noncovalent interactions responsible for the strength and speci-
`ficity of “recognition” among biomolecules are decisively influenced by the
`solvent properties of water.
`
`Weak Interactions in Aqueous Systems
`Hydrogen bonds between water molecules provide the cohesive forces that
`make water a liquid at room temperature and that favor the extreme or-
`dering of moleculesthatis typical of crystalline water (ice). Polar biomole-
`cules dissolve readily in water because they can replace water-water inter-
`actions with more energetically favorable water-solute interactions,
`In
`contrast, nonpolar biomolecules interfere with water-water interactions but
`are unable to form water-solute interactions—consequently, nonpolar mol-
`ecules are poorly soluble in water. In aqueous solutions, nonpolar molecules
`tend to cluster together.
`Hydrogen bonds andionic, hydrophobic (Greek, “water-fearing”), and
`van der Waals interactions are individually weak, but collectively they have
`a very significant influence on the three-dimensional structures of proteins,
`nucleic acids, polysaccharides, and membranelipids.
`
`Hydrogen Bonding Gives Water Its Unusual Properties
`Water has a higher melting point, boiling point, and heat of vaporization
`than most other commonsolvents (Table 4-1). These unusual properties
`
`

`

`Chapter 4 Water
`
`83
`
`table 4—1]
`of Some CommonSolvents
`porization
`Melting Point, Boiling Point, and Heat of Vaporizati
`
`Water
`Methanol (CH,0H)
`Ethanol (CH;CH,OH)
`Propanol (CH;CH,CH,OH)
`Butanol (CH,(CH,),CH,OH)
`Acetone (CH,COCH;)
`Hexane (CH.(CH,),CH,)
`Benzene (CgH,)
`Butane (CH,(CH,).CH;)
`Chloroform (CHCl,;)
`
`:
`Melting
`point (°C)
`0
`—98
`=117
`-127
`—90
`—95
`-98
`6
`-—135
`-63
`
`Boiling
`point (°C)
`100
`65
`78
`97
`117
`56
`69
`80
`-0.5
`61
`
`Heat of
`vaporization
`(J/g)*
`2,260
`1,100
`854
`687
`590
`523
`423
`394
`381
`247
`
`"The heat energy required to convert 1.0g of a liquid at its boiling point, at atmospheric
`PVOSSUTE, jnto ftsgaseous state at the samme temperature. ft is a direct measure ofthe energy
`required te overcome attractive forces between molecules in the liquid phase.
`
`are a consequence of attractions between adjacent water molecules that
`give liquid water great internal cohesion. A look at the electron structure of
`the H,O molecule reveals the cause of these intermolecular attractions.
`Each hydrogen atom of a water molecule shares an electron pair with
`the oxygen atom. The geometry of the molecule is dictated by the shapes of
`the outer electron orbitals of the oxygen atom, which are similar to the
`bonding orbitals of carbon (see Fig. 3-4a). These orbitals describe a rough
`tetrahedron, with a hydrogen atom at each of two corners and unshared
`electronpairs at the other two corners (Fig. 4-1a). The H—O—Hbondan-
`gle is 104.5°, slightly less than the 109.5° of a perfect tetrahedron because
`of crowding by the nonbonding orbitals of the oxygen atom.
`The oxygen nucleusattracts electrons more strongly than does the hy-
`drogen nucleus (a proton); oxygenis more electronegative (see Table 3-2).
`The sharing of electrons between H and O is therefore unequal; the elec-
`trons are more often in the vicinity of the oxygen atom than of the hydro-
`gen. The result of this unequal electron sharing is two electric dipoles in the
`water molecule, one along each of the H—O bonds; the oxygen atom bears
`a partial negative charge (26), and each hydrogena partial positive charge
`(6°). As a result, there is an electrostatic attraction between the oxygen
`atom of one water molecule and the hydrogen of another (Fig. 4-1c), called
`a hydrogen bond. Throughout this book, we will represent hydrogen
`bonds with three parallel blue lines, as in Figure 4—1c.
`
`
`
`Hydrogen bond
`0.177 nm
`
`Covalent bond
`0.0965 nm
`
`
`
`
`figure 4-1
`Structure of the water molecule. The dipolar nature of
`the H,O molecule is shown by (a) ball-and-stick and
`(b) space-filling models. The dashed lines in (a) represent
`the nonbonding orbitals. There is a nearly tetrahedral
`arrangementof the outer-shell electron pairs around the
`oxygen atom; the two hydrogen atoms have localized
`
`(c)
`partial positive charges (8*) and the oxygen atom has a
`partial negative charge (26 ). (c) Two HO molecules
`joined by a hydrogen bond (designated here, and
`throughout this book, by three blue lines) between the
`oxygen atom of the upper molecule and a hydrogen atom
`of the lower one. Hydrogen bondsare longer and weaker
`than covalent O—H bonds.
`
`

