Biomolecules in Water...biochemical components is the most useful method for diagnosis of diseases. (© Klaus Guldbrandsen/Photo Researchers, Inc.) Biomolecules in Water 2.3 Cellular

Blood, an aqueous-based solution, plays an essential

role in the distribution of cells, proteins, nutrients,

gases, and many other biological materials

throughout organisms. Each of us has 5–6 liters of

the complex solution coursing through our arteries

and veins. Blood may be separated into its two parts

by centrifugation. The sediment is composed of three

types of blood cells: erythrocytes, leukocytes, and

platelets (thrombocytes). The supernatant, the liquid

portion called blood plasma, is a solution consisting

of approximately 90% water and 10% dissolved

solutes such as plasma proteins, glucose, hormones,

cholesterol, vitamins, inorganic ions, and many other

biomolecules. Collection of blood samples from an

individual and clinical analysis of changes in the

biochemical components is the most useful method

for diagnosis of diseases.

(© Klaus Guldbrandsen/Photo Researchers, Inc.)

Biomolecules in Water

2.3 Cellular Reactions of WaterIonization of WaterpH and pKTitration CurvesThe Henderson–Hasselbalch EquationDrug Efficacy Depends on Water

Solubility

36

2.1 Water, the Biological SolventNoncovalent Interactions in BiomoleculesExamples of Noncovalent InteractionsCharacteristics of Noncovalent

InteractionsThe Structure of WaterThe Importance of Hydrogen Bonds

2.2 Hydrogen Bonding andSolubilityPhysical Properties of WaterWater as a Solvent

2.4 Buffer SystemsMaintaining a Constant pHEffective Buffering RangeLaboratory Buffers

C H A P T E R

2

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Introduction 37

Table 2.1Percent by weight of water inorgans of the human body

PercentTissue or by WeightOrgan of Watera

Skeletal muscle 79b

Heart 83b

Liver 71Kidney 81Spleen 79Lung 79Brain 77

a In adults.b Fat-free tissue.

Water! Although a seemingly simple and abundant substance containing only theatoms of hydrogen and oxygen, it has extraordinary physical, chemical, and biologi-cal properties. Water is vital to all forms of life and makes up about 70% to 85% ofthe weight of a typical cell (Table 2.1). In addition, extracellular fluids such as blood,cerebrospinal fluid, saliva, urine, and tears are aqueous-based solutions. Many scien-tists believe that life began in an aqueous environment and, during the early stagesof evolutionary development, all living organisms resided in water. Although plantsprobably evolved first, many other forms of life developed lungs and were able tomove to land. Some organisms (unicellular and multicellular) still require not justinternal water but a constant extracellular environment that is aqueous.These organ-isms may live in rivers, lakes, and oceans or sheltered in the aqueous environment ofanother, larger cell.

The Biological Roles of WaterWater plays many roles in cells and organisms and has great influence on the struc-ture and behavior of all biomolecules.

• Water is important as a biological solvent. A study of biomolecules is not completewithout an understanding of the extraordinary properties of water as a solvent.Water provides a medium for metabolic reactions; it is literally the “matrix of life.”Because most biological fluids are aqueous-based, water provides for the deliveryto the cell of nutrients for growth and removal of wastes for general cleansing ofcells.The cellular fluids also assist in the absorption and action of pharmaceuticalcompounds by dissolving them and transporting them to target organs and cells.It is important to note that water, although a very effective solvent, is not a uni-versal biological solvent as is sometimes declared. All biomolecules are not solu-ble in water and this, indeed, is fortunate. Because of this, organisms can developcompartmentation of cell structure and function by creating partitions (mem-branes) using molecules that are water insoluble. One of the newest developmentsin cellular biochemistry is the discovery of aquaporins, membrane proteins thatform channels through which water molecules are transported in and out of cells(Section 9.3).

• Water serves as an essential buffer to regulate temperature and pH. With a high spe-cific heat capacity, water is able to absorb large amounts of energy in the form ofheat released from biochemical reactions. As a result, the water medium acts tomaintain a constant cell temperature. Water also serves as a solvent to dissolve

Water, which plays roles as abiological solvent, reactant molecule,and temperature regulator, is essentialto life. (© Douglas Faulkner/Photo Researchers.)

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38 Chapter 2 Biomolecules in Water

Water, the Biological SolventLearning ObjectiveHave a general knowledge of the chemical and biological properties of water and of the

importance of weak, noncovalent interactions in biochemistry.

Noncovalent Interactions in BiomoleculesBefore we look at the physical and chemical properties of water, we need to con-sider bonding arrangements that are important in water molecules and in how waterinteracts with other molecules. In Chapter 1, we discovered that biomolecules largeand small are made by the covalent combining of elements forming C C, C H,C O, C N, and other strong bonds. These covalent bonds literally hold the atomsof a molecule together. Weaker interactions, called noncovalent interactions, bringtogether whole biomolecules for specific purposes. Four types of noncovalent inter-actions are important: van der Waals forces, ionic bonds, hydrogen bonds, andhydrophobic interactions. Properties and examples of these interactions are reviewedin Table 2.2. One of these types of stabilizing forces (hydrogen bonds) are at workwhen water dissolves polar molecules and when two strands of DNA are broughttogether in the double helix.

Examples of Noncovalent Interactions

Throughout our study of biochemistry we will encounter many examples where non-covalent molecular interactions bring together, in specific ways, two different mole-cules or different regions of the same molecule. Molecules have the ability torecognize and interact (bind) specifically with other molecules.We will use the termmolecular recognition to describe this general phenomenon.The importance of theseinteractions in biology is that the combination of two molecules or the organizedfolding of a single molecule will lead to biological function not present in individualmolecules or in unfolded, randomly arranged molecules.

The interactions important in molecular recognition are often between a smallmolecule (called a ligand, L) and a macromolecule (M):

L � M LM∆

¬¬¬¬

2.1

substances that regulate hydrogen ion concentration (pH). Biological moleculeswill function properly only in an environment of constant pH. Buffering sub-stances as simple as bicarbonate and as complex as proteins react with water tomaintain a remarkably constant pH level in intracellular and extracellular fluids.

• Water is a participant in many biochemical reactions. One of the most commonbiological reactions is cleavage of a chemical bond by water (hydrolysis) asobserved in the initial steps of digestion of proteins, carbohydrates, and nucleicacids. Water is also a principal reactant in the photosynthesis process:

Here water acts as a reducing agent, a source of electrons to reduce carbon dioxidefor the manufacture of glucose. The process of respiration, the final stage of energymetabolism in animals, generates water from O2 by oxidation–reduction reactions(see Chapter 17).

Clearly, water is not just an inert bystander in the cell, but a selectively reactivemolecule with unique properties that greatly affect biochemical molecules and bio-logical processes. In this chapter, we shall examine some of the unusual properties ofwater and learn how they influence structure and reactivity of biomolecules.

Carbohydrates

6 CO2 � 6 H2O hn

" C6H12O6 � 6 O2

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2.1 Water, the Biological Solvent 39

LM represents a complex, held together by noncovalent interactions, with special-ized biological function. The action of hormones is a good example. A hormoneresponse is the consequence of weak, but specific, interactions between the hormonemolecule and a receptor protein in the membrane of the target cell (Special Topic I).

Biochemical reactions also provide many examples of the importance of nonco-valent interactions. Before a metabolic reaction can occur, a small substrate mole-cule must physically interact in a certain well-defined manner with a macro-molecularcatalyst, an enzyme (Chapters 5 and 6).The biochemical action of a drug also dependson molecular interactions. The drug is first distributed throughout the body via thebloodstream. Drugs in the bloodstream are often noncovalently bound to plasmaproteins, which act as carriers.When the drug molecules are transported to their siteof action, a second molecular interaction is likely to occur. The drug will likely bindto a receptor protein or other proteins. Many drugs elicit their effects by then inter-fering with biochemical processes. This may take the form of enzyme inhibition,where the drug molecule binds to a specific enzyme and prohibits binding of normalreactant and, therefore, inhibits catalytic action. Intramolecular noncovalent inter-actions (those within a molecule) also play significant roles in biomolecularprocesses—for example, stabilizing the folding of protein, DNA, and RNA moleculesinto regular, three-dimensional arrangements.

Table 2.2Properties and examples of noncovalent interactions

Stabilization LengthType Brief Description and Example Energy (kJ/mol) (nm)

Hydrogen bonds Between a hydrogen atom covalently bonded to an 10–30 0.18–0.30electronegative atom and a second electronegative atom (H-bond donor � red atoms, H-bond acceptor � green atoms)

Between neutral groups

Between peptide bonds

van der Waals Between molecules with temporary dipoles induced by fluctuating 1–5 0.1–0.2interactions electrons. May occur between any two atoms in close proximity

Hydrophobic The tendency of nonpolar groups and molecules 5–30 —interactions to stick together or cluster in aqueous solutions

Ionic bonds Attractive interactions that occur between oppositely charged 20–30 0.25atoms or groups

Na�Cl�

R COO� H3 R¬N�

¬

HC

H3C CH3

CH2

HC

H3C CH3

CH2

C O H N

H

C O H O

H

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Characteristics of Noncovalent InteractionsAll molecular interactions that are the basis of molecular recognition have at leastthree common characteristics:

• First, the forces that are the basis of these interactions are noncovalent and rela-tively weak. The strengths of these interactions are in the range of 1–30 kJ/molcompared to about 350 kJ/mol for a carbon–carbon single bond, a typical covalentbond. A single noncovalent bond is usually insufficient to hold two moleculestogether. DNA, RNA, and protein molecules have numerous functional groupsthat participate in noncovalent interactions. A collection of many of these inter-actions will lead to greatly stabilized complexes.

