chemical basis of metabolism

The Chemical Basisof Metabolism

A logical approach tobiochemical reactions

FAT

CARBOHYDRATE

PROTEIN

CO2

H2O

UREA

Robert Warner Chambers

The Chemical Basis of Metabolism

A LOGICAL Approach to biochemical reactions

The chemical basis of metabolism

A LOGICAL Approach to biochemical reactions

Version 1.9

Robert Warner ChambersProfessor of Biochemistry and Molecular Biology

Dalhousie University

Tangier Smith, publisherLake of the Woods, Nova Scotia

Canada

Tangier Smith, publisherLake of the Woods, Nova Scotia, Canada 2010 by Robert Warner ChambersAll rights reserved

This book is intended for the personal use of students and teacherswithout charge. Use of any part of this book for profit is strictly forbiddenunder Copyright Law. However, permission is granted to print the bookverbatim for individual use. The book may not be changed in any waywithout express, written permission of the author.

ISBN: 978-0-9866880-0-3

To the Reader: I need your help to purge errors and improve the book. Please send

comments and/or corrections to Professor Robert W. Chambers athttp://www.biochemistrybob.com/

The text is set in Garamond Premier Pro, 12 pt. The type was set to beeasily readable on the printed page. You may find it more readable on acomputer screen if you zoom to 200% in Adobe Reader. All the structureswere drawn with ChemDraw (CambridgeSoft, Cambridge, MA, USA),using Helvetica 10 pt. for atom labels and Times 12 pt. for captions

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The Chemical Basis of Metabolism i

TABLE OF CONTENTS

PREFACE.. ix

INTRODUCTION....1The Unity of Biochemistry. 1Electron Pushing and Metabolic Reactions .1Prerequisites.. .2

part 1. fundamentals

CHAPTER 1. Some Principles of Bonding .61.1. Breaking Covalent Bonds.61.2. The Heisenberg Uncertainty Principle. 71.3. The Shape of Orbitals.. 71.4. Molecular Orbitals. 81.5. Bonding and Antibonding Orbitals.. 91.6. Molecular Orbital Energy Diagram for Hydrogen.. . 101.7. MO Diagram for Helium ... 111.8. Acetylene. 121.8.1. Acetylene Uses sp Hybrid Orbitals to Form -Bonds... 131.8.2. Acetylene Uses p Orbitals to Form -Bonds.. 161.8.3. Some Important Conclusions... 191.9. How Reactions Occur. . 201.10. Ethylene... 201.11Geometrical Isomers . 231.12. Tautomerization. 231.13. Esters, Amides, Ketones and Aldehydes Have -Bonds.. . 241.14. Electronegativity . 241.15. sp3 Hybrid Orbitals.. . 251.16. Phosphate.. . 271.17. Rules to Remember. 311.18. Summary. 32

Table of Contentsii

CHAPTER 2. Some Stereochemistry. 3342.1.Dealing with Asymmetric Atoms in 3D. 342.2. The Fisher Convention.. 372.3. Enantiomers 392.4. D and L Configurations Are Based on D-glyceraldehyde. 412.5. The R/S Convention.. 422.6.1 The Sequence Rule 422.6.2. Assigning Priority to Atoms with Multiple Bonds.. 432.6.3. The Chirality Rule. 442.7.1. Compounds with More than One Chiral Center 452.7.2. Diasteriomers 462.7.3. Assigning the Absolute Configuration of Diasteriomers.. 472.8.1. Prochirality 482.8.2. Prochiral Hydrogens. 502.9. The Stereospecific Numbering System.. 522.10. Summary. 57

CHAPTER 3. Addition, Elimination and Substitution 603.1. Addition to Isolated Double Bonds 603.2. Addition to Conjugated Double Bonds 613.3. Addition Require a Nucleophile and a Leaving Group.. 633.4. Some Properties of Addition. 643.5. Stereochemistry of Addition to C-C Double Bonds... 643.6. Syn and Anti Addition 663.7. Face Specificity. 673.8. Isomers Provide Information About the Mechanism 693.9. Elimination is the Reverse of Addition 703.10. Three Elimination Mechanisms Are Known 713.11. C-C Double Bonds Can Be Formed by Elimination Involving

Saturated Carbon Atoms.. 723.12. Substitution Reactions Occur at Saturated Carbon Atoms 733.13. SN2 Reactions Occur by a Concerted Mechanism 733.14. SN1 Reactions Occur by a Stepwise Mechanism 743.15. The Reaction Mechanism Can Influence the Stereochemistry. 743.16. The Mechanism Can Influence the Product Distribution... 743.17. Summary.. 753.18. Problems for You to Solve. 763.19. Answers. 80

The Chemical Basis of Metabolism iii

CHAPTER 4. Carbon-Carbon Bond Reactions. . 8864.1. Illustrative Problems.. 874.2. Stabilization of Carbanions... 884.3. The Aldol Condensation.. 904.4. Aldol is Set Up for -Elimination... 914.5. The Claisen Ester Condensation . 924.6. Decarboxylation of -Keto Acids. . 944.7. The Benzoin Condensation.... 954.8. Decarboxylation of -Keto Acids . 974.9. The Role of Thiamin Pyrophosphate in Decarboxylation

of -Keto Acids.. . 994.10. Ketol Condensations. . . 1024.11. Some Important Rules... . 103

CHAPTER 5. Redox Reactions. . 11055.1. Definition of Oxidation. . . 1055.2. Balancing Equations.. . 1055.3. Generalized Oxidation... 1085.4. Oxidation of Alkanes.. 1085.5. Oxidation of Alkenes 1095.6. Oxidation of Alcohols. . 1095.7. Oxidation of Aldehydes. . 1115.8. Oxidation of -Keto Acids.. . 1135.9. Hydride Ion Acceptors . 1165.10. Oxidation by Electron Transfer Requires Lipoamide. . 1195.11. Summary.. . 123

CHAPTER 6. Energetics. 11256.1. Free Energy. . 1256.2. Activation Energy . 1266.3. An Equation for Free Energy and What It Tells Us.. . 1306.4. The Effect of Concentration of Free Energy.. 1326.5. Entropy. . 1326.6. The Importance of Protons.. . 1346.7. Hydrolysis of G-1-P and G-6-P . 1366.8. Hydrolysis of Phosphoenol Pyruvate. 1396.9. Hydrolysis of Acetyl-phosphate.. . 1396.10. Hydrolysis of ATP.. 1406.11. High Energy Bonds.. . 1426.12. Summary of Structure/Free Energy Relationships . 143

Table of Contentsiv

6.13. The Importance of Entropy in Reactions That Do Not Involve Protons.. 143

6.14. Phosphates React by Subsitution , 1476.15. Coupled Reactions. . 1496.16. Enzymes Adjust the Concentrations of Reactants

and Products in a Metabolic Pathway... 1516.17. Summary.. . 156

CHAPTER 7. Electron Transport and Oxidative Phosphorylation. . 11597.1. Some Facts. 1597.2. Oxidation of NADH or FADH2 by O2................................. 1607.3. Mitochondria. . 1607.4. Protons Drive the Synthesis of ATP . 1617.5. The Electron Transport System.. 1657.6. Complex I. . 1667.7. Coenzyme Q. 1687.8. Complex II Does Not Pump Protons.. . 1697.9. Complex III and IV Each Pump Two Protons/Electron.. . 1707.10. Facts to Remember . 1707.11. The Structure of Molecular Oxygen. . 1717.12. Reduction of Molecular Oxygen. . 1747.13. Transition Metal Ions.. 1777.14. Coordinate Complexes . 1797.15. Iron-sulfur Proteins Have a Coordination Number of 4.... 1797.16 Octahedral Fe Ion Complexes.... 1827.17. Iron-porphyrin Proteins. . 1847.18. Cytochrome Oxidase . 1857.19. Summary.... 187