`

`Hydrogen bonding in ice. Each water molecule forms the
`maximum of four hydrogen bonds,creating a regular
`crystal lattice. In liquid water at room temperature and
`atmospheric pressure, by contrast, each water molecule
`hydrogen bonds with an average of 3.4 other water mole-
`cules. The crystal lattice of ice occupies more space than
`that occupied by the same number of HO molecules in
`liquid water; ice is less dense than—and thus floats on—
`liquid water.
`
`Foundations of Biochemistry figure 4-2
`
`Hydrogen bonds are weaker than covalent bonds. The hydrogen bonds
`in liquid water have a bond dissociation energy (the energy required to
`break a bond) of about 20 kJ/mol, compared with 348 kJ/mol for the cova-
`lent C—C bond. At room temperature, the thermal energy of an aqueous so-
`lution (the kinetic energy of motionof 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 mo-
`tion of individual water molecules. Although at any given time most of the
`moleculesin liquid water are engaged in hydrogen bonding, the lifetime of
`each hydrogen bondis less than 1 * 107® s. The apt phrase “flickering clus-
`ters” has been applied to the short-lived groups of hydrogen-bonded mole-
`cules in liquid water. The sum ofall the hydrogen bonds between molecules
`nevertheless confers great internal cohesion on liquid water.
`The nearly tetrahedral arrangement of the orbitals about the oxygen
`atom (Fig. 4—1a) allows each water molecule to form hydrogen bonds with
`as many as four neighboring water molecules. In liquid water at room tem-
`perature and atmospheric pressure, however, water molecules are disorga-
`nized and in continuous motion, so that each molecule forms hydrogen
`bonds with an averageof only 3.4 other molecules. In ice, on the other hand,
`each water moleculeis fixed in space and forms hydrogen bonds with four
`other water moleculesto yield a regular lattice structure (Fig. 4-2). Break-
`age of a sufficient number of hydrogen bondsto destabilize the crystal lat-
`tice of ice requires much thermal energy, which accounts for the relatively
`high melting point of water (Table 4-1). When ice melts or water evapo-
`rates, heat is taken up bythe system:
`
`H,0(s) —> H,0(1)
`
`AH = +5.9 kJ/mol
`
`H,0O(1) —> H,0(g)
`
`AH = +44.0kJ/mol
`
`During melting or evaporation, the entropy of the aqueous system in-
`creases 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 evapo-
`ration of water occur spontaneously; the tendency of the water molecules
`to associate through hydrogen bondsis outweighed by the energetic push
`toward randomness. Recall that the free-energy change (AG) must have a
`negative value for a process to occur spontaneously: AG = AH — T AS,
`where AG represents the driving force, AH the enthalpy change from mak-
`ing and breaking bonds, and AS the change in randomness. Because AH is
`positive for melting and evaporation, it is clearly the increase in entropy
`(AS) that makes AG negative and drives these transformations.
`
`