• Second, noncovalent interactions are reversible. Noncovalent interactions are ini-tiated when diffusing (wandering or moving) molecules or regions of a moleculecome into close contact. Diffusion is brought about by thermal motions.An initialclose encounter may not always result in the successful formation of a complex.Afew weak bonds may form but may be disrupted by thermal motion, causing themolecules to dissociate. Therefore, bonds may constantly form and break untilenough bonds have accumulated to result in an intermediate with a transient butsignificant existence.The complex can then initiate a specific biological process.Anintermediate rarely lasts longer than a few seconds. Eventually, thermal motionscause the complex to dissociate to the individual molecules. Reversibility is animportant characteristic of these interactions so that a static, gridlock situationdoes not occur. The biological process initiated by the complex LM must have astarting time and an ending.

• Third, the binding between molecules is specific. Imagine that the interactionsbring together two molecular surfaces. The two surfaces will be held together ifnoncovalent interactions occur. If on one surface there is a nonpolar moleculargroup (phenyl ring, hydrophobic alkyl chain, etc.), the adjacent region on the othersurface must also be hydrophobic and nonpolar. If a positive charge exists on onesurface, there may be a neutralizing negative charge on the other surface.A hydro-gen bond donor on one surface can interact with a hydrogen bond acceptor onanother. Simply stated, the two molecules must be compatible or complementaryin a physical and chemical sense so the development of stabilizing forces can holdmolecules together.The concept of molecular recognition will take on many formsin our continuing studies of biochemistry.

The Structure of WaterThe arrangement of hydrogen and oxygen atoms in the water molecule is nonlinearwith an H O H bond angle of 104.5° (Figure 2.1a). The electronegativity valuefor the oxygen atom (3.5) is approximately one and one-half times that of hydrogen(2.1). Therefore, the electrons of the two covalent bonds are not shared equally; theoxygen atom has a stronger pull on the electrons and takes on a partial negativecharge (��).The hydrogen atoms are left with a partial positive charge (��), becausethey do not have equal access to the bonding electrons. This gives rise to a moleculewith a dipolar structure: a negative end sometimes called the “head” (oxygen) andpositive ends sometimes called “tails” (hydrogens) (Figure 2.1b).The water moleculeis electrically neutral (no net charge) but has a relatively large dipole momentbecause of its bent geometry.Water can be contrasted with another molecule of bio-chemical significance, CO2, which also has polar bonds caused by electronegativitydifferences between the carbon and oxygen atoms but no dipole moment because itis linear (Figure 2.2).

¬¬

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2.1 Water, the Biological Solvent 41

H

O

(b)

Hδ�δ�H

van der Waals envelope

O H covalentbond distance � 0.958 Å

Oδ�

(a)

104.5°H δ�δ�

δ�

van der Waals radius of H � 1.2 Å

van der Waals radius of O � 1.4 Å

Figure 2.1 (a) The structure of thewater molecule showing the relativesize of each atom by the van der Waalsradius. Covalent bonds hold togetheroxygen and hydrogen atoms. The polarcharacter, which is the result ofelectronegativity differences betweenoxygen and hydrogen, is indicated bythe partial charges (�� and ��) onatoms. (b) Water has a dipole momentbecause of its bent geometry. Thearrows pointing to the moreelectronegative atom are used to showbond polarity.

δ�

Cδ�O Oδ�

Figure 2.2 The CO2 molecule,although composed of polar bonds, hasno dipole moment because it is linear.The electronegativity differencesbetween C and O atoms are indicatedby the partial charges (�).

O OH

H H

��

H

Figure 2.3 Hydrogen bond betweentwo water molecules. The hydrogenatom (partially charged) of one watermolecule interacts with a lone pair ofelectrons in an orbital of the oxygenatom of another water molecule.

The Importance of Hydrogen BondsThe characteristics of water as described here have profound consequences for itsstructure and interactions with biomolecules.Water molecules can interact with eachother by attraction of a positive tail (hydrogen) with a negative head (oxygen) asshown in Figure 2.3. This favorable interaction results in a hydrogen bond (seeTable 2.2).This type of bond, which may be represented as X H A, is formed whenan electronegative atom (A), such as oxygen or nitrogen, interacts with a hydrogenatom that is slightly positive or acidic as in X H��, where X may be nitrogen, oxy-gen, or sulfur.The electronegative atom A is defined as the hydrogen bond acceptorand the X H group as the hydrogen bond donor. The hydrogen bond is strongestwhen the three atoms, X H A, are in a straight line (180°) with the hydrogen atominteracting directly with a lone pair electron cloud of A.The hydrogen bond distancein water is about 0.18 nm (1.8 Å) and has a bond energy of about 20 kJ/mol(5 kcal/mol) compared to 0.096 nm (0.96 Å) and 460 kJ/mol (110 kcal/mol) for anO H covalent bond.

Water structure has great significance in biochemistry because many biomoleculeshave atoms that can hydrogen bond with water, with themselves, and with other mol-ecules. Some biochemical examples of hydrogen bonding are shown in Figure 2.4.Functional groups that participate in hydrogen bonding include (1) the hydroxylgroups in alcohols, organic acids, and carbohydrates; (2) carbonyl groups in aldehy-des, ketones, acids, amides, and esters; and (3) N H groups in amines and amides.The specific hydrogen bonding that occurs between complementary base pairs inDNA and RNA (Figure 2.4d) is discussed in Chapter 10. Although the strength of asingle hydrogen bond may be small, the enormous number of potential hydrogenbonding groups in biomolecules more than makes up for their individual weaknesses.

¬

¬

S¬¬

¬

S¬

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O

O

CH

N

RH

N

R

H

C

C

HC

(c)

CO

C

H

H

N H

H

N

CC

O

N

N

CH3

H

N

N CH

NC

CC

C

Adenine

Thymine

(d)

H

O

H

O

R

R

O

H

O

R

H H(a)

O

O

H

C

H

R X

(b)

Figure 2.4 Hydrogen bonds ofbiological importance: (a) between analcohol and water or between alcoholmolecules; (b) between a carbonylgroup and water (X � H, R, OH, OR,or NH2); (c) between two peptidechains, the carbonyl group of onepeptide bonds to an N H of another;(d) between complementary base pairsin DNA. Atoms involved as hydrogenbond donors are in red; hydrogen bondacceptors are in green.

¬

Hydrogen Bonding and SolubilityLearning ObjectiveBe able to explain how water functions as a solvent.

Physical Properties of WaterThe unusual physical properties of water are best understood by comparison to sub-stances of similar structure and molecular weight (Table 2.3). Water has a higherboiling point, melting point, and viscosity than any other hydride of a nonmetallicelement. These peculiar properties are the result of water’s unusually high internalcohesiveness or the tendency of water molecules to “stick together,” which is due toan extensive network of hydrogen bonds. Each water molecule theoretically canhydrogen bond with four neighboring water molecules (Figure 2.5a). In reality, theaverage number of hydrogen bonds to each molecule in liquid water at 10°C is aboutthree. The number of hydrogen bonds decreases with increasing temperature.The theoretical number of four interacting neighbors for each water molecule isapproached in crystalline ice (Figure 2.5b).

A close examination of water structure by X-ray and neutron diffraction tech-niques reveals more detailed features. The network of hydrogen bonds as shown inFigure 2.5 is a snapshot representing an instant in time. The actual structure isdynamic, with water molecules undergoing constant geometrical reorientations andforming new hydrogen bonds with other neighboring water molecules. This changehappens for each molecule about once every 10�12 s. The term “flickering clusters”has sometimes been used to describe the constantly changing network of hydrogenbonds in liquid water. Using powerful X-rays generated at the EuropeanSynchrotron Radiation Facility in Grenoble, France, it has been shown that thehydrogen bonds that hold water molecules together in ice have substantial covalentcharacter, so they are a lot stronger than previously thought.

2.2

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2.2 Hydrogen Bonding and Solubility 43

(a) (b)

H

H

O

H

HH

O

H

O

H

O

H

H

H

δ�

δ�

δ�

δ�δ�

δ�

O

Hydrogen bonds

Figure 2.5 The network of potential hydrogen bonds in water. (a) The center water molecule may form hydrogenbonds with up to four neighboring molecules, but the average is about three. The network structure is constantlychanging, with water molecules undergoing geometrical reorientations and forming new hydrogen bonds withother neighboring water molecules. (b) In ice, hydrogen bonding leads to the formation of a crystalline lattice.

Water as a SolventWater displays an exceptional capacity as a solvent to dissolve many of the biomole-cules found in living cells. In addition, water is an important medium for dissolving andtransporting other chemicals we ingest such as nutritional supplements and pharma-ceutical compounds. Many natural and synthetic molecules that are ionic as wellas some that are polar and uncharged are soluble in water. Here we will examine theinteractions of water with these compounds.