PART 2. CATABOLIC PATHWAYS

CHAPTER 8. Fatty Acid Oxidation. 1928.1. Balancing the Overall Pathway . 1928.2. Condensation of Two Acetyl-CoA Moieties.. 1938.3. Reduction of the -Keto Group. . 1948.4. Elimination of Water . 1968.5. Reduction of the Double Bond . 1968.6. Hydrolysis of Acyl-CoA.. . . 197

The Chemical Basis of Metabolism v

8.7. -Oxidation of Butyric Acid.. 1978.8. Formation of Fatty Acyl-CoA from a Fatty Acid. 1988.9. -Oxidation of Long Chain Fatty Acid . 2018.10. Summary of -Oxidation Reactions . 2018.11. The Advantages of Working Backward... 2028.12. Enzymes and Other Details of -Oxidation . 2038.13. Stereochemistry.. . 2048.14. The Energy Yield of -Oxidation. . 2078.15. Oxidation of Unsaturated Fatty Acids . 2078.16. Isomerization of the cis-Double Bond. . 2088.17. Summary 209

CHAPTER 9. The Citric Acid Cycle... 2119.1. Balancing the overall Reaction.. . 2119.2. Working Backwards. 2129.3. Working Forward . 2179.4. Ambiguities. . 2229.5. Oxidative Decarboxylation of Isocitric Acid. 2229.6. The Glyoxylate Cycle. . 2239.7. Summary. . 226

CHAPTER 10. Enzymes of the Citric Acid Cycle.... 222810.1 Citrate Synthase... 22910.2. Citrate (pro S) Lyase... . 23010.3. ATP Citrate Lyase. 23210.4. Some Thermodynamic Considerations.... 23210.5. Aconitase. 23310.6. Isocitrate Dehydrogenase. 23410.7. The Oxoglutarate Dehydrogenase Complex. . 23810.8. Succinyl-CoA Ligase . 24210.9. Succinic Dehydrogenase... 24510.10. Fumarase..... 24510.11. Malate Dehydrogenase.. 24510.12. The Citric Acid Cycle. 24510.13. Isocitrate Lyase.. 24610.14. The Glyoxylate Cycle . 24710.15. Production of Energy from Fatty Acids . 24810.16. Summary . 249

Table of Contentsvi

CHAPTER 11. Citric Acid Cycle Stereochemistry... 25111.1. Some Experimental Results.. . 25111.2. Misinterpretation of the Labeling Results. . 25211.3. Fixation of CO2 by Animal Cells. 25411.4. The Role of Biotin in Carboxylation Reactions.. . 25511.5. The Stereochemistry of Citrate Synthesis... . 25611.6. Conversion of Citric Acid to Isocitric Acid.. 25811.7. Conversion of Succinic Acid to Fumaric Acid.. . 26011.8. Conversion o Fumaric Acid to S-Malic Acid. 26211.9. Summary.... 26311.10. The Big Picture. 264

CHAPTER 12. Glycolysis.... 226712.1. The Structure of Glucose . 26712.2. A Balanced Equation for Glycolysis Tells Us Three Things. . 27012.3. Enediol Isomerization of Glucose.. . 27112.4. Cleavage of Fructose. . 27312.5.Cleavage of Fructose 1,6-bis-Phosphate.. . 27412.6. Isomerization of Glyceraldehyde 3-Phosphate.. 27412.7. Oxidation of Glyceraldehyde 3-Phosphate Requires Pi.. . 27512.8. Oxidation of Glyceraldehyde 3-Phosphate Set Up a

Substrate Phosphorylation . 27612.9. Conversion of Glyceric Acid 3-Phosphate to Pyruvic Acid.. 27712.10. The Thermodynamics of PEP Induced Substrate

Phosphorylation. 27812.11. Conversion of Fructose 1,6-bis-Phosphate to Pyruvic Acid . 27812.12. Conversion of Fructose 1,6-bis-Phosphate to Glucose. 27912.13. The Glycolytic Pathway.. . 28112.14. Reversal of Glycolysis.. 28212.15. Conversion of Pyruvic Acid to PEP... . 28312.16. Isomerization of Glyceric Acid 3-Phosphate.. . 28512.17. The Phosphoglycerate Mutase Reaction.. . 28512.18. Formation of 2,3-bis-Glycerate Phosphate.. 28712.19. Oxidation of Glyceraldehyde 3-Phophate.. . 28812.20. Production of Energy from Glucose . 28912.21. Glycolysis in the Absence of Oxygen. 29112.22. The Reactions of Pyruvic Acid. . 29312.23. Gluconeogenesis from Lactate.. . 29412.23. Summary . 294

The Chemical Basis of Metabolism vii

CHAPTER 13. Oxidative Deamination of Amino Acids.. 229613.1. Oxidative Deamination by Hydride Transfer... . 29613.2. Oxidative Deamination b y Electron Transfer... 29813.3. Pyridoxal Phosphate is the Acceptor for Oxidation by Electron Transfer. 29813.4. Transamination. 30213.5. Glutamate Dehydrogenase Plays a Key Role

in Transamination.. . 303The Urea Cycle. 305

13.6. Urea is Formed by Hydrolysis of Arginine.. . 30613.7. Conversion of Ornithine to Citrulline. 30813.8. Formation of Carbamoyl Phosphate... 30813.9. Formation of Citrulline from Carbamoyl Phosphate and

Ornithine.. 30913.10. Conversion of Citrulline to Arginine Requires Aspartic

Acid and ATP... 31113.11. Argininosuccinic Acid is Converted to Arginine

by Elimination. 31513.12. Synthesis of Argininosuccinic Acid Requires Two ATPs. . 31613.13. The Urea Bicycle. . 31813.14. Glutamate Dehydrogenase Plays an Important

Role in the Urea Cycle.. . 31913.15. Summary . 322

CHAPTER 14. Oxidation of Amino Acid Carbons-1Alanine, Aspartic Acid, Asparagine,Glutamic Acid, Glutamine ,Proline .. 3324

14.1. Introduction. . 32414.2. Asparagine and Glutamine.. . 32514.3. Proline.. 32514.4. The POP Triangle . 33014.5. Transamination.. 33114.6. The Glutamate Dehydrogenase Cascade. 33214.7. Essential Amino Acids . 33414.8. Synthesis of Proline from Glutamic Acid.... 335

Table of Contentsviii

14.9. A Map for Biosynthesis of Alanine, Aspartic Acid,Asparagine, Glutamic Acid, Glutamine, andProline 337

14.10. Summary of the Biosynthetic Pathways to Alanine,Aspartic Acid, Asparagine, Glutamic Acid, Glutamine,and Proline. 340