`

`Chapter 4 Water
`
`85
`
`Hydrogen
`acceptor
`Hydrogen
`donor
`
`No”
`“ae
`I wyxny ova V7
`oO
`O
`N
`oO
`0”
`N
`H
`H
`H
`H
`H
`H
`|
`|
`|
`|
`|
`|
`oO
`Oo
`O
`N
`N
`N
`|
`|
`|
`|
`|
`|
`
`figure 4-3
`Common hydrogen bondsin biological systems. The
`hydrogen acceptor is usually oxygen or nitrogen.
`
`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 an-
`other electronegative atom (the hydrogen donor) in the same or another
`molecule (Fig. 4-3). Hydrogen atoms covalently bonded to carbon atoms
`(which are not electronegative) do not participate in hydrogen bonding.
`The distinction explains why butanol (CH3(CH,).CH,OH) hasa relatively
`high boiling point of 117 °C, whereas butane (CH3(CHz)2CHs3) has a boiling
`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 sugars dissolve readily in
`water because of the stabilizing effect of hydrogen bonds between the hy-
`droxyl groups or carbonyl oxygen of the sugar and the polar water mole-
`cules. Alcohols, aldehydes, ketones, and compounds containing N—H
`bondsall form hydrogen bonds with water molecules (Fig. 4—4) and tend to
`be soluble in water.
`
`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
`
`figure 4-4
`Somebiologically important hydrogen bonds.
`
`Thymine
`
`Adenine
`
`H
`
`R
`
`H
`
`cH3
`
`RC.
`x6
`S27 ne i Sow
`v
`Ze~ yh
`i
`OO
`acted
`i
`sigh
`pots ae
`=
`|
`Cc
`C
`N
`NH
`|
`I
`“ot” Se”
`R
`Cc
`Cc
`
`i~
`
`o* “Nn
`a
`
`R®
`
`H
`
`Bi,
`
`0
`H
`=
`O
`H -H
`
`Bho
`C
`Oo
`H
`id
`“H
`
`Hydrogen bondsare strongest when the bonded moleculesare oriented
`to maximize electrostatic interaction, 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. 4-5). Hydrogen bondsare thus highly directional and capable
`of holding two hydrogen-bonded molecules or groups in a specific geomet-
`ric arrangement. As weshall seelater, this property of hydrogen bonds con-
`fers very precise three-dimensional structures on protein and nucleic acid
`molecules, which have manyintramolecular hydrogen bonds.
`
`0
`HStrong
`- hydrogen bond oo
`
`|
`¥ Weaker
`
`hydrogen bond
`
`—P
`xe
`
`figure 4-5
`Directionality of the hydrogen bond. The attraction
`betweenthe partial electric charges (see Fig. 4-1) is
`greatest when the three atomsinvolved(in this case O, H,
`and 0) 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.
`
`

`

`86
`
`Part | Foundations of Biochemistry
`
`WaterInteracts Electrostatically with Charged Solutes
`Wateris a polar solvent. It readily dissolves most biomolecules, which are
`generally charged or polar compounds (Table 4-2); compounds that dis-
`solve easily in water are hydrophilic (Greek, “water-loving”). In contrast,
`nonpolar solvents such as chloroform and benzene are poor solvents for po-
`lar biomolecules but easily dissolve those thatare hydrophobic—nonpolar
`molecules such as lipids and waxes.
`Water dissolves salts such as NaCl by hydrating andstabilizing the Na™
`and Cl- ions, weakening the electrostatic interactions between them and
`thus counteracting their tendencyto associate in a crystalline lattice (Fig.
`
`tanie —2
`tahia 4
`Some Examples of Polar, Nonpolar, and Amphipathic Biomolecules
`(Shown as lonic Formsat pH 7)
`Polar
`Glucose
`
`CH,OH
`
`
`
`—|
`
`__| Polar groups
`en Nonpolar groups
`
`Glycine
`Aspartate
`
`Lactate
`
`Glycerol
`
`Nonpolar
`Typical wax
`
`+NH,—CH,—COO-
`ae
`-00C—CH,—CH—COO-
`iialome
`OH
`
`oH
`HOCH,.—CH—CH,OH
`
`|
`CH,(CH2);—CH=CH—(CH2)g—CH2—C.0
`|
`CH;(CH,);—CH=CH—(CH2);—CHe
`
`Amphipathic
`Phenylalanine
`
`Phosphatidylcholine
`
`*NH,
`
`()-ctt-Gx —CO0O0-
`
`9
`CH,(CH»);;CH,—C—O—CH,
`|
`CHs(CHa)isCHs—¢—O—CH
`
`0
`
`eae
`
`