Table 2.3A comparison of some physical properties of water with hydrides of other nonmetallic elements: N, C, and S

Property H2O NH3 CH4 H2S

Molecular weight 18 17 16 34Boiling point (°C) 100 �33 �161 �60.7Freezing point (°C) 0 �78 �183 �85.5Viscositya 1.01 0.25 0.10 0.15

a Units are centipoise.

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Polar Compounds Many uncharged biomolecules readily dissolve in water becausethey have polar functional groups that form favorable dipole–dipole interactions. Afew examples of these compounds, including alcohols, amines, amides, and esters, arefeatured in Figure 2.4. Further examples are shown in Figure 2.6a. Extensive hydro-gen bond networks are possible where the atoms of the polar functional groups com-bine with identical molecules, similar molecules, and/or water. Because of a favorableattraction to water molecules, ionic and polar compounds are said to be hydrophilic,a word of Greek origin translated as water (hydro) and loving (philic). Not all com-pounds containing polar functional groups are water soluble.Those with a relativelylarge hydrocarbon component (usually greater than four carbon atoms) are usuallyinsoluble unless an ionic group or several polar groups are present. Cyclohexane isinsoluble in water; but if one of the carbon–hydrogen groups is converted to an alde-hyde or ketone and a hydroxyl group is substituted on each of the remaining five car-bon atoms, the molecule becomes similar to a carbohydrate such as glucose and iswater soluble.

Ionic Compounds Ionic compounds such as sodium acetate (CH3COO�Na�) andmonopotassium phosphate (K�H2 ) dissolve in water because their individualions can be solvated (hydrated) by polar water molecules (Figure 2.6b). The nega-tive dipole (oxygen atom) of water binds favorably with the Na� or K� forming adipole–ion interaction. The acetate and phosphate anions are hydrated alsoby dipole–ion interactions. In this case, the partially positive hydrogen ends of waterassociate with the negative charges on the anion. The attraction between most ionsand polar water molecules is strong enough to overcome the tendency of anions andcations to recombine. The great ion-solvating ability of water is illustrated by theNa�, K�, and Cl� concentrations in human blood of 0.14 M, 0.004 M, and 0.10 M,

PO 4�

H

H

H H

H H

H

H

H

H

HH

HH

H

H

O

O

C

O

Hδ�

δ�H Hδ�

X R

δ�

Oδ�H Hδ�

δ�

δ�

Oδ�H Hδ�

Oδ�HHδ� Hδ�

δ�

δ�

δ�

OR Hδ�

δ�

(a)

δ�

O

O

OO

δ�

δ� δ�

δ�

δ�

Na�

δ�

δ�

O δ�

δ�

δ�

δ� δ�

Oδ�

δ� δ�

δ�

δ�

Oδ�

δ�

δ�

δ�

O

δ�

δ�

δ�

(b)

CH3COO–

Figure 2.6 Chemicals are madesoluble in water by noncovalentinteractions. (a) Dipole–dipoleinteractions. The partially chargedpositive atoms (hydrogen) of water andalcohol are attracted to oxygen atomdipoles of water and alcohol. Thecarbonyl group of an aldehyde, ketone,or acid can also be solvated by water.(b) Ion–dipole interactions. Thepositively charged sodium ion issurrounded by water moleculesprojecting their partially negativeoxygen atoms (dipoles). The acetate ioninteracts with the partially positivehydrogen atoms (dipoles) of water.

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2.2 Hydrogen Bonding and Solubility 45

respectively. Hundreds of other ionic substances are dissolved in blood. Even macro-molecules, including proteins, nucleic acids, some lipids, and some carbohydrates,under in vivo conditions exist as hydrated ions.

Nonpolar Compounds Nonpolar compounds are usually not water soluble becausethey contain neither ions nor polar functional groups that can interact favorably withwater molecules. Hence, they are called hydrophobic (water fearing). Decane andbenzene are examples of hydrophobic molecules (Figure 2.7). Some significant bio-chemicals have dual properties; they have both nonpolar and ionic characteristics.They are classified as amphiphilic (amphi, on both sides or ends, and philic, loving).This class of compounds is best illustrated by metal salts of long-chain carboxylicacids (Figure 2.8). A specific example, sodium stearate, has an ionic side or end (thecarboxylate anion associated with the sodium cation) and a nonpolar hydrocarbonend. This molecule must be confused when placed in water solution. The observedresult of this experiment is not the formation of a true solution but the self-assemblyof the acid molecules into aggregates called micelles. The amphiphilic moleculesavoid water contact in the hydrophobic region by pointing their hydrocarbon chains(“tails”) to the water-free interior of the aggregate. The surface of each micelle iscomposed of the ionic “heads” stabilized by electrostatic interaction with metalcations and water.The favorable association of nonpolar hydrocarbon tails inside themicelle is defined as a hydrophobic interaction (see Table 2.2). In simplified terms,this situation is favorable because less energy is required to form the micelle than ifthe hydrocarbon chains were allowed to point out into the water, disrupting the net-work of hydrogen bonds. One important practical application of micellar solutions ofsodium stearate and other similar compounds is their use as soaps to “solubilize” oiland grease in water. Owing to their dual chemical character, soaps are able to trap oilin the nonpolar region of the micelle and yet remain dispersed in aqueous solutionsby hydration of the ionic region. Micelle formation is also the key to construction ofbiological membranes (see Chapter 9).

Much of this section has focused on how the structure of water is affected by thepresence of solute molecules. It is also important to note changes in the structures ofbiomolecules brought about by the presence of water. Because of their dual charac-ter, long-chain carboxylic acid salts and other amphiphilic lipids take on a specialarrangement in water. Proteins and nucleic acids also contain hydrophobic regionsand ionic functional groups; therefore, these are amphiphilic substances. As we shalldiscover in later chapters, these biomolecules in water solution fold into conforma-tions that bury hydrophobic regions in water-free areas and expose ionic and polarfunctional groups to water molecules. Important consequences result because it isoften found that these complex and ordered arrangements of biomolecules are theonly ones with biological activity.

CC

C

CH

H H

H

H

H

H

H

H

H

H

H

C

C

C

C

C CWater moleculesin cage aroundhydrocarbon chain

HH

H

H

H H

H

H

H

H

H H

H

H

CC

C

CH

H H

H

H

H

H

H

H

H

H

H

C

C

C

C

C C HH

H

H

H H

H

H

H

H

H H

H

H

OO

O OO

O

OOOOOO

O

Figure 2.7 Because hydrophobicmolecules have no polar groups tointeract with water, they have to besurrounded by a boundary of watermolecules. The formation of thishighly ordered cage of water requiresmuch energy, which comes fromhydrophobic interactions.

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Sodium Stearate Micelle

O

OO

OO

O

O

O

O

O

O

OO

O

�OC

O

O O

O

OO

H2C

H2C

CH2

OO

O

O

O

O O

O

O

O

O

O

O

O

O

O

O

O

OO

Polar head of sodium stearate

Polar head

Nonpolar tail of sodium stearate

Nonpolar tail

Example:

Key:

Na�

H2C

CH2

H2C

CH2

H2C

CH2

H2C

CH2

H2C

CH2

H2C

CH2

H3C

CH2

⎧ ⎪ ⎨ ⎪ ⎩

Amphiphilic Compound

Nonpolarcore

H H

H

H

H

H

H

H H

H

H

H

H

H

H

H

H

H

H

H

H

HH H

HH

HHH H

H H

H

H

H

H

H

H

H

H

H

H

H

H

H

H

H

HH

HH

HH

H

H

H

H

H

H H

H H

H

H

H

H

H

H

H

H HH

H

H

H

H

H

H

Figure 2.8 Formation of a micellefrom the sodium salt of a long-chaincarboxylic acid. The nonpolarhydrocarbon tails of the acid arrangethemselves to avoid contact with water.The negatively charged carboxyl groupsinteract with water by formingion–dipole interactions.

Before You Go On. . .

1. Show examples of all the possible types of hydrogen bonding that would existbetween the molecules in an ethanol–water solution (50:50).

2. Show how the functional groups of the amino acid glycine, H3 CH2COO�, maybe solvated in an aqueous solution.

3. Show how the chemical salt sodium benzoate may be solvated in an aqueoussolution. Sodium benzoate is an inhibitor of fungal growth that is often used as afood preservative.

COO�Na�

N�

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2.3 Cellular Reactions of Water 47

4. The structure for the neurotransmitter and hormone epinephrine (also calledadrenaline) is shown below. Show how the functional groups could be solvatedby water.

�HO

HO

CH¬ CH2¬NH2CH3

OH

¬

Cellular Reactions of WaterLearning ObjectiveUnderstand how the ionization of water affects the structures and actions of biomole-

cules in the cell.