14.11. Summary 341

CHAPTER 15. Oxidation of Amino Acid Carbons-2Serine, Glycine, and Threonine.. . 3342

Serine . 334315.1. Oxidation.. . 34315.2. Decarboxylation. . 34515.3. Deamination of Serine . 34715.4. Serine Contributes to the One Carbon Pool.. 34715.5. Tetrahydrofolic Acid is the One Carbon Acceptor. . 34815.6. Formaldehyde Reacts with Tetrahydrofolic Acid.. . 34915.7. Summary of Serine Metabolism.. 351

Glycine.. 335215.8. Decarboxylation and Deamination. . 35215.9. Formation of N5,10-Methylene-THF from Glycine

Requires Four Different Protein.. 35315.10. Summary of Glycine Metabolism. 355

Threonine.. . 335615.11. Structures. 35615.12. Direct Oxidation of Threonine . 35615.13. The Role of PALP.. . 35715.14. Partial Summary of Threonine Catabolism. . 35915.15. The -Amino- -ketobutyric Acid Pathway. . 36015.16. The Threonine Aldolase Pathway... 36115.17. The -Ketobutyric Acid Pathway... 36215.18. Summary of Threonine Metabolism. . 362

CHAPTER 16. Metabolism of Propionyl-CoAThe Role of Vitamin B12 Coenzyme. 3365

16.1. Carboxylation of Propionyl-CoA.. . 36516.2. Conversion of Methyl Malonyl-CoA to succinyl-CoA. . 36616.3. Vitamin B12 Coenzyme.. . 36816.4. The Role of Cobaltic Ion . 37016.5. Formation of Vitamin B12 Coenzyme from Vitamin B12.... 37416.6. How B12 Coenzyme Works . 376

The Chemical Basis of Metabolism ix

16.7. A Generalized Mechanism for B12 Coenzyme Reactions . 37816.8.Rearangement of Methyl Malonyl-CoA to Succinyl-CoA.. . 37816.9. Summary of Threonine Metabolism.. . 379

CHAPTER 17. Branched Chain Amino Acids:Isoleucine, Leucine and Valine . 3381

Isoleucine. 338217.1. Transamination.. 38217.2. Oxidative Decarboxylation.. . 38317.3. Oxidation of Saturated Carbons. 38317.4. Hyperconjugation (No-bond Resonance).. . 38417.5. A Nomenclature Problem. The EZ System .. . 386

17.6 The E/Z Nomenclature System for NamingGeometrical Isomers.. . 386

17.7. Reactions of Tiglyl-CoA. . 38717.8. Summary of Isoleucine Catabolism.. . 38917.9. A Metabolic Map for Isoleucine Catabolism . 39017.10. Ketogenesis. . 390

Leucine.. 339317.11. Early Reactions of Leucine. . 39317.12. Carboxylation of -Hydroxy- -methylbutyryl-CoA.. . 39517.13. Summary of Leucine Catabolism. 39717.14. A Metabolic Map of Leucine and Isoleucine Catabolism. 398

Valine 39817.15 Early Reactions of Valine.. 39817.16. Reactions of -Methyl- -hydroxypropionyl-CoA. . 40017.17. Different Kind of Oxidative Decarboxylation. . 40217,18. Summary of Valine Catabolism. . 40317.19. A Metabolic Map for Catabolism of Leucine,

Isoleucine and Valine.. 40517.20. Summary . 405

CHAPTER 18. THE SULFUR AMINO ACIDS.. . 440718.1. Bonding of Sulfur Compounds. 40718.2. Some Rules . 409

Cysteine.. . 441018.3.Transamination. . 41018.4. Removal of Sulfur from Mercaptopyruvic Acid.. . 41018.5. Removal of Sulfur by -Elimination... . 41118.6. Removal of Sulfur by Substitution . 412

Table of Contentsx

Methionine . 441518.7. Transamination . 41518.8. Removal of Sulfur by -Elimination . 41618.9. Removal of Sulfur by -Elimination.. . 417

The Methylation Cycle. 441918.10. Methylation.. . 41918.11 The Role of Tetrahydrofolic Acid. . 42018.12. S-Adenosylmethionine (SAM) . 42318.13. Synthesis of SAM . 42418.14. SAM as a Methylating Agent... 42518.15. Summary of the Methylation Cycle.... 42718.16. The Role of Cobalamin in the Methylation Cycle... 42818.17. The Methylation Cycle in Animal Cells. . 43018.18. Conversion of Methionine to Cysteine. . 43118.19. Cleavage of Cystathionine. . 43318.20. Summary of Conversion of Methionine to Cysteine.. . 435

Some Odds and Ends. 443518.21. Formation of Disulfide Bonds.. . 43518.22. A Clear Case for -Elimination.. . 43618.23. Some Properties of the Cyclopropane Ring. . 43618.24. Aminocyclopropane Carboxylic Acid is Reactive. . 43818.25. Lactonization.... 43918.26. Summary . 439

CHAPTER 19. BASIC AMINO ACIDS-1 Arginine and Lysine. . 4444

19.1. How We Are Going to Proceed.. . 444Arginine.. . 4445

19.2. The Initial Step.. . 44519.3. Reactions of Ornithine . 44619.4. Reactions of Glutamic Acid Semialdehyde and

5-Amino-2-oxo-pentanoic Acid.. . 44619.5. A Tentative Map for Arginine Catabolism.. . 44919.6. Predictions . 44919.7. Enzymes for Arginine Catabolism . 45019.8. Catabolism of Arginine in Bacteria. . 45319.9. A Map for Arginine Catabolism . 456

Lysine. 45719.10. A Tentative Map for Catabolism of Lysine.. . 45719.11. Removal of the -Amino Group.. 45819.12. Removal of the 6-Amino Group.. . 462

The Chemical Basis of Metabolism xi

19.13. The Role of Pipecolic Acid.. . 46819.14. Decarboxylation of -Ketoadipic Acid.. 47119.15. Decarboxylation of -Keto- -aminocaproic Acid.. . 47219.16. Catabolism of Glutaryl-CoA . 47419.17. Final Map for Lysine Catabolism.. . 47519.18. Summary of Lysine Catabolism.. . 476

CHAPTER 20. BASIC AMINO ACIDS-2Histidine.. . 4479

20.1. The Structure of Histidine . 47920.2. The Initial Reactions of Histidine Catabolism.. . 48120.3. The Solvent Can Have an Important Effect on a Reaction... 48320.4. Ring Opening.. . 48320.5. The Role of Tetrahydrofolic Acid.. . 48520.6. The Reaction of Tetrahydrofolate with

N-formimidoyl-glutamate Sets Up a New Pathway.... 48920.7. A Map of Histidine Catabolism.. 49120.8. What Does the Map Tell Us?. . 49320.9 Summary . 494

CHAPTER 21. AROMATIC AMINO ACIDS-1 Phenylalanine and Tyrosine.... 4495

21.1. Structure of Phenylalanine and Tyrosine... 49521.2. Molecular Orbitals of Benzene. 49521.3. Reactions with Aromatic Rings.. . 49721.4. Monooxygenase Reactions. . 49721.5. Dioxygenase Reactions.. 49921.6. Some Facts.. 50121.7. Transamination of Phenylalanine. 501 21.8. Conversion of Phenylalanine to Tyrosine. . 50221.9. Transamination of Tyrosine . 50521.10. Formation of Homogentisic Acid . 50621.11. The NIH Shift in Animal Cells.. . 50821.12. Catabolism of Homogentisic Acid.. . 51221.13. A Map for Catabolism of Phenylalanine and Tyrosine.... 51321.14. Summary . 514