`

`Chapter 4 Water
`
`87
`
`
`
`
`
`
`figure 4-6
`Water dissolves manycrystalline salts by hydrating their
`componentions. 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 necessary forlattice formation are weakened.
`
`4-6). The same factors apply to charged biomolecules, compounds with
`functional groups suchas ionized carboxylic acids (—COO), protonated
`amines (—NH}), and phosphate esters or anhydrides. Water readily dis-
`solves such compounds by replacing solute-solute hydrogen bonds with
`solute-water hydrogen bonds, thus screening the electrostatic interactions
`between solute molecules.
`Water is especially effective in screening the electrostatic interactions
`between dissolved ions because of its high dielectric constant, a physical
`property reflecting 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 (7), and the di-
`electric constant (€) of the solvent in which the interactions occur:
`
`Fe= QQ.
`er?
`
`For water at 25 °C, e (which is dimensionless) is 78.5, and for the very non-
`polar solvent benzene, ¢ is 4.6. Thus, ionic interactions are much strongerin
`less polar environments. The dependence on r? is such that ionic attractions
`or repulsions operate only over short distances—in the range of 10 to
`40 nm (depending on the electrolyte concentration) when the solventis
`water.
`
`Entropy Increases as Crystalline Substances Dissolve
`Asa salt such as NaCl dissolves, the Na* and Cl ions leaving the crystallat-
`tice acquire far greater freedom of motion (Fig. 4-6). The resulting in-
`crease in the entropy (randomness) of the system is largely responsible for
`the ease of dissolving salts such as NaCl in water. In thermodynamic terms,
`formation of the solution occurs with a favorable change in free energy:
`AG = AH— T AS, where AH has a small positive value and T ASa large
`positive value; thus AG is negative.
`k= °A aesT&S
`-~% = x
`
`AY ZO
`igs =O
`
`

`

`88
`
`Part | Foundations of Biochemistry
`
`Nonpolar Gases Are Poorly Soluble in Water
`The moleculesof the biologically important gases COs, O., and Np are non-
`polar, In O, and Ng, electrons are shared equally by both atoms. In CO., each
`C=Obondis polar, but the two dipoles are oppositely directed and cancel
`each other (Table 4-3). The movement of molecules from the disordered
`gas phase into aqueous solution constrains their motion and the motion of
`water molecules and therefore represents a decrease in entropy. The non-
`polar nature of these gases and the decrease in entropy when they enterso-
`lution combine to make them very poorly soluble in water (Table 4-3).
`Some organisms have water-soluble carrier proteins (hemoglobin and myo-
`globin, for example) that facilitate the transport of O. Carbon dioxide
`forms carbonic acid (HyCO3) in aqueous solution andis transported as the
`HCO; (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 H,S, also have biological roles in some or-
`ganisms; these gases are polar and dissolve readily in water.
`
`table 4-3 0
`| Solubilities of Some Gases in Water
`
`
`Solubility
`Gas in water (g/L)'arnaeeeStructure* Polarity
`
`
`
`Nitrogen
`N=N
`Nonpolar
`0.018 (40 °C)
`
`| Oxygen
`|
`_ Carbon
`|
`dioxide
`
`| Ammonia
`i
`i
`| Hydrogen
`|
`‘sulfide
`iL
`
`o=0
`8
`5
`Sarr iTt
`0=C=0
`
`8
`
`HHav
`og |
`N
`H
`H |
`tae
`
`hy
`
`Nonpolar
`
`Nonpolar
`
`Polar
`
`Polar
`
`0.035 (50 °C)
`
`0.97 (45 °C)
`
`|
`4
`
`900 (10°C)
`
`1,860 (40 °C)
`
`*The arrows representelectric dipoles; there is a partial negative charge (6~) at the head of the
`arrow, a partial positive charge (6°; not shown here) at thetail.
`‘Note that polar molecules dissolve far better even at low temperatures than do nonpolar
`molecules at relatively high temperatures.
`
`Nonpolar CompoundsForce Energetically Unfavorable Changes
`in the Structure of Water
`When water is mixed with benzene or hexane, two phases form; neither liq-
`uid is soluble in the other. Nonpolar compounds such as benzene and
`hexane are hydrophobic—they are unable to undergo energetically favor-
`able interactions with water molecules, and they actually interfere with the
`hydrogen bonding among water molecules. All molecules or ions in aqueous
`solution interfere with the hydrogen bonding of some water molecules in
`their immediate vicinity, but polar or charged solutes (such as NaCl) com-
`pensate for lost water-water hydrogen bonds by forming new solute-water
`interactions. The net change in enthalpy (AH) for dissolving these solutes
`is generally small. Hydrophobic solutes, however, offer no such compensa-
`tion, and their addition to water may therefore result in a small gain of en-
`thalpy; the breaking of hydrogen bonds between water molecules takes up
`
`