Ionization of WaterIt may be easy to view water as just a background material without much dynamicactivity in the cell and organism; however, this is not a realistic picture. Althoughwater is not usually considered a substance with robust chemical reactivity, it doesdisplay features of selective reactivity. Several examples of water as a participant inbiochemical processes are presented in later chapters. Perhaps the most importantreaction of water is its reversible self-dissociation or ionization to generate thehydronium ion (H3O�) and the hydroxide ion (OH�):

H2O � H2O H3O� � OH�

Although not as correct, because free H� does not exist in aqueous solution, theequation is often abbreviated as

H2O H� � OH�

The extent to which this ionization reaction takes place is of special interest becauseit helps characterize the internal medium of cells. The more water molecules disso-ciated, the more ionic the medium. We can use the law of mass action to obtain aquantitative measure of the equilibrium point for the dissociation reaction:

Keq � [H�][OH�]/[H2O]

Keq represents the equilibrium constant for the reaction, and brackets for eachchemical entity indicate concentration units in moles per liter (M). If the Keq for theionization of pure water is determined from experimental measurements, it is possibleto calculate a quantity for [H�] and [OH�] and, therefore, to estimate the extent ofself-dissociation. Keq for pure water at 25°C has been determined to be 1.8 � 10�16 M.A value for [H2O] can be estimated by dividing the weight of water in one liter(1000 g) by the molecular weight of water (18).This yields [H2O] � 55.5 M.Therefore,

[H�][OH�] � Keq[H2O][H�][OH�] � (1.8 � 10�16 M)(55.5 M)[H�][OH�] � 1.0 � 10�14 M2

Because according to the chemical equation for dissociation H� and OH� must haveequal concentrations in pure water, then

Hydrogen ion concentrations expressed in exponential form are difficult to work with.A more useful terminology is pH, defined as the negative logarithm of the [H�](see Just in Time Review 2-1).

[H�] � [OH�] � 31.0 � 10�14 M2 � 1.0 � 10�7 M

∆

∆

2.3

ww

w.w

iley.com/c

ol l

e

g e / b o y e r

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2.1 J U S T I N T I M E R E V I E W

Using pH to Define Hydrogen Ion ConcentrationHydrogen ion concentrations calculated for ionization reactions are very small and in the exponential formare cumbersome for use in mathematical expressions. As we saw in the calculation for the ionization ofwater above, the [H�] for water is 1 � 10�7 M. In 1909, Søren Sørensen introduced the term pH to more con-veniently express [H�]. He defined pH as the negative logarithm of the hydrogen ion concentration:

pH ��log[H�]

The [H�] of 1 � 10�7 M becomes a pH of 7. The entire pH scale from 0 to 14 is defined in the figure below.The logarithmic feature of the pH scale is an important characteristic as it allows for its use over a very widerange of hydrogen ion concentrations. Note that each digit increase or decrease of pH represents a 10-foldchange in [H�]. A solution at pH 7 has 10 times greater [H�] than a solution at pH 8.

The determination of pH is one of the most frequent measurements made in the biochemistry laboratory.The structures of biomolecules and the efficiency of biochemical processes are dependent on [H�]. A pHmeasurement is usually taken by immersing a glass or plastic combination electrode into a solution andreading the pH directly from a digital readout or meter. The electrode is calibrated with buffer solutions ofknown pH values.

[H�], M

1.0 � 10

�14

1.0 � 10

�13

1.0 � 10

�12

1.0 � 10

�11

1.0 � 10

�10

1.0 � 10

�9

1.0 � 10

�8

1.0 � 10

�7

1.0 � 10

�6

1.0 � 10

�5

1.0 � 10

�4

1.0 � 10

�3

1.0 � 10

�2

1.0 � 10

�1

1.0

0 1 2 3 4 5 6 7pH 8 9 10 11 12 13 14

Acidic Neutral Basic

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As evident from the small size of [H�] (1 � 10�7 M) the extent of water self-dissociation is slight, but it does influence the ionic character of water-based matricesin cells and other aqueous-based biological fluids. The ionic environment and thepresence of H� and OH� promote the ionization of dissolved acidic and basicbiomolecules. The pH values for several fluids are compared in Figure 2.9.


1. Calculate the [H�] in each of the following fluids.a. Coffee, pH � 5.0

Solution: pH ��log[H�]log[H�] ��5.0

[H�] � 1 � 10�5 M

b. Urine, pH � 6.0c. Saliva, pH � 6.6d. Cola, pH � 2.8

2. Calculate the pH for each of the fluids below.a. Gastric juice, [H�] � 1 � 10�2 Mb. Acid rain, [H�] � 3.2 � 10�3 M

pH and pK

Calculations and measurements of pH as described above can be carried out for allaqueous solutions. The pH of a solution will depend little on the hydrogen ions gen-erated by the self-dissociation of water, but rather on the presence of other substances(acids or bases) that increase or decrease the H� concentration.Acids and bases arechemical substances that change the ionic properties of solutions.A useful definitionof an acid is a substance that releases a proton (H�). A base is a substance thataccepts a proton. An acid HA dissociates in aqueous solution accordingly:

HA H� � A�

Acid Base

∆

BasicAcidic

0 1 2 3 4 5 6 7 8

Neutral

9 10 11 12 13 14

Pancreatic juice 7.8–8.0

Seawater 7.0–7.5

Liver 7.4

Intracellular

fluids

Muscle 6.1

Saliva 6.4–6.9

Urine 5.0–8.0

Coffee 5.0

Tomato juice 4.3

Acid rain 3.5

Grapefruit juice 3.2

Cola soft drink 2.8

Lemon juice 2.3

Gastric juice 1.2–3.0

Blood plasma 7.4

Figure 2.9 The pH values of some substances. Note that many natural fluids, but not all, are grouped around aneutral pH of 7.

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This reaction describes the action of a whole range of acids from the strong mineralacids (HCl, H2SO4) to weak biological acids (acetic acid, lactic acid) to very weakacids In the equation, HA, as an acid, releases a H� to the solution. A� isconsidered a base because it accepts a proton in the reverse reaction. HA and A�

have a special relationship; they are a conjugate acid–conjugate base pair, sometimesshortened to acid–base conjugate pair.

The strength of an acid, or the measure of its tendency to release a proton, can beindicated from its dissociation constant. For a general acid HA, the acid dissociationconstant Ka is defined by

Ka � [H�][A�]/[HA]

Acids common in biochemistry have a wide range of Ka values. Hydrochloric acid(HCl) is a strong acid with an immeasurably large Ka; acetic acid is much weakerwith a Ka at 25°C of 1.74 � 10�5 M; and a very weak acid, has a Ka at 25°C of5.62 � 10�10 M. Note that the larger the Ka , the stronger the acid, and hence thegreater the dissociation. Several acids of biochemical importance are listed with theirdissociation constants in Table 2.4.

Dissociation constants written as exponentials are not convenient for everydayuse. These numbers are modified for easier use by the following definition:

pKa ��log Ka

The negative logarithm of the dissociation constant Ka is defined as pKa (pH has asimilar definition, the negative logarithm of [H�]). Like Ka, the pKa value is a quan-titative measure of acid strength. The most common range of pKa values for bio-chemical acids is from 2 to about 13 or 14.The smaller the value of pKa, the strongerthe acid. Note that this is the opposite of Ka, where a large value indicates a strongacid. Table 2.4 lists pKa values along with Ka values for several biochemical acids.

NH�4 ,

(NH�4 ).

Table 2.4Acids of biochemical importance

Acid Structurea Ka pKa

Formic acid HCOOH 1.78 � 10�4 3.75Acetic acid CH3COOH 1.74 � 10�5 4.76Pyruvic acid CH3COCOOH 3.16 � 10�3 2.50Lactic acid CH3CHOHCOOH 1.38 � 10�4 3.86Malic acid HOOC CH2 CHOH COOH (1) 3.98 � 10�4 3.40

(2) 5.50 � 10�6 5.26

Citric acid

(1) 8.14 � 10�4 3.09(2) 1.78 � 10�5 4.75(3) 3.92 � 10�6 5.41

Carbonic acid(1) 4.31 � 10�7 6.37(2) 5.62 � 10�11 10.26

Phosphoric acid(1) 7.25 � 10�3 2.14(2) 6.31 � 10�8 7.20(3) 3.98 � 10�13 12.40

Ammonium ion 5.62 � 10�10 9.25

a Acidic protons are in red.

H¬

Hƒ

N�¬H

ƒH

HO ¬

O‘Pƒ

OH

¬OH

HO ¬

O‘C ¬OH

CH2 CH2HOOC C

OH

COOH

COOH

¬¬¬

c02_036-062v2 8/26/05 8:54 AM Page 50



1. Write out the proton dissociation reaction(s) for each of the acids below. Identifythe conjugate acid and conjugate base for each reaction.a. HCl, hydrochloric acid

Solution:

b. CH3COOH, acetic acid

c. H4, ammonium ion

2. The Ka for lactic acid is 1.38 � 10�4. What is the value of the pKa?

Titration CurvesValues of pK (and dissociation constants) for acids can be determined experimentallyby the procedure of titration. This consists of adding, with a buret, incrementalamounts of a base to an acid sample dissolved in water and monitoring changes in thepH of the solution with a pH meter.A graph is constructed by plotting the pH on thevertical axis and the amount of base added on the horizontal axis. This yields a titra-tion curve from which the pKa can be determined (Figure 2.10).

