Cambridge University Press is part of the University of ...assets.cambridge.org/052179/9651/sample/0521799651WS.pdf · Second edition 1981 Third edition 1986 Third edition ﬁrst

Cambridge University Press is part of the University of Cambridge. It is a charitable enterprise dedicated to printing and publishing for theadvancement of knowledge throughout the world. In recognition ofthe fact that there are many parts of the world where people havedifficulty in purchasing their own copies of the best internationaltextbooks, Cambridge University Press is now publishing a series ofLow Price Editions. These are reprints of internationally respectedbooks from Cambridge University Press, specially printed and pricedfor the benefit of students and teachers in selected countries.

In this new, 5th edition of a highly popular text, undergraduatestudents are introduced to all the basic experimental techniquesroutinely used in practical biochemistry today. Emphasis is given totechniques encountered in practical classes, with the principles andtheories behind these explained in detail to aid understanding. As afurther aid to students, essential calculations and worked answersappear at the end of each chapter. ‘Key terms to understand’ are alsoincluded to help students thoroughly review each topic.

No contemporary book on modern biochemical techniques wouldbe complete without chapters on molecular biology, recombinant DNAtechnology, genetic analysis and biomolecular interactions, and thesetopics have been extensively covered in this new edition.

This is essential reading for all bioscience undergraduate studentsand pre-clinical medical students for whom practical biochemistry,molecular biology and immunology form part of the syllabus.

Fifth Edition

PracticalBiochemistryPrinciples and techniques

Edited byKeith Wilson and John Walker

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First and second editions © Bryan Williams and Keith Wilson 1975, 1981Third edition © Keith Wilson and Kenneth H. Goulding 1986Fourth edition © Cambridge University Press 1994Fifth edition © Cambridge University Press 2000

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First published by Edward Arnold 1975as A biologist's guide to principles and techniques of practical biochemistrySecond edition 1981 Third edition 1986Third edition first published by Cambridge University Press 1992Reprinted 1993Fourth edition published by Cambridge University Press 1994 asPrinciples and techniques of practical biochemistryReprinted 1995, 1997Low price edition 1995Fifth edition 2000Low price edition 2000

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Contents

Preface to the fi◊h edition xiiiList of contributors xviList of abbreviations xviii