`

`Chapter 4 Water
`
`89
`
`Dispersion of
`lipids in H,O
`Eachlipid
`molecule forces
`surrounding H)O
`molecules to become
`highly ordered.
`
`Clusters oflipid
`molecules
`
`
`
`bo3ReaaeaybPSpiotaie amFgeob
`
`
`& &
`ay)
`
`Only lipid portions
`at the edge of
`the cluster force the
`ordering of water.
`Fewer H0 molecules
`#%, are ordered, and
`entropy is increased.
`
`Hydrophilic
`"head group"
`
`molecules in bulk phase
`
`"Flickering clusters" of Hg0
`
`Highly ordered HzO molecules form
`“cages” around the hydrophobic alkyl chains
`
`4
`
`(a)
`
`
`
`energy from the system. Furthermore, dissolving hydrophobic compounds
`in water produces a measurable decreasein entropy. Water moleculesin the
`immediate vicinity of a nonpolar solute are constrained in their possible ori-
`entations as they form a highly ordered cagelike shell around each solute
`molecule. These water molecules are not as highly ordered as those in the
`crystalline compoundof a nonpolar solute and water (a clathrate), but the
`effect is the same in both cases: the ordering of water molecules reduces
`All hydrophobic
`entropy. The numberof ordered water molecules, and therefore the magni-
`groups are
`tude of the entropy decrease,is proportional to the surface area of the hy-
`sequestered from
`water; ordered
`drophobic solute enclosed within the cage of water molecules. The free-
`shell of HgO
`energy changefor dissolving a nonpolar solute in wateris thus unfavorable:
`molecules is
`minimized, and
`AG = AH — T AS, where AH hasapositive value, AS has a negative value,
`entropyis further
`increased.
`and AG is positive.
`Amphipathie compounds contain regions that are polar (or charged)
`and regions that are nonpolar (Table 4-2). When an amphipathic com-
`pound is mixed with water, the polar, hydrophilic region interacts favorably
`with the solvent andtendsto dissolve, but the nonpolar, hydrophobic region
`tends to avoid contact with the water (Fig. 4-7a). The nonpolar regions of
`the molecules cluster together to present the smallest hydrophobic area to
`the aqueous solvent, and the polar regions are arranged to maximize their
`interaction with the solvent (Fig. 4—7b). These stable structures of amphi-
`pathic 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 attraction between non-
`polar moieties. Rather, it results from the system's achieving greatest ther-
`modynamic stability by minimizing the numberof ordered water molecules
`required to surround hydrophobic portions of the solute molecules.
`
`Micelles
`
`(b)
`
`figure 4-7
`(a) Long-
`Amphipathic compounds in aqueous solution.
`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.
`
`