Monoprotic Acids We will illustrate the construction of a titration curve using aceticacid, a weak monoprotic biochemical acid, and sodium hydroxide, a strong base.Theequation representing the chemistry of titration is

The data from the experiment are shown in Figure 2.10. The general shape of thecurve, which is obtained for all weak acids, reveals useful information about the acid.The information is both structural and quantitative. The beginning pH of the acidsolution, before addition of base, can be used to calculate the concentration of H� in

CH3COOH � NaOH CH3COO�Na� � H2OAcid Base Conjugate base

∆

N�

HCl Cl� � H�

Conjugatebase

Conjugateacid

Conjugate acid–base pair

∆

pH � pKa

Volume of NaOH added (mL)

pH

5 10 15 20 25

10

11

12

13

14

9

8

7

6

5

4

3

2

1

CH3COOH

[CH3COO�]�[CH3COOH]

CH3COO�

Figure 2.10 A titration curve. Theexperimental curve obtained from thetitration of acetic acid with sodiumhydroxide. The predominant structurefor acetic acid is shown in two pHranges, up to 4.76, and above 4.76.Note that at the inflection point the twoforms CH3COOH and CH3COO� arepresent in equal molar concentrations.The pKa (4.76) for the acid–baseconjugate pair is measured here.

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the solution.Acetic acid is relatively weak (there is only slight dissociation), so mostacetic acid molecules just prior to titration are in the form of CH3COOH.The curvechanges direction (inflects) at the midpoint. At the inflection point, 0.5 mol of basehas been added for each mole of acid present. Here, exactly one-half of the originalacid has undergone dissociation so equal amounts of two forms of acetic acid arepresent: the undissociated form CH3COOH (50%) and the conjugate base formCH3COO� (50%). The pH at this inflection point is equal to the pKa of acetic acid.(We will show mathematical proof for this statement in the next section.) At the endpoint (or equivalence point, where 1 mol of base has been added for each mole ofacid), equal amounts of acetic acid and sodium hydroxide have reacted, so essentiallyall molecules of acetic acid are dissociated to the conjugate base CH3COO�Na�.The titration experiment is valuable because it reveals the pKa value as well as theionic forms of acetic acid present at various pH values.

Polyprotic Acids Our discussion to this point has concentrated on a monoprotic acid,which is an acid molecule with only a single hydrogen atom that can dissociate. Manyacids of biochemical importance have two or more acidic protons; that is, they arepolyprotic. Some of these, listed in Table 2.4, include malic acid, citric acid, carbonicacid, and phosphoric acid. All of the acidic protons on a polyprotic acid do notdissociate at the same pKa , but are released in sequence at different pKa values.For example, phosphoric acid has three dissociable protons and, therefore, three pKavalues listed in Table 2.4, one value for each proton. The three-step ionization forphosphoric acid proceeds accordingly:

Note that the conjugate base for the first ionization reaction (H2 ) becomesthe proton donor (acid) for the second reaction, and so on. Titration curves forpolyprotic acids become more complex than those for monoprotic acids, but they areobtained in the same manner and they provide the same useful structural andquantitative information.


1. Draw the titration curve for each of the following acids. See Table 2.4 for struc-tures and pKa values. Assume they are titrated with NaOH as in Figure 2.10.a. Malic acidb. Phosphoric acid

The Henderson–Hasselbalch EquationIn the previous section, we were able to use titration data to define the ionic forms ofacetic acid present at three pH values: the beginning, midpoint, and the end of thetitration. In addition, it was possible to estimate the concentration of each species atthe three pH values.With the use of the Henderson–Hasselbalch equation, it is possi-ble to calculate the concentration of acid and conjugate base at all points of the titra-tion curve.The equation can be derived from the definition for dissociation constant:

pH � pKa � log([A�]/[HA])

where HA is the dissociating acid and A� is the conjugate base of the acid. If onethinks of this equation as containing four unknown but measurable quantities (pH,pKa, [A�], and [HA]), then, if three are determined, the fourth can be calculated.

PO 4�

H3PO4 H2PO4�

H�

HPO4�

H�

� 2�pKa(1) � 2.14 pKa(2) � 7.20PO4�

H�

3�pKa(3) � 12.4

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The Henderson–Hasselbalch equation may be used for many purposes—forcalculating [A�] and [HA] separately, or as a ratio [A�]/[HA], and for use in prepar-ing laboratory buffer solutions. Note that when [A�] � [HA] (at the inflection pointof the titration curve), the equation becomes

pH � pKa � log 1log 1 � 0

pH � pKa

When a solution contains equal concentrations of HA and A�, the pH of that solu-tion is equal to the pKa for the acid.


1. Lactic acid is present in our muscles after strenuous exercise. What is the valueof the ratio [A�]/[HA] in a solution of lactic acid at a pH of 7.2?

Solution: The titration of lactic acid proceeds as follows:

We need to calculate the ratio [lactate]/[lactic acid]. From Table 2.4, the pKa forlactic acid is 3.86. The Henderson–Hasselbalch equation for lactic acid is

pH � pKa � log([lactate]/[lactic acid])

Inserting known quantities gives

In a solution of lactic acid at pH 7.2, there are about 2200 molecules (or moles) ofthe conjugate base lactate for each molecule (or mole) of lactic acid. Therefore, thesolution at pH 7.2 consists of 99.9% lactate and 0.1% lactic acid.

2. What is the pH of a lactic acid solution that contains 60% lactate form and 40%lactic acid undissociated form?

Solution: The ionization of lactic acid proceeds as shown in question 1. TheHenderson–Hasselbalch equation becomes

pH � pKa � log(0.60/0.40)

The pKa for lactic acid is 3.86.

pH � 3.86 � log(0.60/0.40)pH � 3.86 � log 1.5pH � 3.86 � 0.18pH � 4.04

Notice from these two examples how the concentration ratio of conjugate base(lactate) to acid (lactic acid) varies with pH.

These examples illustrate two of many applications for the Henderson–Hasselbalch equation. In the next section, we discover yet another practical use ofthe equation.

[lactate]/[lactic acid] � 103.34 �2188

1

log([lactate]/[lactic acid]) � 7.2 � 3.86 � 3.34

7.2 � 3.86 � log([lactate]/[lactic acid])

OH

CH3CHCOOH � OH� CH3CHCOO� � H2O

Lactic acid(acid, HA)

OHLactate

(conjugate base, A�)

c02_036-062v2 8/26/05 8:54 AM Page 53



1. Write out all of the proton dissociation reactions for carbonic acid, H2CO3. UseTable 2.4 for pKa values.

2. What is the ratio of acetate to acetic acid in a solution at a pH of 5.0?

Drug Efficacy Depends on Water SolubilityThe study of water solvation, ionization, pH, and equilibrium constants may seem tobe theoretical and not of great significance in practical biochemistry. However, itmust be recognized that before we can use ingested energy molecules (carbohydrates,proteins, and fatty acids), nutrients, and vitamins, they must be dissolved in the mostabundant physiological solvent, water, and distributed throughout the body formetabolism.As we study these important biomolecules throughout this book, we willnote that most are water soluble because they contain functional groups that can

Table 2.5Structures and Actions of Common Drugs

Structurea Name Use/Action

Ibuprofen (Advil) Analgesic, anti-inflammatory agent,relief of pain

Sodium pentobarbital Sedative, induces sleep(Nembutal)

4-Androstene-3, 17-dione Anabolic steroid, enhances physical (“andro”) performance

Procaine-HCl (Novocain) Local anesthetic

Methadone-HCl Narcotic analgesic, treatment ofopiate addiction

a Note the number of charged groups and polar functional groups in the molecules. These groups enhance solubility in body fluids. Three of thedrugs are marketed and administered as the salts shown here.

O

“

φ

φ

CH2

¬¬

¬C ¬C ¬CH2CH3

¬

HCl�

H CH3

¬

¬ ¬

H3C ¬C¬¬¬N ¬CH3

φ =

Define:

�

≈√

O

“

¬

H2N C¬O¬CH2CH2

N

CH3CH2

� HCl

CH2CH3

OH3C

H3C

O

CH3CH2

CH3CH2CH2CH

CH3

ƒ

N

NH

O

O�Na�

(CH3)2CHCH2

COOH

CH3

ƒ

ƒCH

c02_036-062v2 8/26/05 8:54 AM Page 54

2.4 Buffer Systems 55

ionize (carboxyl and amino) to form highly soluble, charged salts and/or they havefunctional groups that are polar and can be hydrated. Polar functional groups includehydroxyl, carbonyl, amino, and amide groups.

The same kinds of chemical characteristics must be present in pharmaceuticalcompounds that are used to treat abnormal medical conditions. In order for anadministered drug to be effective, it must dissolve in aqueous body fluids and betransported to its target organs or cells. Therefore, not only must drugs be designedwith chemical functional groups that treat the malady, but they must also havechemical structures that lead to water solubility and the ability to be transportedacross cell membranes.

The chemical structures of several widely used drugs (legal and illicit) are shownin Table 2.5. Although the structures are complex, it is still possible to identify func-tional groups that enhance water solubility. Note that some have polar functionalgroups that can hydrogen bond and be solvated with water.

Some compounds have functional groups that can ionize to form charged salts.Many drugs have basic amino groups that react with HCl and other acids. This leadsto a positively charged salt that can easily dissolve in body fluids such as blood plasmaor cerebrospinal fluid. The pKa of these amino groups is about 10, so they would beprotonated and positively charged in blood plasma (pH about 7.4) and in the stomach(pH about 1–2). Carboxyl groups in drugs may be converted to soluble, negativelycharged groups by reaction with NaOH and other bases. The pKa of these carboxylgroups is about 4, so they would be ionized in blood plasma but not in the stomach.Some of the drugs shown in Table 2.5 are sold and administered as water-soluble salts.These include sodium pentobarbital, procaine hydrochloride, and methadonehydrochloride. The presence of hydrophobic groups (carbon skeletons and phenylrings) assists in transport of the molecules across cell membranes (Chapter 9).