Table of Contentsxii

CHAPTER 22. AROMATIC AMINO ACIDS-2Tryptophan Stage-1. Removal of the Side Chain.. 5518

22.1. Tryptophan is Aromatic. . 51822.2. Constructing a Map Without Any Previous Knowledge of

Tryptophan Catabolism . 51922.3. Predictions . 52422.4. Enzymes Catalyzing Tryptophan Catabolism. 52422.5. A New Map of Tryptophan Catabolism . 528

A Few Final Details... 52922.6. Kynurenine 3-Monooxgenase... 52922.7. Transfer of the Formyl Group from

Formylkynurenine to FH4.. . 53122.8. Formation of Indole Acetic Acid.. 53322.9. Indole Acetic Acid is a Plant Growth Factor. 53322.10. A Final Map for Tryptophan Catabolism... 53422.11. Summary . 535

CHAPTER 23. AROMATIC AMINO ACIDS-3Tryptophan Stage-2. Catabolism of Anthranilic Acid

and 3-Hydroxyanthranilic Acid . 5537Ortho Oxidation in Bacteria . 538

23.1. Deamination of Anthranilic Acid. . 53823.2. Decarboxylation of Protocatechuic Acid. . 54023.3. Opening the Catechol Ring.. . 54223.4. Catabolism of Muconic Acid . 54223.5. Formation of Muconolactone.. . 54423.6. Reactions of Muconolactone. . 54523.7. 3-Oxo-adipate Acetyl-CoA Transferase. . 54723.8. A Map for Ortho Oxidation of Anthranilic Acid

in Bacteria . 54823.9. A Map for Tryptophan Catabolism in Bacteria. . 549

Meta Oxidation.. . 55023.10. What Does the Presence of Pipecholic Acid and

Quinolinic Acid in Urine Tell Us?. 55023.11. Predictions. . 55923.12. A Map for Tryptophan Catabolism via Meta Oxidation.... 56023.13. Summary . 561

The Chemical Basis of Metabolism xiii

PART 3. ANABOLIC PATHWAYS

CHAPTER 24. STARCH AND GLYCOGEN.. . 5564 24.1. The Structure f Starch and Glycogen... 56424.2. Phosphorolysis of Glycogen . 56724.3. Conversion of G-1-P to G-6-P . 56824.4. Biosynthesis of Glycogen Does Not Occur by Reversal of

Phosphorolysis.. 57024.5. UDPG . 57024.6. Adding Glucose Residues to the Existing Glycogen Chains.. . 57124.7.Branching.. . 57324.8. Summary.... 576

CHAPTER 25. BIOSYNTHESIS OF FATS.. 557725.1. Elongation of the Fatty Acid Chain. . 57725.2. Fatty Acid Synthase 57925.3. Introduction of Double Bonds. . 58225.4. Synthesis of Triglycerides. 58725.5. Naming Triglycerides. . 59525.6. Summary.... 595

EPILOGUE... .596INDEX 5597

Preface ix

PREFACE

It has always bothered me that we make students memorize metabolicpathways by rote. I first encountered this practice as a student in aphysiology course. I had to learn the names of all the intermediates inglycolysis, fatty acid oxidation and the Citric Acid Cycle so we could talkabout interesting things like diabetes and insulin. What's more I had tolearn these names without any structural formulas. I did as I was told. Wedid discuss diabetes and other interesting things, but I had no idea what allthose names in the metabolic pathways meant.

The next year I took my first biochemistry course. This time I had tomemorize the structures that go with the names. While this was still anarduous exercise, the structures at least helped establish some biochemicalrelationships that were not obvious from the names alone.

Then a remarkable thing happened. I discovered a book called"Principles of Ionic Organic Reactions" by E.R. Alexander. As I studiedthis book, organic chemistry began, for the first time, to make sense to me.The reactions I had been memorizing became logical and predictable.When Ingold's, "Structure and Mechanism in Organic Chemistry", waspublished, I devoured it.

My thesis, to the amazement of my supervisor, was full of curvedarrows and flying electrons.

"Where did you learn to do that," he asked?Soon I began applying what I had learned to biochemical reactions,

and I discovered mechanisms could be used to predict most of theintermediates in the metabolic pathways. It was great fun! I tried to tellother graduate students and my professors about this exciting game, butthey weren't very interested. In retrospect, I realize now they didn't knowwhat I was talking about since their training had been in classical organicchemistry, and at that time application of mechanisms to organic reactionswas in its infancy.

When I started teaching in 1956 at New York University School ofMedicine, I tried to show graduate students how they could usemechanisms to generalize and predict biochemical reactions. It soonbecame clear there was a problem. Even though the students had studiedorganic chemistry, they didn't know how to "push electrons". So I began

The Chemical Basis of Metabolismx

developing a course to teach students with a minimal background inchemistry how to do this.

The course was based on a series of problem sets containing questionsdesigned to take the student through important principles of mechanisticorganic chemistry in a carefully planned way. After some lectures (4 hours)in which we discussed general mechanisms for the three reaction types(substitution, elimination and addition), the students began applying theseprinciples to simple examples from classical organic chemistry. Then,when the fundamental mechanisms were clear, they began applying themto biochemical pathways.

At each class session the students went to the board individually and,without notes, wrote out their answer to a question from the problem set. Isat in the back of the room and listened. If the student made a mistake andno one objected, I intervened. If the student missed a possibility I hadconsidered and no one else suggested it, I intervened. Often students wrotemechanisms that had escaped me. Sometimes these were reasonable,sometimes they were not, but they always served as an interesting point fordiscussion.

At first, students were apprehensive about going to the board. Theywere afraid their mistakes would brand them as "stupid". They were notused to performing in front of an audience. They were scared! Slowly theylost these inhibitions, and many lively discussions ensued. Mostimportantly the students were participating instead of listening to medrone on.

In the end most students liked the course. And most performedadmirably. Most said they had to work awfully hard, but they learned a lot.

"Metabolism is so logical", they said, "why isn't it taught this way inregular biochemisty courses?"

"Because most students don't know enough about mechanisms to usethis approach", I replied. "Look how long it took for you to learn how topush electrons in the right direction".

"Well", they said, "this is an interesting way to learn metabolism, butwe need a guide to get us started".

There are excellent books on Bio-organic Chemistry and on theMechanisms of Enzyme Action. There is an occasional chapter in abiochemistry text on the "logic" of metabolic pathways. I am not aware ofany book dealing exclusively with the use of mechanistic organic chemistryto predict the intermediates of the major metabolic pathways. That is whatthis book is about.

Electron pushing follows fairly simple rules. We will present the rulesas such, and we will discuss the basis for these rules. Once you know how

Preface xi

to push electrons in the proper direction, the logic of metabolic reactionsemerges automatically. Of course, there is much more to metabolism thanthis. For example, electron pushing will not tell you which of severalpossibilities actually occur. Nor does it tell you how metabolic pathwaysinteract or how they are controlled. That is where biochemistry comes in.There are lots of good books dealing with the biochemistry of metabolism.However, electron pushing will provide you with insight that is notpossible by simply memorizing pathways.