`

`90
`
`Part |
`
`Foundations of Biochemistry
`
`Ordered water
`
`substrate and enzyme
`
`interacting with
`“eea.
`
`p o&
`
`Sr
`Robo
`»
`Seec
`
`Enzyme
`
`Disordered water
`displaced by
`enzyme-substrate
`
`interaction
`
`Enzyme-substrate
`interaction stabilized
`by hydrogen-bonding,
`ionic, and hydrophobic
`interactions
`
`figure 4-8
`Release of ordered water favors formation of an
`enzyme-substrate complex. While separate, both enzyme
`and substrate force neighboring water molecules into an
`ordered shell. Binding of substrate to enzyme releases
`some of the ordered water, and the resulting increase in
`entropy provides a thermodynamic push toward formation
`of the enzyme-substrate complex.
`
`Many biomolecules are amphipathic; proteins, pigments, certain vita-
`mins, and the sterols and phospholipids of membranes all have polar and
`nonpolar surface regions. Structures composed of these molecules are sta-
`bilized by hydrophobic interactions among the nonpolar regions. Hy-
`drophobic interactions amonglipids, and betweenlipids and proteins, are
`the most important determinants of structure in biological membranes. Hy-
`drophobic interactions between nonpolar amino acids also stabilize the
`three-dimensional folding patterns of proteins.
`Hydrogen bonding between water and polar solutes also causes some
`ordering of water molecules, but the effectis less significant than with non-
`polar solutes. Part of the driving force for binding of a polar substrate (re-
`actant) to the complementary polar surface of an enzymeis the entropy in-
`crease as the enzymedisplaces ordered water fromthe substrate (Fig. 4-8).
`
`Van der Waals Interactions Are Weak Interatomic Attractions
`When two uncharged atoms are brought very close together, their sur-
`rounding electron clouds influence each other. Random variations in the po-
`sitions of the electrons around one nucleus may create a transient electric
`dipole, which induces a transient, opposite electric dipole in the nearby
`atom. The two dipoles weakly attract each other, bringing the two nuclei
`closer. These weakattractions are called van der Waals interactions. As
`the two nuclei draw closer together, their electron clouds begin to repel
`eachother. At the point when the vander Waalsattraction exactly balances
`this repulsive force, the nuclei are said to be in van der Waals contact. Each
`atom has a characteristic van der Waals radius, a measure of how close
`that atom will allow another to approach (see Table 3-1). In the “space-
`filling” molecular models shown throughout this book (e.g., Fig. 83-7c) the
`atoms are depicted in sizes proportional to their van der Waals radii.
`
`Weak Interactions Are Crucial to Macromolecular
`Structure and Function
`The noncovalent interactions we have described (hydrogen bonds and
`ionic, hydrophobic, and van der Waals interactions) (Table 4-4) are much
`weaker than covalent bonds. An input of about 350 kJ of energy is required
`to break a mole of (6 x 10°°) C—Csingle bonds, and about 410 kJ to break
`a mole of C—H bonds,butaslittle as 4 kJ is sufficient to disrupt a mole of
`typical van der Waals interactions. Hydrophobic interactions are also much
`weaker than covalent bonds, although they are substantially strengthened
`by a highly polar solvent (a concentratedsalt solution, for example). Ionic
`interactions and hydrogen bondsare variable in strength, depending on the
`polarity of the solvent, but they are alwayssignificantly weaker than cova-
`lent bonds.
`In aqueous solvent at 25 °C, the available thermal energy can
`be of the same order of magnitude as the strength of these weak interac-
`tions, and the interaction between solute and solvent (water) molecules is
`nearly as favorable as solute-solute interactions. Consequently, hydrogen
`bondsandionic, hydrophobic, and van der Waals interactions are continu-
`ally formed and broken.
`Although these four types of interactions are individually weak relative
`to covalent bonds, the cumulative effect of many such interactions with a
`protein or nucleic acid can be very significant. For example, the noncova-
`lent binding of an enzymeto its substrate may involve several hydrogen
`bonds and one or moreionic interactions, as well as hydrophobic and van
`der Waals interactions. The formation of each of these weak bonds con-
`tributes to a net decrease in the free energy of the system. The stability of
`a noncovalentinteraction such as that of a small molecule hydrogen-bonded
`to its macromolecular partneris calculable from the binding energy. Stabil-
`
`