Buffer SystemsLearning ObjectiveBe able to describe how buffers function to maintain a constant physiological pH.

Maintaining a Constant pHThe hydrogen ion concentration of intracellular and extracellular fluids must bemaintained within very narrow limits.A pH change in blood plasma of �0.2 to �0.4may result in serious damage to an organism or even death. A constant pH ensuresthat acidic and basic biomolecules are in the correct ionic state for proper function-ing. This is especially critical for enzymes and other proteins (Section 5.2) that aresensitive to pH changes. Metabolic reactions generate high concentrations of organicacids that would change fluid pH values if buffering agents were not present.Consider the following experiment. If 1.0 mL of 10 M HCl is added to 1.0 L of 0.15 MNaCl solution at pH 7.0, the pH would tumble to pH 2.0. If 1.0 mL of 10 M HCl isadded to 1 L of blood plasma, the pH would fall only from pH 7.4 to pH 7.2. Thereis nothing magical about blood. Its ability to maintain a constant pH is due to a het-erogeneous mixture of biomolecules that can act to neutralize added acids and bases.Blood and other biological fluids contain buffer systems: reagents that resist changesin pH when H� or OH� are added. Chemically, buffer systems contain acid–baseconjugate pairs. (See Window on Biochemistry 2-1.)

Effective Buffering RangeThe titration curve for acetic acid (see Figure 2.10) displays a region in which thepH changes little with addition of OH�. The center of this region is an inflectionpoint that coincides with the pKa of the titrated acid.The acetic acid solution at that

2.4

c02_036-062v2 8/26/05 8:54 AM Page 55


point (0.5 mol OH� per mol of acetic acid; pH � 4.76) shows the smallest change inpH and represents the most effective buffering range for the acetic acid–acetate con-jugate pair. Remember that at this point there is an equal concentration of aceticacid (50%) and acetate (50%).The solution has a large amount of acetic acid to reactwith and neutralize added base and a large amount of acetate (base) to react with andneutralize added acid. Therefore, the solution is buffered or protected against pHchanges caused by added acids or bases. In general, an acid–base conjugate pair ismost effective as a buffer system at a pH equal to its pKa. The effective bufferingrange for an acid–base conjugate pair can be estimated:

effective buffering range (pH) � pKa � 1

Acetic acid–acetate is an effective buffer in the pH range of 3.76–5.76. Phosphoricacid is an effective buffer in three pH ranges: (1) pH � 1.14–3.14, (2) pH � 6.20–8.20,and (3) pH � 11.40–13.40.

Laboratory BuffersBiochemical processes that occur in cells and tissues depend on the strict regulationof the hydrogen ion concentration provided by natural buffers. When in vitro studiesare done in the laboratory, artificial, buffered media are used to mimic the cell’s naturalenvironment. When cell components are isolated by using cell homogenates(see Window on Biochemistry 1-2), the constituent biomolecules are most stable whenmaintained in the natural pH range, which is usually 6–8. Sodium and potassium saltsof phosphoric acid are among the most widely used buffers in biochemical research.The phosphate acid–base conjugate pair, H3 – has a pK of about 7.2,which provides a useful pH range of 6.2–8.2 (see Table 2.4). Because phosphate is anormal constituent of the cell, its presence provides a natural environment. Other nat-ural buffering agents that may be used in the laboratory include the acid–baseconjugate pairs of acetic acid, citric acid, carbonic acid, and others listed in Table 2.4.

HPO42�,PO 4

�

Table 2.6Some synthetic buffers

Name (Abbreviation) pKa Useful pH Range Ionization Reactiona

N-(2-Acetamido)-2-aminoethanesulfonic 6.9 6.4–7.4 H2NCOCH2 H2CH2CH2acid (ACES) H2NCOCH2NHCH2CH2 � H�

3-(Cyclohexylamino)propanesulfonic 10.5 10.0–11.0acid (CHAPS)

N-(2-Hydroxyethyl)piperazine- 7.5 7.0–8.0N�-2-ethanesulfonic acid (HEPES)

Tris(hydroxymethyl)aminomethane (TRIS) 8.3 7.5–9.0 (HOCH2)3C H3 (HOCH2)3CNH2 � H�

a Each reaction shows the two predominant forms (acid and base) present in the useful pH range.

∆N�

NCH2CH2SO3 H��HOCH2CH2N �

NCH2CH2SO3 ∆�HOCH2CH2N�

H

NHCH2CH2CH2SO3 H��

�NH2CH2CH2CH2SO 3 ∆�

SO 3�

∆SO 3�N

�

Sometimes natural buffers like phosphate and citrate interfere with laboratory invitro studies and it is necessary to use synthetic buffers. N. E. Good at Michigan StateUniversity synthesized and tested several acid salts that are now widely used asbuffering agents (Table 2.6). The use of the synthetic buffer tris(hydroxymethyl)aminomethane (TRIS) is now probably greater than that of phosphate salts.

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2.4 Buffer Systems 57

2 - 1

Window on BiochemistryBuffers in Blood and Other Cellular Fluids

The major buffer system of blood and other extracellularfluids is the carbonic acid–bicarbonate conjugate pair:

A base added to blood would be neutralized by thefollowing reaction:

The addition of an acidic substance to blood also resultsin neutralization:

These reactions (see figure) illustrate how blood is pro-tected against pH changes. The actual pH of blood(pH � 7.4) is at the upper limit of the buffering range ofcarbonic acid–bicarbonate (6.4 � 1 � 5.4–7.4) and maynot seem as efficient as desired. This inefficiency is reme-died by a reserve supply of gaseous CO2 in the lungs, whichcan replenish H2CO3 in the blood by the following series ofequilibrium reactions:

The reactions also work in reverse. A major product ofmetabolism, H� is removed from cells by the blood plasma.

Lungs Blood

H2CO3(aq) ∆ HCO 3(aq)� � H(aq)

�

CO2(aq) � H2O(l) ∆ H2CO3(aq)

CO2(g) ∆ CO2(aq)

HCO 3� � H� ∆ H2CO3

H2CO3 � OH� ∆ HCO 3� � H2O

H2CO3 3HCO� � H�

BicarbonateCarbonic acid

pKa � 6.4

It is neutralized by reaction with and leads to even-tual release of CO2(g) , which is exhaled from the lungs.

The carbonic acid–bicarbonate conjugate pair is themost important buffer system in biological fluids; however,it is not the only one. A diverse array of amino acids,peptides, and proteins with ionizable functional groups( COOH and ) assist in buffering. A major proteinconstituent of blood, hemoglobin also serves as a buffer-ing agent. Details of the function of hemoglobin are dis-cussed in Chapter 4.

Medical conditions caused by changes in blood pH areacidosis and alkalosis. An increase in the [H�] of blood(acidosis) may have causes that are of metabolic or respira-tory origin. Metabolic acidosis occurs in individuals withuntreated diabetes or in those on starvation diets or on high-protein, low-fat diets. All of these metabolic conditions leadto ketosis, the excessive generation of ketone bodies, whichare acidic and increase the [H�] of blood. Respiratoryacidosis is caused by a change in [CO2] that is often asymptom of pulmonary problems associated with emphyse-ma or asthma. Untreated acidosis leads to coma andeventually death.

An increase in blood pH (alkalosis) also has metabolic orrespiratory origins. Clinical administration of the saltsof metabolic acids (sodium lactate or sodium bicarbonate)in excessive amounts or cases of severe vomiting causemetabolic alkalosis. Respiratory alkalosis is induced byhyperventilation (heavy breathing), which may result fromhysteria, anxiety, or altitude sickness.

NH3�¬¬

HCO3�

Blood

Artery

CO2

O2

CO2

CO2

H2CO3

H� � HCO 3

H2O H2O

VeinVein

O2

O2

Lungs Outsideair

CO2

O2

CO2

O2

Active tissuesand organs

Fuel � O2

H2O � CO2

�

Control of blood pH by the carbonic acid–bicarbonateconjugate pair. There is some CO2 in arterial blood andsome O2 in venous blood; not all of the CO2 is exhaledas blood flows through the lungs. CO2 in the lungs is inequilibrium with CO2 in the blood. A high concentrationof CO2 in the lungs leads to respiratory acidosis while alow CO2 concentration causes respiratory alkalosis.