1 General principles of biochemical investigationsi . s impkins

1.1 The nature of biochemistry 1

1.2 Bioenergetics 4

1.3 Methods for investigating metabolism 21

1.4 Practical considerations 28

1.5 In vivo models 42

1.6 In vitro models 44

1.7 Microscopy 66

1.8 Key terms 75

1.9 Calculations 76

1.10 Suggestions for further reading 78

2 Molecular biology and basic techniques 80r. rapley

2.1 Introduction 80

2.2 Components and primary structure of nucleic acids 80

2.3 Genes and genome complexity 87

2.4 The nature of the genetic code 90

2.5 Cellular location of nucleic acids 90

2.6 The cellular functions of DNA 93

2.7 The manipulation of nucleic acids: basic tools and techniques 103

2.8 Isolation and separation of nucleic acids 105

2.9 Restriction mapping of DNA fragments 110

v

2.10 Nucleic acid blotting methods 111

2.11 Gene probe derivation 113

2.12 Labelling DNA gene probe molecules 114

2.13 The polymerase chain reaction 116

2.14 Nucleotide sequencing of DNA 125

2.15 Bioinformatics and the Internet 131

2.16 Key terms 134

2.17 Calculations 135


3 Molecular cloning and gene analysis 138r. rapley

3.1 Introduction 138

3.2 Constructing gene libraries 138

3.3 Cloning vectors 148

3.4 Hybridisation and gene probes 167

3.5 Screening gene libraries 167

3.6 Applications of gene cloning 171

3.7 Expression of foreign genes 178

3.8 Analysing genes and gene expression 184

3.9 Analysing whole genomes 196

3.10 Molecular biotechnology and applications 202

3.11 Key terms 204


4 Immunochemical techniques 206r. thorpe and s. thorpe


4.2 Production of antibodies 211

4.3 Purification and fragmentation of immunoglobulins 222

4.4 Immunoprecipitation 229

4.5 Labelling antibodies 234

4.6 Immunoblotting 240

4.7 Immunoassays 244

4.8 Immunohisto/cytochemistry 254

4.9 A∞nity and avidity 260

4.10 Immunochemical use of surface plasmon resonance 260

4.11 Key terms 261

vi Contents

4.12 Calculation 262


5 Centrifugation techniques 263a. griffiths


5.2 Basic principles of sedimentation 264

5.3 Centrifuges and their uses 271

5.4 Design and care of preparative rotors 276

5.5 Sample containers 284

5.6 Separation methods in preparative ultracentrifuges 285

5.7 Performing density gradient separations 290

5.8 Selection, e∞ciency and applications of preparative rotors 296

5.9 Analysis of subcellular fractions 300

5.10 Some applications of the analytical ultracentrifuge 302

5.11 Safety aspects in the use of centrifuges 305

5.12 Key terms 306



6 Protein structure, purification and characterisation 312j . walker

6.1 Ionic properties of amino acids and proteins 312

6.2 Protein structure 316

6.3 Protein purification 318

6.4 Protein structure determination 338

6.5 Key terms 353



7 Biomolecular interactions: I Enzymes 357k. wilson

7.1 Receptor–ligand binding 357

7.2 Enzymes: characteristics and nomenclature 358

7.3 Enzyme steady-state kinetics 359

7.4 Enzyme assays 378

7.5 Substrate assays 385

7.6 Enzyme pre-steady-state kinetics 386

vii Contents

7.7 Enzyme active sites and catalytic mechanisms 389

7.8 Immobilised enzymes 394

7.9 Cellular control of metabolic activity 396

7.10 Key terms 400



8 Biomolecular interactions: II Cell surface receptors and transporters 403k. wilson

8.1 Cell surface receptor classification 403

8.2 Quantitative aspects of receptor–ligand binding 406

8.3 Receptor structures 417

8.4 Mechanisms of signal transduction 420

8.5 Signal amplification 430

8.6 Key terms 432

8.7 Membrane transport processes 432

8.8 Physical di≠usion 433

8.9 Facilitated transport 435

8.10 Active transport and ion channels 438

8.11 Receptor-mediated endocytosis 445

8.12 Key terms 449



9 Spectroscopic techniques: I Atomic and molecular electronic spectroscopy 453d. b . gordon


9.2 g-Ray spectroscopy and g-ray resonance spectroscopy 456

9.3 X-ray spectroscopy 458

9.4 Ultraviolet and visible light spectroscopy 459

9.5 Spectrofluorimetry 471

9.6 Circular dichroism spectroscopy 479

9.7 Turbidimetry and nephelometry 482

9.8 Luminometry 483

9.9 Atomic spectroscopy 485

9.10 Lasers 490

viii Contents

9.11 Key terms 491



10 Spectroscopic techniques: II Vibrational spectroscopy and electron and

nuclear spin orientation in magnetic fields 498d. b . gordon


10.2 Infrared and Raman spectroscopy 498

10.3 Electron spin resonance spectroscopy 501

10.4 Nuclear magnetic resonance spectroscopy 508

10.5 Key terms 525


11 Mass spectrometric techniques 527d. b . gordon


11.2 The mass spectrometer 527

11.3 Electron impact ionisation 529

11.4 Chemical ionisation 538

11.5 Field ionisation 539

11.6 Ion desorption methods 539

11.7 Ion evaporation methods 551

11.8 Analysers 555

11.9 Detectors 564

11.10 Tandem mass spectrometry 566

11.11 Key terms 573



12 Electrophoretic techniques 580j . m. walker

12.1 General principles 580

12.2 Support media 584

12.3 Electrophoresis of proteins 588

12.4 Electrophoresis of nucleic acids 607

12.5 Capillary electrophoresis 612

12.6 Key terms 617

ix Contents

12.7 Calculation 618


13 Chromatographic techniques 619k. wilson


13.2 Chromatography theory and practice 623

13.3 Low pressure column chromatography 631

13.4 High performance liquid chromatography 637

13.5 Adsorption chromatography 647

13.6 Partition chromatography 649

13.7 Ion-exchange chromatography 656

13.8 Molecular exclusion (permeation) chromatography 661

13.9 A∞nity chromatography 665

13.10 Gas–liquid chromatography 672

13.11 Thin-layer (planar) chromatography 678

13.12 Selection of a chromatographic system 681

13.13 Key terms 682



14 Radioisotope techniques 687r. j . slater

14.1 The nature of radioactivity 687

14.2 Detection and measurement of radioactivity 693

14.3 Other practical aspects of counting radioactivity and analysis of data 713

14.4 Inherent advantages and restrictions of radiotracer experiments 717

14.5 Safety aspects 718

14.6 Applications of radioisotopes in the biological sciences 721

14.7 Key terms 726



15 Electrochemical techniques 729p. k. robinson


15.2 Principles of electrochemical techniques 734

15.3 Redox reactions 742

x Contents

15.4 The pH electrode 745

15.5 Ion-selective and gas-sensing electrodes 748

15.6 The Clark oxygen electrode 750

15.7 Electrochemical detectors for HPLC 757

15.8 Biosensors 760

15.9 Key terms 768



Index 771

xi Contents

Chapter 1

General principles of biochemicalinvestigations

1.1 THE NATURE OF BIOCHEMISTRY

Biochemistry is an interdisciplinary science that integrates systematically theprinciples of mathematics, physics and chemistry to attempt to explain the dis-tinctive characteristics of life processes in terms of structure–function correla-tions. Advances in biochemistry have therefore largely exploited principles andtechniques first applied in the physical sciences. Biochemistry is no longer simplyan academic discipline but an applied science in which scientific discovery isexpected to lead to material benefits. In recent years the fusion of biochemistry,cell biology and microbiology to form molecular biology has led to spectacularadvances in the understanding and control of biological processes in medicine,agriculture, pharmaceutics and the food and drinks industry. This justifies tomany the widely held belief that biotechnology will become the pre-eminentindustry of the new millennium.

Biochemistry is quintessentially both analytical and quantitative in usingmodel biological systems of di≠erent physiological complexities to explain cause-and-e≠ect relationships in molecular terms. Analysis means literally ‘getting tothe bottom of things’, i.e. taking to pieces. Analysis is useful only, however, whencombined with synthesis, the piecing together, through interpretation andextrapolation of observations made on the disassembled parts, into the workingwhole. Analysis and synthesis therefore in combination define the boundariesbetween systems and surroundings, i.e. what components are part of the systemand consequently a≠ect the system directly in contrast to external factors thata≠ect it only indirectly. Repeated investigations lead in many cases to the defini-tion of a system.

In analytical biochemistry, for example, experimental models are subjected firstto qualitative analysis, in which predominantly heterogeneous biologicalmaterial is subjected to disruption techniques and the constituent parts separated,concentrated and identified. Qualitative analytical biochemistry is concernedwith identifications mainly at the molecular level but sometimes at the electroniclevel. Quantitative analysis is concerned with measuring amounts and/or concen-trations (amount per unit of volume) of constituents identified by qualitativeanalysis. The technique relies heavily on assay methods and instruments measur-ing the values of biochemical samples, representative of the whole population of

1

biomolecules, within desirable limits of accuracy and precision of the popula-tion’s true mean values. Quantitative estimations must therefore be supported byappropriate statistical analysis to establish whether or not numerical values aremeaningful in terms of the objectives being set, whether data may have beenobtained by chance or whether there is a valid pattern to the results. Accuracy(nearness to the true value) and precision (the variation in quantitative measure-ment around the norm) are both statistical terms.

Analysis, combined with microscopy (Section 1.7) and centrifugation (Chapter5), not only has identified the cell as the basic unit of life but has led to an appreci-ation of biochemical compartmentalisation within subcellular systems of thecytoplasm and cell membranes. Biochemical compartmentalisation is leastwithin prokaryotes, simple organisms lacking a true nucleus and possessing asingle periplasmic membrane, and greatest in eukaryotes, which normally havenot only a nucleus but also a nuclear membrane cytoplasmic extension, the endo-plasmic reticulum, together with associated Golgi bodies and other membrane-bound organelles, ribosomes, mitochondria, peroxisomes, and plastids (in plants),all held together by an intricate web of microtubules and filaments.

Extensive study of membrane structure and function substantiated the classicalSinger–Nicholson fluid mosaic model for membranes in which a hydrophobiclipid bilayer is loosely bound by mainly weak van der Waals’ forces to both surface(extrinsic) and deep-seated (intrinsic) proteins. ‘Proteins floating in a sea of lipid’ isa well established, but still extremely apt, image for cell membranes. Intensivestudy has been made of the structure–function correlation between membraneproteins and transport, which has revealed the importance of a-helices (Section6.2), which completely span the membrane, creating hydrophilic channels andthus allowing passage of a huge array of polar molecules (Section 8.4). Examplesinclude lactose permease in Escherichia coli, glycophorin in erythrocyte mem-branes, bacteriorhodopsin in Halobacterium halobium, cytochrome oxidase inmitochondria, the H1-ATPase of chloroplasts and mitochondria, the Na1, K1

pump of the nerve cells and Ca21 channels in secretory epithelial tissue. Entry intoa channel is, of course, selective, being controlled either directly by stereochemi-cal binding at a receptor site (Section 8.3) or indirectly by gating mechanisms thatopen or close the channel by means of conformational change in the protein,which is induced occasionally by hormonal intervention (Section 8.4.1).