`

`Chapter 4 Water
`
`table 4-4
`
`Four Types of Noncovalent (“Weak”) Interactions
`among Biomolecules in Aqueous Solvent
`
`Hydrogen bonds
`Between neutral groups
`
`.
`Between peptide bonds
`
`\
`ae 1 H-—O—
`
`NN
`a
`
`C=O'''H—N
`
`rs
`
`lonic interactions
`
`Attraction
`
`°
`
`~O—C—
`
`—+*NH;
`
`~<
`
`Repulsion
`
`Hydrophobic interactions
`
`Van der Waals interactions
`
`Any two atoms in
`close proximity
`
`ity, as measured by the equilibrium constant (see below) of the binding re-
`action, varies exponentially with binding energy. The dissociation of two
`biomolecules associated noncovalently by multiple weak interactions (such
`as an enzyme andits bound substrate) requires all these interactions to be
`disrupted at the same time. Because the interactions fluctuate randomly,
`such simultaneous disruptions are very unlikely. The molecular stability be-
`stowed by two or five or 20 weak interactions is therefore much greater
`than would be expected intuitively from a simple summation of small bind-
`ing energies.
`Macromolecules such as proteins, DNA, and RNA contain so manysites
`of potential hydrogen bonding or ionic, van der Waals, or hydrophobic in-
`teractions that the cumulative effect of the many small binding forces is
`enormous. For macromolecules, the most stable (native) structure is usu-
`ally that in which weak-bonding possibilities are maximized. The folding of
`a single polypeptide or polynucleotide chain into its three-dimensional
`shape is determined bythis principle. The binding of an antigen to a spe-
`cific antibody depends on 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 neurotransmitter to its cellular receptor protein
`is the result of weakinteractions. One consequence of the large size of en-
`gvmes and receptors Is that their extensive surfaces provide many oppor-
`tunities for weak interactions. At the molecular level, the complementarity
`between interacting biomolecules reflects the complementarity and weak
`interactions between polar, charged, and hydrophobic groups on the sur-
`faces of the molecules.
`
`
`
`

`

`Qo = H,0
`Q= Solute
`
`Forming
`
`ice crystal
`
`(a)
`
`(b)
`
`In pure water,
`every molecule at
`the surface is H,O, and
`all contribute to the
`vapor pressure. Every
`molecule in the bulk
`solution is H,O, and
`can contribute to
`formation of ice crystals.
`
`In this solution, the
`effective concentration
`of HO is reduced; only
`3 of every 4 molecules
`at the surface and in the
`bulk phase are H,O.
`The vapor pressure of
`water andthe tendency
`of liquid water to enter
`a crystal are reduced
`proportionately.
`
`figure 4-9
`Solutes alter the colligative properties of aqueous
`solutions. (a) At 101 kPa (1 atm) pressure, pure water
`boils at 100 °C and freezes at 0 °C. (b) The presence of
`solute molecules reduces the probability of a water mole-
`cule leaving the solution and entering the gas phase,
`thereby reducing the vapor pressure of the solution and
`increasing the boiling point. Similarly, the probability of a
`water molecule colliding with and joining a forming ice
`crystal is reduced when someof the molecules colliding
`with the crystal are solute, not water, molecules. The
`effect is depression of the freezing point.
`
`figure 4-10
`Osmosis and the measurement of osmotic pressure.
`(a) Theinitial state. The tube contains an aqueous solu-
`tion, the beaker contains pure water, and the semiperme-
`able membrane allows the passage of water but not
`solute. Water flows from the beaker into the tube to
`equalize its concentration across the membrane. (b) The
`final state. Water has moved into the solution of the non-
`permeant compound, diluting it and raising the column of
`water within the tube. At equiliorium, the force of gravity
`operating on the solution in the tube exactly balances the
`tendency of water to move into the tube, where its con-
`centration is lower. (c) Osmotic pressure (II) is measured
`as the force that must be applied to return the solution in
`the tube to the level of that in the beaker. This force is
`proportional to the height, h, of the column in (b).
`
`Solutes Affect the Colligative Properties of Aqueous Solutions
`Dissolved solutesofall 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 (“tied together”) proper-
`ties because the effect of solutes on all four properties has the samebasis:
`the concentration of water is lower in solutions than in pure water. Theef-
`fect of solute concentration on the colligative properties of water is inde-
`pendent of the chemical properties of the solute; it depends only on the
`numberof solute particles (molecules, ions) in a given amount of water. A
`compound such as NaCl, which dissociates in solution, has twice the effect
`on osmotic pressure, for example, as an equal number of moles of a nondis-
`sociating solute such as glucose.
`Dissolved solutes alter the colligative properties of aqueous solutions
`by lowering the effective concentration of water. For example, when a sig-
`nificant fraction of the molecules at the surface of an aqueous solution are
`not water but solute, the tendency of water molecules to escape into the va-
`por phase—the vapor pressure—is lowered (Fig. 4—9). Si