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2 - 1

Biochemistry in the ClinicAspirin Overdose and Salicylate Poisoning

The first day on your new job as a physician in the emer-gency department at Riverside County Hospital has beencalm—an ankle sprain, a cut finger, and a broken rib—noth-ing life threatening. Your morning coffee break is interruptedby an ambulance bringing in a mother with her 4-year-old,semiconscious son, Sam. The mother reported what hadhappened: two hours earlier, Sam had complained of stom-achache, had vomited, was hyperactive, began breathingheavily and fast, developed a fever, became drowsy, andfinally collapsed. Sam’s mother also brought in what shethought had caused the accident, an empty bottle that oncecontained about 30 orange-flavored, children’s aspirintablets. Careful examination of the patient showed classicaspirin-overdose symptoms—a fever, rapid heart rate, dis-orientation, and hyperventilation. You know that ingestionof aspirin, acetylsalicylic acid (or any acid), disturbsacid–base balance (see Window on Biochemistry 2–1), soyou order measurements of Sam’s blood gases, pH, bicar-bonate, and salicylate:

Sam (after 2 hrs) Normal

Partial pressure of CO2 20 mm Hg 35–45 mm Hg

Partial pressure of O2 115 mm Hg 75–100 mm Hg

pH 7.45 7.35–7.4518 mM 24–28 mM

Salicylate 75 mg/100 mL 0 mg/100 mL

While you are awaiting the test results, your medical schoollectures on aspirin overdose flash through your mind. Aspirinin the stomach is hydrolyzed by acid and enzymes to salicy-late, another acid as strong as aspirin. Salicylate is an uncou-pling agent that increases respiration—it enhances oxygenuptake and thus increases the rate of aerobic metabolism(see Section 17.3). The energy from metabolism is not storedand used in the normal way (to make ATP), but is channeledto the production of heat (fever). The hyperventilation ini-tially induces respiratory alkalosis, a rise in blood pH anddecline in carbon dioxide. With the increased production ofmetabolic acids (lactate, pyruvate), the respiratory alkalosisis quickly changed to metabolic acidosis (lower blood pH).The presence of salicylate in Sam’s blood confirms your sus-picions about aspirin overdose and you begin treatment:

1. Perform a gastric lavage (stomach pumping) with asodium bicarbonate solution. The basic solution speedselimination of the acid salicylate in the urine and helpsneutralize the metabolic acidosis condition.

HCO 3�

2. Administer activated charcoal to bind the aspirin and sal-icylate and to enhance their elimination before they areabsorbed into cells.

3. Use bicarbonate infusion to maintain urine pH above 7.5to hasten the elimination of the acids.

4. Let Sam rest, but administer fluids (juices, milk, etc.)and assess response to treatment by monitoring changesin metabolic indices (see table). Concentrations of CO2,O2, and and pH should move to normal values inabout 24 hours.

5. If symptoms do not improve after several hours, use akidney machine to perform hemodialysis to clear theblood of aspirin and salicylate.

The prognosis (probable outcome) for a patient like Samis usually very good unless the level of aspirin is very high(above 150–300 mg/kg). For levels below this range, themortality rates are less than 2%. For levels above the range,the mortality rates can be as high as 25%.

Study Questions

1. Aspirin in the stomach is hydrolyzed under the acidicconditions (pH about 2) to salicylic acid:

HCO3�


COOH

OCCH3

O

“

COOH

OH

+CH3COOHH�, H2O

Acetylsalicylicacid

Salicylicacid

a. Write the reaction for the proton ionization of the car-boxyl group in salicylic acid.

b. Use the Henderson–Hasselbalch equation to calculatethe ratio of salicylate to salicylic acid in Sam’s stom-ach (pH about 2.0) and in gastric lavage sodium bicar-bonate solution (pH about 8.5). The pKa for thecarboxyl group is about 3.0.

c. Explain why salicylate is more water soluble than sal-icylic acid.

ReferencesCornely, K., 1999. Acute aspirin overdose: Relationship to the blood

buffering system. In Cases in Biochemistry, pp. 1–2. New Jersey:John Wiley & Sons.

(© FeliciaMartinez/Photo Edit.)

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Study Exercises 59

SUMMARY

■ Water, a nonlinear, polar molecule, serves at least threefunctions in the cell: It is an effective solvent, it is a reactantmolecule, and it is a temperature buffer.As a solvent, wateris able to dissolve biomolecules that are ionic and polar.■ The most important reaction of water is its reversibleionization to generate the hydronium ion (H3O� or theproton, H�) and the hydroxide ion (OH�).The extent ofionization is quantified by the pH scale (pH ��log[H�]).■ The strength of an acid is defined by its pKa, the nega-tive log of its dissociation constant.The pKa for an acid isequivalent to the pH at which there is an equal concen-tration of an acid and its conjugate base.■ Blood and other cellular fluids are maintained at aconstant pH by natural buffer systems including thecarbonic acid–bicarbonate conjugate pair and the hydro-genphosphate–dihydrogenphosphate conjugate pair.

STUDY EXERCISES

Understanding Terms2.1 Define each of the following terms and give a specificexample, if appropriate.

a. An acidb. A basec. pHd. pKe. Henderson–Hasselbalch

equationf. Noncovalent interactionsg. A hydrophobic moleculeh. A hydrogen bond

i. Hydrophobic interactionsj. Buffer systemk. Acid–base conjugate pairl. Metabolic acidosism. Micellen. Good buffero. Amphiphilic moleculesp. Ionization of acidsq. Molecular recognition

Reviewing Concepts2.2 Calculate the hydrogen ion concentration [H�] in each ofthe following solutions.

2.5 Write acid dissociation reactions for each of the followingbiochemically important molecules. Show the ionization of allacidic protons.

a. HClb. CH3COOH (acetic acid)c. (ammonium ion)d. CH3(CH2)13CH2COOH (palmitic acid)

e.

f. H3PO4 (phosphoric acid)g. H2Oh. H2CO3 (carbonic acid)

2.6 Arrange the following natural solutions in decreasing orderof acidity.

Gastric juiceBloodAcid rainColaCoffee

Hint: See Figure 2.9.

2.7 The nucleotide bases shown here form hydrogen bondsbetween the two strands of the double helix of DNA. Circlethose atoms that may become involved in hydrogen bonding.Distinguish between acceptor atoms and donor atoms.

2.8 Determine whether each of the statements about nonco-valent interactions is true or false. If false, change the statementso it is true.a. Ionic bonds are the result of electrostatic attraction between

two ionized functional groups of opposite charge.b. Hydrogen bonds result from interaction of an anion with a

hydrogen atom.c. Hydrophobic interactions are electrostatic attractions

between nonpolar functional groups and water.d. H� and OH� interact together by ionic bonding to form

water.e. Hydrophobic interactions are important in the formation

of micelles when the detergent sodium dodecanoate,CH3(CH2)10COO�Na�, is added to water.

2.9 Identify the type of interaction that holds each of thefollowing atoms or groups of atoms together. Choose fromdipole–dipole, ion–ion, or dipole–ion.

NH2CH3

N NH

N N

H

O

ON

H

N

Adenine Thymine

H3N

R

CHCOOH� (an amino acid)

NH 4�

a. NaClb. Na�(H2O)n

c. CH3COO�Na�

d.

e. Cl�RNH 3�

R O

H

H O R

a. Gastric juice, pH � 1.80b. Blood plasma, pH � 7.40c. Cow’s milk, pH � 6.6

d. Tomato juice, pH � 4.3e. Urine, pH � 5.0f. Maple tree sap, pH � 7.1

2.3 The amino acid glycine, NCH2COO�, has pKa valuesof 2.4 and 9.8. Estimate the effective buffer range(s) for glycine.

2.4 Aspirin (acetylsalicylic acid) has the following structure:

a. Draw the ionic structure for the predominant form of aspirinas it would exist in blood plasma.

b. Draw the ionic structure for the predominant form of aspirinas it would exist in gastric juice.

Hint: The pKa for aspirin is 3.5.

OCCH3

O

COOH

H 3�

c02_036-062v2 8/26/05 8:54 AM Page 59


2.10 For each pair of molecules listed below, determine whichone is more polar than the other.

2.16 Determine whether each of the statements is true or false.If false, rewrite it so it is true.a. Noncovalent bonds are usually more easily broken then

covalent bonds.b. The strength of a typical noncovalent bond is usually at least

300 kJ/mol.c. Noncovalent interactions often occur between only certain

molecules.d. Noncovalent bonds reversibly break and re-form at room

temperature.

2.17 Define and contrast four types of noncovalent bonds thatpair biomolecules for molecular recognition.

2.18 Describe the differences between hydrogen bonds andionic bonds.

2.19 Predict which of the following compounds are solublein water.a. CH3CH2 OHb. CH3(CH2)10CH2 OHc. CH3CH2COOH

d.

e.

f.

g.

2.20 Which of the following molecules will form micelles inwater?a. CH3(CH2)10CH2 Cl�

b. CH3(CH2)10CH2COO�Na�

c.

d.

Hint: Look for polar head and nonpolar tail.

CH3(CH2)12CH2

O

C NH2

CH

CH2 O

O

C

P

(CH2)10 CH3

C (CH2)8 CH3O

CH2

OO

O�

O CH2 CH2 NH 3O �

NH 3�

HOCholesterol

H3C

CH3

OH

CH2CHCH2

OHOHGlycerol

H 3NCH2COO��

Glycine

OCH2OH

OH

OHHO

OHGlucose

¬¬

a. H2O, CH3OHb. H2O, CH3COOHc. CH3(CH2)3CH3, CH3CH2OH

d.

e.

f.

g. H2NCNH2, CH3CNH2

O O

CH2CH3, CH3CH3

NH2

CH3CNH2, CH3CH

O O

CH2CH2, CH3CH2OH

HO OH

2.11 What is the molar concentration of the HOH species inpure water?

2.12 Write the molecular formula for the conjugate acid of eachof the following bases.

a. OH�

b.c. NCH2COO�

d. CH3COO�

e. H2PO 4�

H 3�

HCO 3�

2.13 For each pair of molecules listed below, determine whichone is the less polar (more nonpolar).a.

b.

c.

d.

e. HOOCCH2CH2COOH, CH3CH2CH2COOH

2.14 Write structures for the conjugate base of each of thefollowing acids.

CH3CH2COOH, CH3CHCOOH

OH

H 3NCHCOO�, H 3NCHCOO�

CH3

CH3 CH3

CH

��

CH2CH3, H2O

OH

CH3CH3, CH2CH3

OH

a. H2Ob. NCH2COOHc. CH3(CH2)10COOHd.

e.