The self-assembly of non-polar lipids (glycolipids, sulpholipids and sterols) intolipid bilayers and spherical micelles was a key event in biochemical evolution, notonly because it allowed membranes to organise biochemical processes into dis-crete subsystems but also because it facilitated the establishment of ion gradients.Such gradients were formed initially between the primaeval cell and its environ-ment and later, a◊er the establishment of the multicellular state, intercellularly.These ion gradients represent primary energy stores that can be exploited for per-forming biological work (transport, motion, biosynthesis, etc.). The uniqueness ofliving organisms lies in their ability to trap environmental energy and transformit through the process of anabolism into chemical bonds. These bonds typicallycomprise C 22 C, C 22 N, C 22 O, N 22 H, O 22 H, O 22 P. The energy conserved in these

2 General principles of biochemical investigations

bonds is capable of being released to perform work. Compared to eukaryotes, theprokaryote kingdom displays a much greater capacity to exploit di≠erent environ-mental niches and to use alternative energy sources. Primitive organisms wereprobably subject to natural selection based on biochemical pathways for energytransduction and biosynthesis of macromolecules.

Prokaryotes established standardised mechanisms for energy generation viasubstrate level, oxidative and photosynthetic phosphorylation. These permittednatural selection within eukaryotes, a◊er the arrival of endosymbiosis with pro-karyotes, to operate physiologically on the basis of di≠erentiation of cells intoorgans for performing specialised physiological functions. Each di≠erent cell typein a multicellular organism must reflect accompanying biochemical and physio-logical di≠erences operating within these cells and invoke mutual cooperation ofcells in physiological processes. As a consequence, a large part of developmentalbiochemistry is concerned with elucidating, at the molecular level, mechanismsof selective gene expression that lead to di≠erentiation.

In genetic terms there are significant di≠erences between species that preventcross-fertilisation and gene mixing through DNA hybridisation (phylogeneticdi≠erences). Despite these di≠erences there remain su∞cient common chemicalconstituents and processes to allow for a mode of biochemical deduction based onextrapolation of results obtained in one species to be applied to another.Frequently, extrapolation is made to humans from, for example, microorganisms,animal tissue cultures or laboratory animals for monitoring the biochemical,physiological, pharmacological or toxicological responses to foreign exogenouscompounds (xenobiotics) as a prelude to their commercial or medical use. Thisapproach must be treated with caution however, since biological variationbetween cell types or species is possible and there may be gross physiologicaldi≠erences, particularly between unicellular and multicellular species. Any formof either chemical analysis or sample preparation for microscopy inevitablycreates artefacts by destroying biological integrity, which means that in vivoextrapolation might be unjustified. Despite this criticism, however, the ‘outside-in’ analytical approach with ‘bottom-up’ extrapolation to the whole, has facili-tated great insight into the unity and diversity of life forms and providedtheoretical underpinning for other biological disciplines such as ecology, physiol-ogy, pharmacology, microbiology, botany, zoology and the environmental sci-ences which comprise biological systems.

Bioanalysis of the central information-processing molecules, DNA and RNA, inrecent years, has led to an unequivocal and clearer appreciation of the connectionbetween definable base sequences in the hereditary material, DNA, and howspecific proteins are formed, thereby allowing for a more accurate explanation ofhow cells di≠erentiate during development. Complementary advances in the arti-ficial synthesis of DNA, exploiting the polymerase chain reaction (PCR) andmethods for both multiplying and transferring DNA to unnatural recipients ingenetic engineering (Section 3.7), have facilitated an ‘inside-out’ approach tosolving problems in biochemistry that were hitherto much more di∞cult topursue using the more traditional biochemical approaches.

3 1.1 The nature of biochemistry

1.2 BIOENERGETICS

1.2.1 Laws of thermodynamics

Bioenergetics is the study of energy and biomatter transformation in biosystemsand is very closely related biochemically to metabolism, the sum of all chemicalreactions in an organism. Bioenergetics is founded on thermodynamic laws, so-called because they were concerned historically with heat (enthalpy) exchangesin chemical and physical processes in closed systems. Closed systems allowenergy, but not matter, to pass the system’s boundary. The principal interrelatedvariables involved in closed systems are pressure, volume, temperature and con-centration changes that establish an equilibrium condition, at which point thereis no exchange of energy with their surroundings. An example of a closed systemand its surroundings is a chemical solution in a sealed phial transferred from 25 °Cto a beaker of water at a higher temperature. A closed system removed from itsequilibrium state is therefore capable of either taking up energy from or giving upenergy to its surroundings.

Thermodynamics is concerned solely with measuring the di≠erence in energystatus across a boundary under non-equilibrium conditions and hence not witheither the method of achieving this state or the rate at which the process occurs.Cells are, by contrast, open systems that exchange both energy and matter withtheir surroundings when temperature, volume and pressure remain constant.

Bioenergetics is concerned with the interconversion of kinetic, thermal,mechanical, electrical, electromagnetic (light) and osmotic energy in biosystems.Bioenergetics also incorporates the energy changes associated with biomoleculeschanging from one chemical form to another in metabolism. Degradative reac-tions, involving release of energy on breaking of bonds, constitutes catabolism,whereas energy conservation in bond formation constitutes anabolism. Anabolicreactions typically involve adding hydrogen atoms (reduction) and removingwater molecules (dehydration) both to stabilise and to concentrate molecules,thus providing more energy on subsequent oxidation. Enzymic reduction in vivois facilitated by reduced forms of coenzymes, principally NADPH. Oxidation notonly invokes coenzymes such as NAD or the flavin prosthetic groups FMN or FADin dehydrogenases but also the agency of proteins that combine either withoxygen, for example cytochrome oxidase , or with alternative electron acceptorsas, for example, in the sulphate-reducing bacteria that grow anaerobically andproduce large amounts of hydrogen sulphide (H2S) from sulphate. Terminal elec-tron acceptors that are alternatives to oxygen are found exclusively in prokar-yotes.

The first and second laws of thermodynamics underpin bioenergetics. The firstlaw states that ‘Energy can be neither created nor destroyed but can be convertedfrom one form to another’, implying that the energy of the system plus its sur-roundings must remain constant. As a consequence, chemical reactions, in whichbonds are broken and reformed, move to a position of equilibrium either by takingheat from the surroundings (as in endothermic reactions) or release heat to the


surroundings (exothermic reactions). Endothermic reactions are recognised by atemperature decrease in the surroundings but show an increase in enthalpy (H)within the system, i.e. the change in enthalpy (DH) is positive. Exothermic reac-tions represent the converse, in which DH is negative when the external solutionheats up. Enthalpy changes are expressed normally in J mol21. An enthalpychange DH always accompanies the establishment of an equilbrium involvingeither an exothermic or an endothermic reaction. At equilibrium, conditions areisothermal. Enthalpy changes are related to the equilibrium constant Keq for areaction, as given by the van’t Ho≠ equation:

(1.1)

which states that the rate of change of Keq with temperature is logarithmic.Equation 1.1 may be integrated to give

ln Keq 52 1 c (1.2)

where c is an integration constant, R is the molar or gas constant (8.314 J mol21

K21) and T the absolute temperature (K). Consequently a plot of ln Keq against 1/Tfor a reversible reaction would produce a straight line graph with a slope of2DH/RT allowing DH to be calculated. Some biochemists prefer to work in log10

rather than natural logarithms. In such cases it must be remembered thatln x 5 2.303 log10 x.