With all the excitement surrounding Molecular Biology, biochemicalmechanisms are not as fashionable as they once were. Still from thestudents point of view the need for understanding the chemical basis ofmetabolism is as great as ever. Metabolism is the motor that drives all livingorganisms. Students of the Life Sciences need a working knowledge ofmetabolic pathways.

Electron pushing provides a logical approach to metabolic reactions.Not only does it provide new insights into metabolism, it is essential forunderstanding modern theories of enzyme action. I hope to show you thatthere is a chemical logic to metabolism and that application of this logic isboth useful and fun.

Robert Warner Chambers

Professor of Biochemistry and Molecular BiologyDalhousie UniversityHalifax. Nova ScotiaCANADA

The Chemical Basis of Metabolism 1

INTRODUCTION

The essence of life is complicated chemistry and nothing moreJames D. Watson

The Unity of BiochemistryBiochemists like to talk about the unity of biochemistry. The

concept arose during the golden years of research on metabolism. As thepieces of the complicated puzzle were put together, it became clear thatorganisms from unicellular bacteria to Homo sapiens utilize the samemetabolic reactions. There are differences to be sure, but the similaritiesare remarkable.

Metabolism is a series of linked chemical reactions designed to produceenergy for cell function or to synthesize components that are necessary forcell structure. These reactions involve making and breaking carbon-carbon,carbon-oxygen, carbon-nitrogen, carbon-sulfur, carbon-hydrogen andphosphorus-oxygen bonds. Metabolic reactions follow a set of rules. If weknow the rules and apply them carefully, we can predict all sorts of things.This can get pretty complicated, but there are surprisingly fewfundamental rules we must know in order to predict most of theintermediates that occur in metabolic pathways. In this book we willdevelop these rules and apply them to biochemical reactions. In theprocess, we will build complete metabolic pathways that include not onlythe reactions involved, but the enzymes that catalyze them. Each chapter inParts 2 and 3, contains a metabolic map that summarizes our conclusionsand agrees, for the most part, with the accepted pathway.

Electron Pushing and Metabolic ReactionsWe are going to use a technique known as electron pushing.

Electron pushing (EP) enables us to keep track of the electrons involved inmaking and breaking chemical bonds. Organic chemists use EP to design aroute to some synthetic compound. They also use the technique to explainthe experimental observations they have made.

We can do the same thing if we know the EP rules. In fact our task ismuch simpler than that of the organic chemist. The pathway an organicreaction takes is controlled by the choice of reactants and reactionconditions. We have no control over biochemical reactions under normalin vivo conditions. If A and B yield C under one set of conditions and D

The Chemical Basis of Metabolism2

under another set of conditions, the chemist can, in many cases, controlwhich reaction occurs by selection of the appropriate solvent, temperatureand catalyst. In the body, biochemical reactions go on in an aqueousenvironment at 37. So the solvent is fixed and under normal conditionsthe temperature is constant. Enzymes catalyze the metabolic reactions and,in most cases, determine the mechanism. Since we cannot controlbiochemical reactions, we do not need to know how various conditionsaffect the outcome of a reaction.

It might appear that we have to memorize how each enzyme works inorder to understand how metabolic reactions work. In fact, metabolism isusually taught this way. However, this classical approach obscures thechemical logic that underlies metabolism. As we shall see, with a fewexperimental facts, some EP rules and careful thought, we can predict mostof the intermediates occurring in a metabolic pathway.

PrerequisitesIn order to understand the principles we are going to develop, it is

necessary to have some background in organic chemistry. You need toknow the structure of the functional groups and some of their generalproperties. You should be familiar with atomic and molecular orbitals,resonance, -bonds, electronegativity, and the like. We will discuss thesethings, but these discussions are intended to be a review of concepts youhave encountered before.

We are going to talk about stereochemistry quite a bit. There are tworeasons for doing this. First, molecules are not flat, as we write them onpaper. They are three-dimensional objects, each with a characteristic shape.It is very important to have a mental picture of these shapes. Many of thestructures in this book are written in a three dimensional representation tohelp you develop this skill. Second, the shape of a molecule determines itsstereochemistry and its stereochemical predisposition. Stereo-selectivity isone of the most remarkable features of metabolic reactions. Enzymes,because they are dissymmetric reagents, are able to distinguish betweengroups that are chemically equivalent and in so doing direct the formationof a particular isomer. Understanding the principles behind this selectivityis crucial to understanding the chemistry of metabolic reactions. We willdevelop these principles as we go along, but it will be helpful if you knowthe difference between enantiomers and diasteriomers as well as betweencis and trans isomers. We will review of these things in Part 1 for those whoare a little rusty.

The Chemical Basis of Metabolism 3

You must understand that we are going to use EP to write a reasonablemechanism for a particular reaction. Our purpose is to predict the productof the reaction. We are not concerned with whether the mechanism wehave written is the one that actually occurs. It turns out that there areusually several mechanisms that lead to the same product. In fact, youcannot prove a given mechanism directly. One tries to rule out as manyalternatives as possible to arrive at the most likely explanation.Experimental data is required to do this. In biochemistry, this comes underthe heading of Enzyme Reaction Mechanisms. We are not going to talkmuch about this; there are already many good books on the subject. Ourgoal is to predict, the intermediates in a metabolic pathway withoutknowing anything about the enzymes that are involved. This simplifies ourtask.

If there are several mechanisms that correctly predict the products of areaction we have only to find one of these, and it need not be the one thatactually occurs. This approach may be disturbing at first. How can we use amechanism that does not actually occur to predict what does occur? Howcan we predict the product of an enzyme-catalyzed reaction if we don'tknow anything about the enzyme involved? Application of EP tometabolic reactions leads to a realization that there is a chemical logic tothe metabolic pathways that has nothing to do with enzymes. This doesnot in anyway impugn the importance of enzymes. Enzymes are a criticalpart of metabolism, but they do not alter the fundamental chemistrydriving the reaction.

Although we are not going to talk enzyme mechanisms per se, we willuse two enzyme data banks1 to check our predictions or, in some cases, toresolve ambiguities that arise in developing a metabolic pathway with EPalone. These data banks list all the reactions where the enzyme activity hasbeen detected. We can check any reaction we predict by looking in thesedata banks to see if the enzyme catalyzing the reaction has been described.This is also useful in determining the exact nature of any coenzyme thatmight be involved and in resolving stereochemical ambiguities. However,we will not rely on these sources for any of our initial predictions.

It will make things easier if you have had or are taking a generalbiochemistry course. However, it is essential to keep an open mind. We 1 Enzyme Nomenclature recommended by the Nomenclature Committee of theI n t e r n a t i o n a l U n i o n o f B i o c h e m i s t r y a n d M o l e c u l a r B i o l o g yhttp://www.chem.qmul.ac.uk/iubmb/. Academic Press publishes this in book form.Brenda (http://www.brenda-enzymes.info/ ) is a similar data bank, but contains moredetailed information.

The Chemical Basis of Metabolism4

want to see where EP will take us. If, instead, you try to force reactions tofit pre-conceived notions about a pathway, you may fail to see theimportant logic we are trying to develop. You may also miss someimportant alternatives that EP predicts, and one or more of thesealternatives may actually occur in a certain organism under certainconditions.