�NH3�NH3

CH2(CH2)10CH2

HCO 3�

H 3�

Solving Problems2.15 Identify the type of bonding between each pair of atomsand molecules. In examples with more than one type, name thetype of bonding indicated by the arrow.

a. Na�Cl�

b. H O H

c.

d.

Hint: Select from covalent, hydrogen, and ionic bonding.

H2N C NH

(i)

O

O

H

H

H(ii)

H O H O

H H

¬¬

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Study Exercises 61

2.21 You need to prepare a buffer for use at pH 4.0.Which sub-stance would be most effective?

a. Lactic acid b. Acetic acid c. Phosphoric acid

Hint: Check Table 2.4 for pKa values.

2.22 The acid–base conjugate pair H2CO3– maintainsthe pH of blood plasma at 7.4. What is the ratio of bicarbonateto carbonic acid in blood?

2.23 Can you think of a possible emergency room treatmentfor a patient in a state of acidosis?

2.24 A buffer solution was prepared by mixing 0.05 mol of thesodium salt of the amino acid alanine with 0.1 mol of free alaninein water. The final volume is 1 L. The equilibrium reaction is

What is the pH of the final solution?

Hint: Use the Henderson–Hasselbalch equation.

2.25 The pH of normal rainwater is approximately 5.6. Theslight acidity compared to pure water is caused by dissolvedCO2. In some regions of the world, the acidity of rain hasincreased to a pH of approximately 3.5. This is caused by thepresence of polluted air containing SO2, SO3, and NO2. Theseoxides react with rainwater to form acids:

SO2 � H2O H2SO3

SO3 � H2O H2SO4

2 NO2 � H2O HNO3 � HNO2

Calculate the ratio of bicarbonate to carbonic acid in normalrainwater and in acid rain (pH � 3.5).

2.26 Excess “stomach acid,” a result of our hectic, fast-pacedlifestyles, is often treated with antacids. From your knowledge ofacid–base chemistry, predict which of the following compoundswould be an ingredient in over-the-counter antacids.

¡¡¡

�NH3

CH3CHCOO� CH3CHCOO� � H�

Alanine

pKa � 9.9

NH2Salt of alanine

HCO 3�

a. NaHCO3

b. Ascorbic acid (vitamin C)c. Mg(OH)2

d. CH3COOH (acetic acid)

e. NaAl(OH)2CO3

f. Aspiring. Lemon juiceh. CaCO3

2.27 The aspirin product Bufferin contains magnesium car-bonate, MgCO3. What is the purpose of this ingredient?

2.28 Predict which of the compounds below can form hydrogenbonds with water. Show examples of hydrogen bonding for eachone you select.

a.

b.

c.

d. H3N CHCOO��

CH2

SHCysteine

H2N C

O

NH2

Urea

CH3CH2CH3

Propane

CH3CH2OHEthanol

2.29 We will discover in the next chapter that amino acids aredi- or triprotic acids. Classify each of the amino acids shownbelow as a diprotic or a triprotic acid and write all dissociationreactions showing removal of all acidic protons.

a.

b.

c.

d. CHCOOH

CH2

NH2

C O

Asparagine

H 3N�

CHCOOH

(CH2)4

NH3

Lysine�

H 3N�

CHCOOH

CH2

COOHAspartic acid

H 3N�

H 3N CHCOOH

CH3Alanine

�

2.30 Which of the compounds shown below would function assoaps or detergents?a. CH3(CH2)12CH3b. CH3(CH2)9CH2COO�K�

c. CH3(CH2)10CH2OH

d.

2.31 Shown below is the structure of a dipeptide formed withthe amino acids serine and cysteine. Circle all atoms that maybecome involved in hydrogen bonding with H2O. Distinguishbetween those atoms that are hydrogen bond donors andhydrogen bond acceptors.

2.32 What is the numerical value of the ratio [lactate]/[lacticacid] in a solution of lactic acid at a pH of 5.0?

Hint: Use the Henderson–Hasselbalch equation:pH � pKa � log([lactate]/[lactic acid]).

2.33 What is the pH of a lactic acid solution that contains 75%lactate form and 25% lactic acid?

2.34 Phosphate buffers are often used in biochemical researchbecause they can be prepared in the physiological pH range of7 and because phosphates are naturally occurring biomolecules.How many moles of monobasic sodium phosphate (NaH2PO4 ·H2O) and dibasic sodium phosphate (Na2HPO4 · 7H2O) mustbe added to a liter of water to prepare a 0.5 M phosphate bufferof pH 7.0?

2.35 When compared to the hydrides of other nonmetallicelements (N, C, S), water has:a. The lowest boiling pointb. The lowest freezing pointc. The highest boiling point and the lowest freezing pointd. The highest viscosity

CHCOOHCHCN

CH2

OH

CH2H

SH

O

Ser-Cys

H3N�

CH3(CH2)10CH2OSO�Na�

O�Na�

O

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2.36 Which of the following statements about water as asolvent is not true?a. Ions in water are solvated by formation of hydration shells.b. Uncharged but polar compounds form favorable

dipole–dipole interactions with water.c. Nonpolar compounds like decane are soluble because water

can form a hydration cage around the decane molecules.

2.37 Compare the [H�] of gastric juice (assume pH � 1.4) tothat of blood plasma (assume pH � 7.4).a. The [H�] of gastric juice is 6 times higher (7.4 � 1.4).b. The [H�] of gastric juice is 106 times higher.c. The [H�] of blood plasma is 6 times higher than gastric juice.d. The [H�] of blood plasma is 106 times higher.

2.38 You prepare a sodium phosphate buffer by mixing 100 mLof 0.1 M Na2HPO4 with 100 mL of 0.1 M NaH2PO4. The pH ofthe final solution is 7.20. What is the pKa?

2.39 Which of the following acids would make the most effec-tive buffer at pH � 6? See Tables 2.4 and 2.6 for pKa values andstructures.a. Acetic acidb. Lactic acidc. Histidine (Ka � 9.12 � 10�7 M)d. TRIS (Ka � 8.08 � 10�9 M)

2.40 Why are the carboxyl groups in drugs and other com-pounds ionized in blood plasma but not ionized in the stomach?

Writing Biochemistry2.41 Your roommate does not have time to read this chapterbefore the next class exam. She has offered to pay you a hand-some reward for writing a brief summary for her to review.Write a 100-word summary of Chapter 2 that contains a reviewof the most important concepts. Use the chapter summary as aguide, but be creative in your approach and style.

2.42 The Sixty-Second Paper. Your instructor will probablyspend about one lecture on this chapter. Immediately afterthe lecture, use 1–2 minutes to write a paper answering the

following questions: (1) What were the central points outlinedin the lecture? (2) What concepts were confusing?

2.43 Although methane and water have similar molecularmasses (16 and 18 daltons, respectively), they have quite differ-ent properties. Explain why methane is unable to form hydro-gen bonds with itself or with other molecules.

2.44 Your roommate, an English major, makes heavy use ofantacid tablets especially just before her creative writing papersare due in class. She has asked you to explain how the tabletswork to neutralize stomach acid.Write the explanation in a 100-word paragraph.

FURTHER READINGBall, P., 2000. Life’s Matrix: A Biography of Water. New York: Farrar,

Straus & Giroux.Boyer, R., 2000. pH, buffers, electrodes, and biosensors. In Modern

Experimental Biochemistry, 3rd ed., pp. 29–41. San Francisco:Benjamin/Cummings.

Colson, S. and Dunning, T. Jr., 1994. The structure of nature’s solvent:water. Science 265:43–44.

Daviss, B., 2004. Structured water is changing models. The Scientist,Nov. 8:14–15.

Gerstein, M. and Levitt, M., 1998. Simulating water and the moleculesof life. Sci. Am. 279(11):101–105.

Kegley, S. and Andrews, J., 1997. The Chemistry of Water. Sausalito, CA:University Science Books.

Mattos, C., 2002. Protein–water interactions in a dynamic world. TrendsBiochem. Sci. 27:203–208.

Norby, J., 2000. The origin and the meaning of the little p in pH. TrendsBiochem. Sci. 25:36–37.

Po, H. and Senozan, N., 2001. The Henderson–Hasselbalch equation:Its history and limitations. J. Chem. Educ. 78:1499–1503.

Pollack, G., 2001. Cells, Gels, and the Engines of Life. Seattle,WA: Ebner& Sons.

Rawls, R., 2001.Watching water line dance. Chem. Eng. News, Dec. 17:14Wettlaufer, J. and Dash, J., 2000. Melting below zero. Sci. Am. 282(2):

50–53.Zeuthen, T., 2001. How water molecules pass through aquaporins,

Trends Biochem. Sci. 26:77–79.

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