The second law of thermodynamics relates to the direction of energy flow bystating that ‘All chemical and physical processes are pulled towards equilibriumby the forces of entropy (S ) which represent a condition of chaotic, random molec-ular movement’. Entropy is consequently a function of temperature and in energyterms is represented as TDS expressed in J mol21 K21. Entropy is greatest, forexample, when a chemical reaction in a closed system is at equilibrium. Entropy isa part of enthalpy that cannot be used directly to perform useful work.

Biosystems are, however, orderly, organised states in which energy flow isdirected towards the conversion of precursor molecules of greater entropy intobiopolymers with decreased entropy. Furthermore, biosystems exist characteris-tically as non-equilibrium steady states. The living state does not contradict thesecond law of thermodynamics, since life is made possible by the exploitation ofincreases in entropy in the surroundings, i.e. ultimately the universe. As a con-sequence, entropy is, paradoxically, the driving force operating on livingsystems.

Josiah Gibbs (1839–1903) was responsible for introducing the notion of a mea-surable Gibbs free energy component of enthalpy, designated G, which couldperform useful work and could also act as a reliable indicator of whether a reac-tion would proceed to equilibrium spontaneously or not from a given set of condi-tions, i.e. without the input of energy from the surroundings. Gibbs combinedenthalpy, entropy and free energy in the equation

DH 5 DG 1 TDS (1.3)

DHRT

d ln Keq

dT5

DHRT 2

5 1.2 Bioenergetics

which uses the first and second laws of thermodynamics but operates only underconditions of constant temperature, pressure and volume. This equation isfrequently rearranged as

DG 5 DH 2 TDS (1.4)

Provided the free energy status of the products of a reaction is less than for thereactants, i.e. DG has a negative sign, the reaction is termed exergonic and willproceed spontaneously towards equilibrium, irrespective of whether the reactionis exothermic or endothermic. The converse process in which the free energystatus of the products of a reaction are greater than the reactants, where DG has apositive value, is termed endergonic and will not proceed to equilibrium withoutthe input of energy from the surroundings. A negative value for DG does not givean indication of the rate of the reaction, however, which is governed by kineticfactors relating mainly to the energy of activation being lowered significantly inenzyme-catalysed reactions (Section 7.3.6). Metabolism operates through the cou-pling of exergonic and endergonic reactions to achieve chemical synthesis at theexpense of an increase in the entropy of the universe. All biochemical reactionsare spontaneous therefore when viewed from this perspective.

Spontaneous reactions accompanying positive enthalpy changes are demon-strated, for example, when crystalline salts such as potassium chloride (KCl) aredissolved in water. The solution becomes cold by taking heat out of the solvent tobreak hydrogen bonds in the crystal lattice. In this case DS resulting from theincreased mobility of the ions in water is greater than DH, making DG negativeoverall. A spontaneous reaction associated with a negative DH is observed,however, when anhydrous calcium sulphate dissolves: the temperature of thesolution increases. Spontaneous reactions induced by a positive DS of the sur-roundings, in contrast, are seen in protein folding and in the organisation of polarlipids into micelles. The principal driving force in both of these cases is the hydro-phobic action of non-polar groups (aromatic rings and alkyl groups of aminoacids, and alkyl residues in fatty acids) forcing water molecules to organise aslayers or shells on the outside surface. This happens because water moleculesinteract more strongly with each other, forming hydrogen bonds, with the non-polar groups. The initial interaction is not spontaneous in each case, however,because DH is slightly positive and entropy is decreased as the water molecules areforced to organise externally. Protein folding follows when hydrophobic groupsexposed at the surface become more stabilised within the interior of the protein bynon-covalent van der Waals’ forces. Any hydrophilic amino groups drawn into thehydrophobic interior reduce their polarity by hydrogen bonding to form a-helicesand b-pleated sheets, thus stabilising the protein structure. This infolding allowsexternal water molecules to become more mobile and consequently the entropy ofthe solvent increases, making the overall process spontaneous. A similar increasein S of the surrounding water is achieved when phospholipids aggregate sponta-neously first into monolayers and then into spherical micelles.

Gibbs free energy values in J mol21 can be calculated in relation to four princi-pal functions: chemical potential, electrical potential, hydrostatic potential and


gravitational potential. Chemical potential and electrical potential are the mostimportant in biochemical terms.

1.2.2 Standard energy status

The chemical potential of a solution is the product of its molar concentration (M)multiplied by its activity coe∞cient a. Activity coe∞cients for physiological solu-tions, which are usually in the mM or mM range, are su∞ciently close to 1 that con-centrations will substitute e≠ectively for chemical activity. The standard energystatus (or standard electrochemical potential) of m0, a unique property of the mole-cule resulting from the interaction of its chemical bonds, relates to the free energyof a solution containing 1 mol dm23 at atmospheric pressure and at a given tem-perature, normally 25 °C, representing the standard state. Increases or decreases inthe molar concentration changes the energy status to m by a logarithmic factoraccording to the equation

m 5 m0 1 RT ln [A] (1.5)

where [A] is the concentration of solute A in mol dm23. When [A] equals 1 M inequation 1.5, m 5 m0, since ln 1 5 0.

The di≠erence in m, Dm, if the concentration were to be changed from [A] to alower concentration [A1], would be

(m0 1 RT ln [A1]) 2 (m 0 1 RT ln [A])new status old status

Dm 5 RT ln (1.6)

assuming the molecule remained chemically unaltered. DG would have a nega-tive sign in this case and free energy would be available potentially to performwork. Biological membranes play an important role in energy conservation bymaintaining such solute gradients when selectively preventing solute transport.When membranes are permeable to the substance in question it would movespontaneously from the higher to the lower concentration, moving down the con-centration gradient to dissipate the potential energy.

Chemical potential can be used to work out the thermodynamic feasibility of achemical reaction, i.e. whether or not it will proceed spontaneously to equilib-rium, when the concentration of the reactants and products are known, providedthe equilibrium constant for the reaction is also known. This relationship may bederived from considering the hypothetical reaction A ⇀↽ B, in which case

m[B] 5 m0[B] 1 RT ln [B]and m[A] 5 m0[A] 1 RT ln [A]

Dm 5 (m0[B] 1 RT ln [B]) 2 (m0[A] 1 RT ln [A])

Dm 5 (m0[B] 2 m0[A]) 1 RT ln (1.7)

where [B]/[A] represents the mass action ratio.