One of the most common mistakes you can make is to look at thereactants and products of a pathway and try to figure out all theintermediates, in your head, without pushing electrons. This hardly everworks. The beauty of EP is that it can predict each successive reaction in anorderly fashion. One begins by analyzing possible reactions that can occurwith a given compound. Then one tries to make a "best guess" as to whichis the most likely. You will be surprised how often you get it right if youproceed in a stepwise manner and apply the principles correctly.

Many students find EP very confusing at first. If you do, do notdespair. You cannot play concertos until you have mastered certainfundamental techniques of music. Learning those fundamentals takes timeand hard work; there is no substitute. Learning EP is like learning to playscales. Anyone can learn to do it though it maybe easier for some than forothers. Once you become competent with EP, you can make reasonablepredictions and gain new insight into metabolic reactions.

5

The Chemical Basis of Metabolism

PART 1

FUNDAMENTALS

Start with fundamentals and from this build generalities

Alexander Graham Bell

Part 1. Fundamentals6

1

SOME PRINCIPLES OF BONDING

1.1. Breaking Covalent BondsMetabolic reactions, like all organic reactions, involve breaking and

remaking covalent bonds. Since covalent bonds are formed by sharingelectron pairs between the nuclei of two or more atoms in a molecule, itfollows that metabolic reactions proceed by redistributing electron pairsbetween atoms. This can occur in two fundamentally different ways.Homolytic splitting breaks the electron pair so that each of the atomsparticipating in the covalent bond comes away with one unpaired electron.Heterolytic fission cleaves the bond so that one atom comes away with bothelectrons. For example,

A:B = A. + .B A:B = A+ + :B-Homolytic Heterolytic

Homolytic reactions produce free radicals. Heterolytic reactionsproduce ions. If the starting molecule is neutral, then a cation and an anionare produced, as in the example above. If the starting molecule is charged,then new ions may be produced. For example,

A:B+ = A+ + :B

Regardless of the exact nature of the charges, electrical neutrality mustbe preserved so there is no net change in charge in the reaction. Heterolyticreactions are often called ionic reactions because they involve ions in somemanner. Most metabolic reactions proceed in this way. Very few importantreactions proceed homolytically. Unless I point out differently, we willassume that the reactions proceed heterolytically.

Ionic organic reactions occur by one of three mechanisms: substitution,addition or elimination. Before we consider these (Chapter 3), it isimportant to understand some basic concepts of bonding. To do achievethis we need to discuss some elementary molecular orbital theory. These

Chapter 1. Some Principles of Bonding 7

principles are important for understanding the 3-dimensional structure ofvarious compounds and for predicting the products of a given reaction. Inthis discussion, I have simplified things in order to make the concepts clear.Even without the rigor of quantum mechanics, we can make someimportant predictions and draw some interesting conclusions.

1.2. The Heisenberg Uncertainty PrincipleChemical reactions are governed by energetics. In general, the higher

the energy level of an electron pair, the more reactive it will be. Quantummechanics tells us that electrons are located at discrete energy levels. Weneed to know what these levels are, at least in relative terms, to predictwhich reactions will occur and which will not. However, the HeisenbergUncertainty Principle says that we cannot know the energy and theposition of an electron at the same time. This fundamental principle ofquantum mechanics is not a statement of incompetence. It says that anyattempt to measure the position of an electron changes its momentum(kinetic energy) in an unpredictable way. If we want to know the energy ofan electron, we must settle for some uncertainty in its position. Soundshopeless, but it is not. The position of an electron can be expressed as theprobability of finding it in a finite space called an orbital. We can representthe probability as an electron density with a size and shape that can becalculated from quantum mechanics. Or we can represent the probabilityas a line tracing the outer boundary of the electron density. The chargedensity diagram is more realistic, but the line diagram is usually moreconvenient. It is important for us to have a mental picture of the variouskinds of orbitals because they define many of the properties of an atom or amolecule including its stereochemistry.

1.3. The Shape of OrbitalsAn electron has a measurable mass so it has the properties of a particle

like a very tiny grain of sand. However, quantum mechanics tells us that theproperties of an electron can also be described as wave-like. We must acceptthat an electron has these two properties and is not like any of the objectswe observe around us. Interestingly, we can guess the shape of some simpleatomic orbitals from the wave-like character of the electron. If we draw twoidentical wave functions and adjust them so they overlap, they look likethis1

1This is an oversimplification. Waveforms do not look quite like this. We will have tochange these shapes a little bit, but the simplification does lead to the correct conclusion.


+ =

Fig. 1.1

We can see two distinct shapes: a figure 8 lying on its side, which ismade up of two circles. If we separate the circles, each circle represents theboundary of an s orbital. The figure 8 lying on its side represents a p orbital.The size of the orbital as well as its shape is related to the energy level of theorbital. For example, a 1s orbital is at the first quantum level. The electronsin this kind of orbital are closer to the nucleus than the electrons in a 2sorbital, which is at the second quantum level. It follows that a 1s orbital issmaller than a 2s orbital.

We should think of these orbitals, as well as all others, as a negativelycharged cloud. We can represent this pictorially with charge densitydiagrams. The s orbitals at the first and second quantum levels look likethis:

1s orbital 2s orbital

Fig. 1.2

These charge density diagrams represent the probability of finding anelectron with a particular energy in a given space. We should rememberthese are three-dimensional objects so s orbitals are ball-shaped.

1.4. Molecular OrbitalsMolecules form by combining electrons from two or more atoms.

For example, sharing an electron pair between two hydrogen nuclei formshydrogen gas, H: H. This means the atomic orbitals of hydrogen mustcombine somehow to form a molecular orbital that can accommodate theshared electron pair. We can see how this occurs by a simple algebraicmanipulation. Let A be the atomic orbital of one hydrogen atom, HA, and

B be the atomic orbital of the other, HB. Let = a molecular orbitalformed by combining two atomic orbitals. We can combine these atomicorbitals in two ways:


A+ B or A - B

However, a molecular orbital, which represents the probability offinding an electron in a defined space, is actually the square of the atomicwave function2 so we can write:

1 = ( A+ B)2 = A

2 + 2 A B + B

2

2 = ( A - B)2 = A

2 - B2

1.5. Bond and Antibonding OrbitalsThe difference between 1 and 2 is obvious. One contains an

overlapping region, denoted by 2 A B, while the other has no overlap. If weexamine the charge density diagrams for these orbitals, we find 1 lookslike this:

Fig. 1.3

This molecular orbital is called a bonding orbital because of the overlapdensity. Orbitals of this kind are often called orbitals. When electronsoccupy this kind of molecular orbital a so-called -bond is formed.

2 This simplification is useful for our purposes. Actually the molecular orbital is the squareof the wave function integrated over all space. Since the probability of finding an electronin such an orbital is 1.0, the function we write for a particular orbital must be normalizedto 1.0. However, these details need not concern us here.


The other orbital ( 2) looks like this:

Fig. 1.4

This molecular orbital is called an a n t ibonding orbital becauseelectrons in this kind of molecular orbital tend to repel each other asindicated by the region between the two nuclei where the electron densityis zero. An antibonding orbitals of this kind are often designated as a *orbital3.