1[B][A]2

1[A1][A] 2

7 1.2 Bioenergetics

But if both B and A are at standard 1 M concentrations then

RT ln 5 0

from which Dm 5 m0[B] 2 m0[A] 5 DG0, which defines the standard free energychange for a reaction. The standard free energy change, DG0, constitutes thedi≠erence in free energy between the equilibrium state where G 5 0 and thatwhen reactants and products are maintained artificially under standard states atm0, i.e. at a uniform concentration of 1 mol dm23 at atmospheric pressure (101

kPa), hydrogen ion concentration of 1 mol dm23 (pH 0, Section 1.2.5) and at a giventemperature, usually 25 °C (298 K).

For any other molar concentrations of B and A

Dm 5 DG 5 DG0 1 RT ln (1.8)

Under equilibrium conditions

DG 5 0 and K eq 5

therefore DG0 52 RT ln 52 RT ln K eq (1.9)

1.2.3 Standard free energy change

DG0 has very little relevance in biological systems because the value of cellular pHis normally around neutrality (6.0 to 7.5). As a consequence DG09 is used as thestandard free energy change at pH 7.0, when the equation that applies is

DG09 52RT ln Keq (1.10)

The units for the equilibrium constant in equation 1.10 will change accordingto whether a single product is formed from a single reactant (no units) to M(where two products are produced from a single reactant) to M21 (where twoproducts give rise to a single product). This does not a≠ect the units of measure-ment for DG09, which remain J mol21. For this reason units for Keq are frequentlyomitted.

For non-standard, non-equilibrium mass action ratios, the relevant free energyequation is therefore

DG9 5DG09 1 2.303RT log (1.11)

The phosphorylation of fructose 6-phosphate (F6P) by ATP to fructose 1,6-bisphos-phate (FBP) by phosphofructokinase-1 (PFK-1) may be used to illustrate the sig-nificance of this equation.

1[products in mol dm23][reactants in mol dm23]2

1[B][A]2

[B][A]

1[B][A]2

1[B][A]2


The equilibrium constant Keq 5

5 3.08 3 102 M21

If FBP, ATP and F6P were to be maintained at pH 7.0 away from equilibrium at stan-dard 1 M concentrations, the Gibbs free energy at t (K) is given by the equation

DG9 5DG09 1 2.303RT log

but, since

5 1

because reactants and products are at unit concentration, DG9 5DG09.At equilibrium DG9 50 and DG09 522.303RT log Keq (equation 1.10). Hence in

the example for PFK-1 above, at 25 °C, pH 7, DG0 can be calculated as follows:

DG 09 522.303 3 81314 J mol21 K21 3 298 K 3 log (3.08 3 102)

5 214 199.3 J mol21

Putting the mass action ratio at exactly 1 would make approximately 14.2 kJ offree energy available to perform work if the reaction were allowed to move to theequilibrium position. Whenever Keq . 1, therefore, an equimolar solution of prod-ucts and reactants would proceed spontaneously to equilibrium. The converseapplies for reactions where Keq ,1, in which case energy would need to be pro-vided from the surroundings to attain the equilibrium condition.

For coupled reactions involving a common intermediate

i.e. Keq1 Keq2

A −⇀↽− B −⇀↽− C,

where B is the common intermediate, the overall Keq for the conversion of A to C isthe product Keq1 3 Keq2. In coupled reactions that include a reaction proceeding inthe backward direction, a reciprocal value for Keq is used to calculate the overall Keq.Examples of calculations of DG09 involving coupled reactions are shown below.

The pentose phosphate pathway (PPP) is the metabolic sequence for the synthesisof pentose sugars. Ribose 5-phosphate and xylulose 5-phosphate are typicalpentose sugars found in the PPP. What is the standard free energy for the forma-tion of ribose 5-phosphate from xylulose 5-phosphate at 37 °C, given

ribulose 5-phosphate Keq 5 2.30

ribose 5-phosphate−−−−−−−⇀↽−−−−−−−

ribulose 5-phosphate Keq 5 0.67

xylulose 5-phosphate−−−−−−−⇀↽−−−−−−−

Example 1

1 [FBP ][ATP][F6P]2

1 [FBP ][ATP][F6P]2

1 [FBP ][ATP][F6P]2

9 1.2 Bioenergetics

The overall Keq 5 [(1/0.67) 3 2.30] 5 3.433

Hence DG09 522.303 3 8.314 J mol21 K21 3 310 K 3 log 3.433

523179.55 J mol21

Hence the di≠erence in free energy between products and reactants poised at 1 Mconcentrations would be 3179.55 J less than when they are at equilibrium condi-tions. This represents a form of potential energy that is theoretically available toperform work such as coupling to endergonic reactions.

An important stage in the breakdown of sugars via the Embden–Meyerhof (gly-colytic) pathway occurs by the coupling of glyceraldehyde 3-phosphate dehydro-genase with phosphoglycerate kinase to form ATP in a substrate levelphosphorylation. If the coupled reactions occurred together in solution at pH 7and 38 °C, what would be the standard free energy change for the formation ofATP from glyceraldehyde 3-phosphate and Pi?.

1,3-bisphosphoglycerate 1 ADP Keq 5 1.746 3 10

3

3-phosphoglycerate 1 ATP−−−−−−−−−−−⇀↽−−−−−−−−−−−

The overall Keq 5 0.0875 3 1746 5 152.78

DG09 522.303 3 8.314 J mol21 K21 3 311 K 3 log 152.78

5213.005 kJ mol21

The DG09 generated from the phosphoglycerate kinase reaction alone is equal to218.5 kJ mol21 under the conditions specified. This is considerably less than theDG09 for ATP synthesis from ADP and Pi under the same conditions, whichamounts to approximately 230.5 kJ mol21. Substrate-level phosphorylation isachieved in vivo in the phosphoglycerate kinase reaction by exploiting massaction ratios.

In the case of PFK-1 cited earlier, if the concentrations of FBP, ATP and F6P were,respectively, 80 mM, 8 mM and 1.5 mM, what would be the actual free energychange?

DG9 5DG091 2.303 RT log

5214199.3 J mol21 1 2.303 3 8.314 J mol21 K21 3 298 K

3 log 1 80 3 1026

1.5 3 1023 3 8 3 10

232

1 [FBP][F6P][ATP]2

Answer

Example 3

Answer

K eq 50.8 31021

glyceraldehyde 3-phosphate 1 inorganic phosphate

NAD NADH1 H 1

1, -bisphosphoglycerate

75

3

Example 2

Answer


5214199.3 J mol21 1 4701.1J mol21

5 29498.2 J mol21

Physiological concentrations of FBP are maintained in a steady state, wellremoved from near-equilibrium conditions, making this reaction metabolicallyirreversible. The constraint on PFK in vivo is achieved by allosteric regulation(Section 7.3.4).