We should think of these two molecular orbitals, and *, assuperimposed upon one another. So the area of the antibonding orbitalwhere the probability of finding an electron is zero tends to cancel theoverlapping area of the bonding orbital. The result is that disruptiveinfluence of electrons in an antibonding orbital tends to cancel theattractive forces of electrons in the corresponding bonding orbital. As weshall see this will lead us to some important conclusions.

1.6. Molecular Orbital Energy Diagram for HydrogenSince the electrons in 1 "feel" the attractive forces of both atomic

nuclei we can guess, without making any calculations, that the bondingmolecular orbital, 1, will lie at a lower energy level than the antibondingorbital, 2. An energy diagram of these orbitals looks like this4:

energy

2

1

Fig. 1.5

In molecular orbital theory the energy levels of the different molecularorbitals (MO) are calculated without considering the electrons that aregoing to occupy the orbitals. Then the available electrons are distributed

3Antibonding orbitals are molecular orbitals formed by combining atomic orbitals in aparticular way. There are no antibonding atomic orbitals.4Actually the energy levels do not split equally, but this refinement need not concern ushere.


into as many orbitals as needed to accommodate them. This is known asthe Aufbau Principle5.

We know that each hydrogen atom contributes one electron to thehydrogen molecule, H2. Since we have concluded that 1 lies at a lowerenergy level than 2 both electron go into the bonding orbital, 1, leavingthe antibonding orbital, 2, empty6. Quantum mechanics also tells us theelectrons occupying the same orbital must have opposite spins. Ourdiagram now looks like this:

2

1

A(1s) B(1s)

ener

gy

(bonding)

(antibonding)

bond

electron pair

Fig. 1.6

From this simple discussion, we predict that hydrogen atoms willcombine to form a hydrogen molecule, and that the molecule will be linear.We indicate this as H-H. The bond formed by overlapping two s orbitals iscalled a sigma bond and it is designated by . The antibonding orbital isdesignated *.

1.7. MO Energy Diagram for HeliumNow let us predict some properties of helium (He). A helium atom has

two electrons. In order to form a helium molecule, He2, four electrons mustbe accommodated. This would fill both 1 and 2; the energy diagramwould now look like this:

5 Aufbau is German for building.6 This is true as long as there is a significant difference in the energy levelsof the empty orbitals. If the orbitals are at the same energy level(degenerate orbitals) then each of the degenerate orbitals will have oneelectron before any one of them has paired electrons. This is known asHunds Rule.


2

1

A(1s) B(1s)

ener

gy

(1 ) bonding

(1 ) antibonding

electron pair

Fig. 1.7

From this diagram and what we have just learned about the effect ofantibonding electrons on bonding electrons, we can predict thisarrangement is unstable. The bonding tendency of the electrons in 1 iscancelled by the antibonding properties of the electrons in 2. Thus, we donot expect He to form He-He. Our prediction is correct. Under ordinaryconditions helium exists only in its atomic form.

1.8. AcetyleneNow let us construct a molecule of acetylene, which has two carbons

and two hydrogens. Carbon is a first row element. It has six electrons;hydrogen has one so we must place fourteen electrons in appropriateorbitals to form acetylene. Four of these electrons go into the twomolecular orbitals, 1 and 2, that occupy the first quantum level. Theantibonding properties of the electron pair in 2 tend to cancel thebonding properties of the electrons in 1. Therefore, these kernel electrons,as they are sometimes called, do not participate directly in bonding7. Ourtask now is to construct the orbitals at the second quantum level that willhold the remaining ten (valance) electrons and to see what conclusions wecan draw from this. Let us begin by joining two carbon atoms.

Two new molecular orbitals, one bonding and one antibonding, can beformed by combining the 2s atomic orbitals, but these can only account forfour of the remaining eight electrons. Clearly some other atomic orbitalsmust be involved.

7Kernel electrons do affect the bond strength indirectly by shielding the bondingelectrons (valence electrons) located at the second quantum level from the attractive forceof the nucleus. This detail need not concern us.


1.8.1. Acetylene Uses sp Hybrid Orbitals to Form -BondsEarlier we deduced there should be two kinds of orbitals, one with

spherical symmetry (s orbitals) and one that looks like a figure eight lyingon its side (p orbital). At the first quantum level, p orbitals lie at very highenergy levels so they are not used. At the second quantum level the three porbitals, which are only slightly higher in energy than the 2s orbital. Allthree 2p orbitals lie at exactly the same energy level. Unlike s orbitals,which are spherical, p orbitals have direction. The three p orbitals areoriented as shown below:

px py pz

Fig. 1.8

It is possible to form a MO by overlapping p orbitals with the samesymmetry:

C C C C

+ = =

px px

Fig. 1.9

However, a better overlap can often be established if an s atomic orbitalis mixed first with one of the p orbitals to form an sp hybrid orbital. Sincetwo atomic orbitals are used, two hybrid orbitals must form. We can makea pretty good guess at the shape of these orbitals:

=2px

2s

sp hybrids

Fig. 1.10


It is important to remember that these two orbitals belong to a singlecarbon atom, as shown in Fig. 1.108.

Let us place the new hybrid orbitals on our energy diagram:

Atomic Orbitals Atomic OrbitalsMolecular Orbitals

A(2s)

A(sp)

(2px,y,z) B(2px,y,z)

B(sp)sp hybrids sp hybrids

B(2s)

2

1

A(sp) B(sp)

Hybrid Orbitals

(2py,z) (2py,z)

Hybrid Orbitals

First Quantum Level

Second Quantum Level

ener

gy

Fig. 1.11

The highly directional nature of the hybrids increases the overlap withsome other atom. Let us join the carbons first. The bonding orbitals looklike this:

+ =+

Ahy1 Ahy2

C CC

C

Bhy1 Bhy2

Fig. 1.12

Now we can add the hydrogens. The bonding orbitals look like this:

+ ++ =

H H H HC C C C

Fig. 1.13

8There is no reason why the orbital mixing should be equal. In fact, if one of the hybridshas more s character it will overlap better with the 1s orbital of hydrogen. We will ignorethis refinement because it does not change the argument.


The bonds formed by the sp-sp overlap as well as the sp-s overlap arecalled -bonds and a single line in our usual structural formulas denotesthem.

Three bonding MO's ( 3. 5 and 7) formed by the new overlaps9.Three antibonding orbitals ( 4. 6 and 8) are also formed. They have anodal plane (where the probability of finding electrons is zero) much likethe one we showed for the hydrogen MOs. We can put these MOs on ourenergy diagram although without calculations and experimental(spectroscopic) data, we will have to guess at the relative positions. It takesmore energy to break a C-H bond (99 kcal/mol) than a C-C (83 kcal/mol). Since the C-H bond is more stable than the C-C bond, we shouldplace the C-H bonding orbitals ( 3 and 5) below the C-C bondingorbital ( 7):


A(2pX,Y,Z) B(2pX,Y,Z)

A (sp hybrids) B (sp hybrids)

Ene

rgy

1s (H) 1s (H)

1s (C)1s (C)

2s2s

FirstQuantumLevel

SecondQuantumLevel

4 6

8

7

3 5

2

1

Fig. 1.14

Once we have constructed such a diagram, we can distribute theelectrons in pairs with opposite spins into the molecular orbitals startingwith the lowest energy orbital:

9It is customary to number the MOs according to increasing energy. For this discussion,we will number the bonding/ antibonding orbital pair in sequence.