DG°9 for ATP synthesis under standard conditions at 25 °C 5 30 540 J mol21. Theconcentrations of ATP, ADP and Pi in cells is normally such that considerablymore energy than 30 540 J mol21 is necessary to synthesise 1 mole of ATP fromADP and Pi. If the steady state concentrations of ATP, ADP and Pi in isolatedchloroplasts under full illumination at pH 7.0 and 25 °C were, respectively, 8 mM,0.5 mM and 5 mM, what would be DGp (the phosphorylation potential), the freerequirement for the synthesis of 1 mole of ATP under these conditions?

DGp 5DG091 2.303 RT log

5230 540 J mol21 1 2.303 3 8.314 J mol21 K21 3 298 K

3 log

5230 540 J mol21 1 19 999 J mol21

5 50.54 kJ mol21

Frequently, in vivo DGp is of the order of 50 kJ mol21.

Many reactions that represent a thermodynamic barrier as judged by positiveDG09 values can be made spontaneous by manipulating the mass action ratio. Anillustration occurs in the synthesis of glyceraldehyde 3-phosphate (G3P) from 3-phosphoglycerate (3PGA) via 1,3-bisphosphoglycerate (1,3-BPGA) in the photo-synthetic carbon reduction cycle (Calvin–Benson cycle). The first stage involvesphosphoglycerate kinase:

3PGA 1 ATP −−−⇀↽−−− 1,3-BPGA 1 ADPDG09518.5 kJ mol21

This reaction is endergonic under standard conditions and consequently ATP isnot being used to drive the reaction. However,

since DG9 5DG0912.303 RT log

the thermodynamic barrier is reduced in vivo and the reaction made exergonic fora number of reasons. First, 3PGA levels are high following ribulose bisphosphate

1[ADP][1,3 BPGA][3PGA][ATP] 2

1 8 3 1023

0.5 3 1023 3 5 3 10

232

1 [ATP][ADP][Pi]2

Answer

Example 4

11 1.2 Bioenergetics

carboxylase (Rubisco) activity; secondly, ATP levels are high as a result of photo-phosphorylation; thirdly, 1,3-BPGA is low through coupling of phosphoglyceratekinase to glyceraldehyde-3-phosphate dehydrogenase. This enzyme representsthe second stage of the process:

1,3-BPGA 1 NADPH −−−⇀↽−−− G3P 1 NADPDG09526.3 kJ mol21

Despite the low concentrations of 1,3-BPGA, the reaction is made more exergonicunder physiological conditions by the generation of large amounts of NADPHfrom NADP as a result of the light reactions of photosynthesis. Glyceraldehyde-3-phosphate dehydrogenase also has a very high a∞nity for both 1,3-BPA andNADPH (Km 1,3-BPA 5 1 mM, Km NADPH 5 4 mM). (Details on the importance ofKm for enzyme-catalysed reactions are given in Section 7.3.2.)

1.2.4 Reduction–oxidation (redox)

Many biological molecules are known to exist in either a reduced (red) or oxidised(ox) form as a result of either gaining or losing electrons to molecules with similarproperties. Such substances are known as redox couples or half-cells. (Furtherdetails on theoretical and practical aspects of redox substances are given inSection 15.3). When two half-cells are mixed, electron transfer takes place in aredox chemical reaction such that the oxidising half-cell becomes reduced and thereducing half-cell oxidised, thus establishing an equilibrium condition when elec-tron transfer ceases. Free energy changes accompanying such redox reactions aretypified hypothetically as

in which A/AH and B/BH represent two substances in their oxidised and reducedforms, respectively.

The free energy expressed in electrical terms for a 1 M standard solution con-taining a monovalent cation (C1) is the product nFE, where nF coulombs is thecharge on one mole of the substance. In the case of a monovalent ion n 5 1. TheFaraday constant, F, is the product of the charge on a single electron(1.602 3 10

219 C) and Avogadro’s number (6.022 3 1023) (the number of mole-

cules in one mole; see Section 1.4.4). Thus F 5 9.648 3 104 C. The term E relates to

the electric potential to which the charge is subjected. An individual half-cellcan be set up such that its oxidised and reduced forms are both at 1 M concentra-tion (50% oxidised), i.e. under the same standard state as for chemical potentialdescribed earlier. When two half-cells are arranged to form a cell, such that each

reducing half-cell oxidising half-cell

BH A

B AH


half-cell is in its standard state, a potential di≠erence is set up between them.This DE is dependent upon the di≠erent a∞nities that each half-cell has for elec-trons, as measured by their individual electrode potentials. Electrode potentialsare measured for any given half-cell against the hydrogen half-cell in a standardstate, in which a 1 M solution of HCl (at pH 0) is equilibrated with H2 gas (Section15.2). The standard electrode potential for a given half-cell is known as the redoxpotential and is designated as E0 (pH 0) or E09 (pH 7.0). The E0 for the hydrogenhalf-cell is 0.0 V at pH 0; at pH 7.0 it is 20.3200V (Section 15.3). Redox potentialsare consequently ordered as either being more negative or more positive thanthe hydrogen half-cell. In moving towards an equilibrium position, any sub-stance with a more negative standard redox potential than the hydrogen half-cell will spontanously reduce it, causing H2 gas to be liberated at the hydrogenhalf-cell. Conversely any substance having a standard redox potential more posi-tive than the hydrogen half-cell will spontaneously oxidise it in moving towardsequilibrium, but H2 gas will not be produced. If equilibrium were obtained, asingle redox potential would be achieved and electron flow would cease. Table15.1 (p. 743) lists some standard redox potentials (E09 values) of biological impor-tance.

For two half-cells maintained under standard states, the potential energy avail-able represents the standard free energy change DG09, and is equal to nF DE09,where DE09 is the di≠erence in the standard redox potentials between the two half-cells. Subtracting the equilibrium state where there is no free energy availablefrom the standard state produces a negative term in the expression

DG09 52 nF DE09 (1.12)

If DE09 is negative, i.e. E09 for the oxidising half-cell is less than E09 for the reducinghalf-cell DG09 would be positive and equimolar concentrations of reactants andproducts would require an input of energy from the surroundings.

In general terms, the redox potential of a given half-cell will change accordingto the proportion of its oxidised and reduced forms as given by the Nernst equa-tion:

i.e. E9 5 E09 1 2.303 log (1.13)

where E09 is the half-oxidised (midpoint redox potential) and nF is the charge permole. Physiological values for redox potential relate therefore to the logarithmicratio of [ox]/[red] which means that they are independent of concentration as longas dilute solutions are used. Equation 1.13 also shows that, at 50% oxidation for agiven half-cell, E9 5 E09. The Nernst equation may be used to facilitate exergonicreactions, for example when the redox potential is lowered in chloroplasts whenlarge amounts of light-generated NADPH are used for the glyceraldehyde-3-phos-phate dehydrogenase reaction described earlier.

In cases where products and reactants are away from their half-oxidised condi-tion, the mass action ratio has to be taken into account to calculate the new valuefor DE9. Hence for the hypothetical reaction

1 [ox][red]2

RTnF


E9A/AH 5 E09A/AH 1 2.303 log

E9B/BH 5 E09B/BH 1 2.303 log

DE9(ox–red) 5 E9A/AH 2 E9B/BH 2 (E09A/AH 2 E09B/BH) 1 2.303 log 2 log

DE9(ox–red) 5 DE09 2 2.303 log (1.14)

allowing the value of DG9 to be calculated from DG9 52 nF DE9.

Some examples of the application of redox potential for calculating DG 09 areshown below:

What is the equilibrium constant for the oxidation of acetaldehyde (ethanal) toacetate (ethanoate) by ferredoxin at pH 7.0 and 30 °C?

acetaldehyde −−−⇀↽−−− acetate 1 2H1 1 2e (E09520.60 V)ferredoxin (Fe21) −−−⇀↽−−− ferredoxin (Fe31) 1 e (E09 520.432 V)

The redox reaction is represented by

DG09 52nFDE09 522.303 RT log Keq

log Keq 5

5

5

5

log Keq 5 5.6097

Keq 5 4.07 3 105

32544.96

5801.58

(2 3 96 860 J V21 mol21 3 0.168 V)(2.303 3 8.314 J mol21 K21 3 303 K)

{2 3 96 860 J V21 mol21 3 [( 2 0.432 V) 2 ( 2 0.60 V)]}(2.303 3 8.314 J mol21 K21 3 303 K)

nF DE 09

2.303 RT

reducing half-cell oxidising half-cell

acetaldehyde 2 ferredoxin (Fe31)

2 ferredoxin (Fe21)acetate

Answer

Example 5

1[AH][B][A][BH]2

RTnF

1 [B][BH]21 [A]

[AH]2RTnF

1 [B][BH]2

RTnF

1 [A][AH]2

RTnF

BH A

B AH


In mitochondria, energy resulting from electron transport is made available forproton pumping by the oxidation of protons arising from dehydrogenase activityin the tricarboxylic acid (or Krebs) cycle to water. How many moles of ATP couldtheoretically be produced under standard conditions of pH 7.0, and 25 °C from theproton gradient established, given DG09 for ATP synthesis from ADP 5 30 540 Jmol21?

NADH → NAD1 1 H 1 2e (E09 5 20.320 V)

H2O → O2 1 2H1 1 2e (E09 5 10.816 V)

DG09 52nFDE09 5 22 3 96 860 J V21 mol21 3 [(0.816 V) 2 (20.320 V)]522 3 96 860 J V21 mol21 3 (1.136 V)52220.066 kJ mol21

Number of moles of ATP 5

5 7.21 moles

This is considerably greater than the experimentally observed value of 2.5, high-lighting ine∞ciencies in the coupling process.

In photosynthetic electron transport, light quanta are transformed into ATP bythe light reactions of photosynthesis involving the excitation of chlorophyll mol-ecules. The e≠ective wavelength range for higher plant photosynthesis is appro-ximately 400 nm to 710 nm for O2 evolution resulting from the photolysis of H2O,i.e. oxygenic photosynthesis. Show by means of calculations that one photon of710 nm light could provide enough energy to reduce NADP following oxidationof water, assuming E09 values for the H2O/O2 and NADPH/NADP half-cells to be,respectively, 10.816 V and 20.342 V (DE09 5 11.14 V).

The energy per quantum is given by the equation

E 5 (1.15)

where c is the velocity of light (2.998 3 108 ms21), h is the Planck constant 5

6.626 3 10234 J s, and l is the wavelength. The energy per quantum (E ) is

expressed in J photon21.Hence, when l 5 710 nm,

E 5

5 2.798 3 10219 J photon21

Since this unit is very small, it is more convenient to express energy asJ einstein21, where 1 einstein is equivalent to one mole of photons, which isAvogadro’s number of photons.

6.626 3 10234 J s 3 2.998 3 10

8 ms21

710 3 1029 m

hcl

Answer

Example 7

2 220.066 kJ mol21

2 30.540 kJ mol21

Answer

1

2

Example 6


The energy per einstein of 710 nm light 5 2.789 3 10219 3 6.023 3 10

23 photons5 mol21

5 168.52 3 103 J mol21

5 168.52 kJ mol21

This amount of energy may be converted into V by dividing it by the Faraday con-stant, 9.6485 3 10

4 J V21 mol21.

Hence 168.52 kJ mol21 5

5 1.75 V

Note that this value of 1.75 V can also be obtained by considering the displace-ment of a single electron by a single photon when

5 1.75 V

where 1.602 3 10219 J V21 is the charge on a single electron 5 Faraday constant

(F )/Avogadro’s number.This means, in principle, that one photon of 710 nm wavelength light could

provide enough energy to reduce NADP a◊er oxidation of water. In practice,however, two photons are involved cooperatively in transferring a single electronfrom water to NADP. The reasons for this include the facts that redox carriersmore positive than H2O/O2 (in particular an oxidised form of chlorophyll at themanganese reaction centre of photosystem II) and more negative thanNADPH/NADP are involved (notably the non-haem protein ferredoxin), redoxcarriers are not in their standard states, energy transfer losses are incurred, andATP, ADP and Pi mass action ratios require a DGp value of around 50 kJ mol21

under physiological conditions. In vitro chloroplast preparations have shownthat 710 nm represents a cut-o≠ point for oxygenic photosynthesis, although710 nm will continue to reduce ferredoxin in a cyclic flow involving a chloro-phyll reaction centre (Section 15.3).

1.2.5 Electrochemical potential

Chemical and electrical potential for an electrolyte (A) may be combined as theelectrochemical potential by the equation

m 5 m0 1 RT ln [A] 1 nFDE (1.16)

As a consequence, charged ions in solution are able to exert an electromotive forceon other ions by virtue of their electrochemical potential. This equation isexploited in cells by the di≠erential permeability of membranes that establishequilibrium conditions of electrochemical potential by creating di≠erent electri-cal potentials in adjacent compartments whilst holding solutions at di≠erent

2.789 3 10219 J photon21

1.602 3 10219 J V21

168.52 3 103 J

9.6485 3 104 J V21


concentrations. This property is vital in, for example, maintaining osmoticbalance between cells to facilitate water absorption. The di≠erence in electricalpotential across a membrane associated with such asymmetric solute concentra-tions represents a potential energy source and is defined in volts (V) as the Nernstpotential.

The Nernst potential may be derived by considering the electrochemical poten-tials of a cation [C1] distributed in two compartments inside (i) and outside (o),separated by a membrane freely permeable to [C1] but not to any counterion [A2].If the concentration of [C1] at equilibrium is represented by [C1]i and [C1]o respec-tively, then the di≠erence in molar energies in the two compartments is given by

Dm 5 mi 2 m0

5 (RT ln [C1]i 1 nFEi) 2 (RT ln [C1]o 1 nFEo)

5 (RT ln [C1]i 2 RT ln [C1]o) 1 (nFEi 2 nFEo)

5 RT ln 1 nF (Ei 2 Eo)

5 RT ln 1 nF (EM) (1.17)

where EM 5 (Ei 2 Eo) is the membrane potential.But at equilibrium Dm 5 0 when EM 5 EN, the Nernst potential. Hence the equil-

brium Nernst potential equation may be written as

EN 5 ln (1.18)

This is the electromotive force in volts associated with maintaining a particularion gradient across a membrane. RT/nF is a constant for a univalent ion equal to0.0295 V at 25 °C.

A term may be found for ln ([C1]i/[C1]o) by inverting the previous equation to

2 EN 5 ln

when ln 52EN

Equation 1.17 may therefore be rewritten as

Dm 52ENnF 1 nFEM

5 nF (EM 2 EN) (1.19)

This equation represents the energy stored in a membrane ion gradient main-tained under non-equilibrium conditions. Considerable amounts of energy maybe used in vivo to set up such conditions.

The Nernst equation has considerable physiological application where potentialdi≠erences are established between cells that are subject to change under exitable

nFRT1 [C1]i

[C1]o2

1 [C1]i

[C1]o2RT

nF

1[C1]o

[C1]i2RT

nF

1 [C1]i

[C1]o2

1 [C1]i

[C1]o2


conditions. For example, the resting potential of the squid giant axon, a favouritemodel system, may be calculated from the Nernst equation at 16 °C to be approxi-mately 275 mV when the concentrations of [K1] outside is 20 mM and inside is 400

mM. Experimental values are typically nearer 260 mV at this temperature,however, a◊er the concentrations and di≠erential permeability of other ions suchas Na1, Cl2 and endogenous amino acid anions are taken into account. Frog muscleis another example of tissue that produces a resting potential across the membraneof approximately 290 mV when the outside concentrations of Na1, K1, and Cl2 arerespectively 120 mM, 2.5 mM and 120 mM and when the internal concentrations ofthe same ions are, respectively, 10 mM, 140 mM and 4 mM.

When this principle is applied to the distribution of H1 across biological mem-branes, for example the inner mitochondrial membrane, another useful term maybe derived, which Mitchell and Moyle, in 1968, called the proton motive force(p.m.f.). In this case H1 accumulates in the intermembrane space as electrons areconveyed from reduced hydrogen carriers NADH or FADH towards oxygen(Section 15.3). If the su∞xes i and o are taken to represent the matrix side (wherethe Kreb’s cycle is located) and the intermembrane space respectively, then

Dm[H1] 5 m[H1]i 2 m[H1]

o

Substituting in equation 1.17, in which n 5 1, gives

Dm[H1] 5 RT ln 1 F(Ei 2 Eo)

52 RT ln 1 F(Ei 2 Eo)

522.303 RT log 1 F(Ei 2 Eo)

But pH 52log H1 (Section 1.4.5).

Therefore, Dm[H1] 5 2.303 RT DpHo2i 1 F(Ei 2 Eo)

where DpHo2i is the di≠erence in pH across the membrane.This is usually expressed the other way around as

Dm[H1] 5 F EMi2o 1 2.303 RT DpHo2i (1.20)

Dm[H1] is converted into volts by dividing by F to give

p.m.f. 5 EMi2o 1 2.303 DpHo2i

i.e. p.m.f. 5 EMi2o 1 zDpHo2i (1.21)

where z is a constant for a given temperature.An electrochemical gradient of any ion represents a potential energy store, since

energy release would follow translocation of the ions down an electrochemical

RTF

1[H1]o

[H1]i2

1[H1]o

[H1]i2

1 [H1]i

[H1]o2


gradient through a channel in the membrane. In oxidative and photosyntheticphosphorylation, electrochemical gradients are established through H1 accumula-tion against a concentration gradient in the periplasmic space of bacteria, the inter-membrane space of mitochondria and the interthylakoid (lumen) of chloroplasts.This accumulation exploits the equation DG9 52nFDE9. The proton gradient thenallows protons to traverse the membrane through a proton channel represented bythe protein complex ATP synthase. In vivo this enzyme is also capable of ATPaseactivity that actively pumps protons in the opposite direction. The coupling ofenergy-releasing electron flow to the establishment of electrochemical gradientswas accommodated in Peter Mitchell’s chemiosmotic theory. It involved oxidationof hydrogen carriers by electron carriers, releasing protons into the intermem-brane space, with electron flow continuing until it reached the terminal electronacceptors. Evidence in favour of the chemiosmotic theory has derived mainly frommitochondrial and chloroplast preparations using O2 electrodes, pH electrodes andsophisticated spectrophotometric techniques (Section 15.1.2).

1.2.6 Group transfer molecules

In metabolism, exergonic and endergonic reactions are frequently coupledthrough the agency of a class of intermediates known as group transfermolecules. These comprise nucleoside di- and triphosphates (e.g. adenosine tri-phosphate and diphosphate), acylphosphates (e.g. acetyl phosphate, 1,3-bisphos-phoglyceric acid), enol phosphates (e.g. phosphoenolpyruvate) and acyl-CoAderivatives (e.g. acetyl-CoA and succinyl-CoA). Synthesis of such group transfermolecules is dependent on the ambient conditions of concentration and thereduction–oxidation states of participating molecules, allowing for energyrelease to be captured in chemical bonds. Table 1.1 shows the major classes ofgroup transfer molecule that act as the energy currency in metabolism. Thesemolecules are especially energy rich because they have a specific set of chemicalbonds that are readily capable of being changed as a result of hydrolysis to mole-cules of lower energy status (greater entropy) with the concomitant release ofenergy. Factors contributing to free energy changes include lowering of negativecharge repulsion, stabilisation of products through isomerisation and electronsharing by resonance. In vivo hydrolysis of group transfer molecules does not nor-mally occur, but instead energy is conserved in the biosynthesis of new chemicalbonds.

The structures of ATP and 1,3-bisphosphoglycerate are shown in Fig. 1.1,which illustrates the principles involved. The enzymic transfer of a phosphategroup from a group transfer molecule to ADP is known as substrate-level phos-phorylation (SLP). Substrate level phosphorylation is ubiquitous in livingorganisms but is especially important in anaerobes. ATP and ADP are the mostimportant group transfer molecules because they are very stable in water andare frequently covalently bonded to the active site of many enzymes, usually asthe magnesium salt, whenever phosphorylation/dephosphorylation is involved.This includes the phosphorylation of proteins by protein kinases (Section 8.4.3).


Cambridge University Press is part of the University of ...assets.cambridge.org/052179/9651/sample/0521799651WS.pdf · Second edition 1981 Third edition 1986 Third edition ﬁrst

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