A(2p)

Ene

rgy

FirstQuantumLevel

SecondQuantumLevel

1

2

3 5

8

4 6

B(2p)

(C-C)

(C-H)(C-H)

7

kernal electrons

Fig. 1.15

1.8.2. Acetylene Uses p Orbitals to Form -BondsWe have placed for ten of the fourteen electrons in acetylene in

appropriate orbitals. Each carbon still has two unused atomic 2p orbitalsthat can overlap to form four new molecular orbitals (two bonding andtwo antibonding). In constructing these orbitals, we must pay attention tothe symmetry of the atomic orbitals. We recall the 2p orbitals look likethis:

px py pz

Fig. 1.16

The px orbital was used to make the sp hybrids. If we superimpose the2pz orbitals, one from each carbon, on the -bond framework and allowthe lobes to overlap, we get:


C C HH C C HH

Fig. 1.17

This kind of orbital, where the two lobes of the 2p orbitals overlap, iscalled a orbital, and the bond it creates is called a -bond . Two bonds, a

-bond and a -bond, now join the two carbons. We indicate this as C=C,but now we know these bonds are different. We also conclude thatformation of a -bond requires that the carbon atoms move closer togetherso a double bond must be shorter than a single bond; it is, by about 0.2Angstroms. Finally rotation about the C-C -bond requires disruption ofthe p orbital overlap. This takes energy so we conclude a -bond inhibitsrotation about the -bond.

The 2py orbitals are oriented 90 to the 2pz orbital, and they canoverlap as well; we will leave the coordinates in place to aid our perspective:

py py

+ ==

Fig. 1.18

This creates another -bond. If we combine these two orbitals weget:

Fig. 1.19

Cutting through the orbitals from top to bottom and removing thefront half reveals the -bond framework:

Fig. 1.20


Collectively, the orbital formed from the sp hybrid overlap and thetwo orbitals formed from the p orbital overlap establish a triple bond.The complete charge density looks something like this:

C C HH

Fig. 1.21

We can add the two new bonding orbitals ( 9 and 11) and theirantibonding counterparts ( 10and 12) to the energy diagram:


A(2pX,Y,) B(2pX,Y,)

Ene

rgy

2s

FirstQuantumLevel

SecondQuantumLevel

1

2

3 5

7

9 11

10 12

8

4 6

(C-H)(C-H)

(C-C)

Fig. 1.22

We can now distribute the remaining four electrons into the twobonding orbitals:


Molecular Orbitals

Ene

rgy

FirstQuantumLevel

SecondQuantumLevel

(C-C)

*(C-C)

(C-H) (C-H)

(C-C)

*(C-C)

*(C-H)(*C-H)

*(C-C)*(C-C)

(C-C) (C-C) bonding

antibonding

1

64

8

12

11

10

9

7

53

2

Fig. 1.23

Both the bonding and antibonding orbitals at the first quantum levelare full. These form the kernel electrons, and there is no net bonding fromthem. All of the bonding orbitals at the second quantum level are nowoccupied. 3,5,7 form the H-C-C-H -bond framework. 9,11 form two bonds. The antibonding orbitals at the second quantum level are empty.

1.8.3. Some Important ConclusionsWe can draw some important conclusions from this diagram10 and the

principles we used to construct it. Even if calculations require us to modifythe diagram, the conclusions will not change.

1. Since the p orbitals that form the -bond framework lie in a line,acetylene must be a linear molecule.

10 These energy diagrams are based on guestimates not on actual calculations. They arereasonable, and they are useful in our discussion, but they should not be used or quoted asan accurate representation of acetylene.


2. Because of the overlap of p-orbitals that form the -bonds,significant energy must be used to rotate the molecule. We conclude thetriple bond strongly inhibits rotation about the -bond.

3. Since all the bonding orbitals are full, a bond must be broken beforea new one can be formed.

1.9. How Reactions OccurIn the above diagram, *(C-C) orbitals are the Lowest Unoccupied

Molecular Orbitals (LUMO). A reagent with a free electron pair in itsHighest Occupied Molecular Orbital (HOMO) interacts with one ofthese * orbital first. As it does so, the antibonding properties of thisorbital cause one of the bonds to begin breaking, and two atomicorbitals, one on each carbon, begin to reform. As the interaction continues,rearrangement of the orbitals occurs until the process of addition iscomplete. Let us predict what will happen when water reacts withacetylene. We will use curved arrows to keep track of the electronmovement:

C CH H

H2O

C C

H

H2O

H

C C

HO

H

H

H

acetylene vinyl alcohol

Fig. 1.24

The product is vinyl alcohol. If you have guessed this is not the end ofthe reaction, you are correct; we will take this up little later.

1.10. EthyleneAlthough we will not encounter triple bonds in biochemical reactions,

we have learned some important things about multiple bonds in general byconsidering the orbitals involved in forming acetylene. We can use theseprinciples to draw some conclusions about double bonds, which are quitecommon in biochemical reactions.

We can predict the geometry of double bonds by introducing one newprinciple: Mixing a 2s orbital and two 2p orbitals gives three 2sp2 hybridorbitals. Each of these sp2 orbitals looks like this:

Fig. 1.25


In order to keep the electrons in sp2 orbitals as far apart as possible wecan overlap them so the large lobes (the regions of highest electron density)point toward the apices of an equilateral triangle:

Fig. 1.26

This arrangement is said to have trigonal symmetry, which means sp2orbitals lie in a plane and point toward the apices of an equilateral triangle..

The large lobes of the sp2 hybrid orbitals can combine with orbitals ofother atoms to form -bonds. Therefore, a trigonal carbon should formthree -bonds. The resulting molecule should be "flat." For example the -bond framework of ethylene is:

Fig. 1.27

An atomic 2p orbital remains at each carbon.

H H

HH

Fig. 1.28

These can overlap to form and * MOs. Formation of the p-bondcompletes the structure of ethylene:

H H

HH

Fig. 1.29


We can make a reasonable guess that the energy levels of the orbitalsinvolved in forming ethylene will not be very different from those ofacetylene. However, there will be four bonding orbitals for the C-Hbonds (instead of only two) and there will be only one bonding orbital(instead of two). Furthermore, we must place sixteen electrons instead ofonly fourteen. The following is a first approximation:

(C-H)

(C-C)

(C-C)

(C-C)

(C-C)

(C-H)

ener

gy

First Quantum Level

Second Quantum Level

Molecular Orbital

Fig. 1.30

Because of the -bond, we expect ethylene to be reactive. Since rotationabout the double bond is inhibited by the bond, and since the -bondframework is formed from sp2-sp2 and sp2-s overlaps, we predict that all theatoms attached to a double bond will lie in the same plane with bondangles of 120. Physical measurements (e.g. X-ray diffraction) indicatepredictions are essentially correct.


1.11. Geometrical IsomersThe properties we have just described for ethylene, based on its

molecular orbitals, are common to all carbon-carbon double bonds, andthis allows us to make some generalizations.

Because of restricted rotation and the bond angles, we expect geometricisomers to exist when a double bond has the general structure:

B

A

B

B

cis transB

A

A

B

Fig. 1.31

When identical groups are on the same side of the double bond, theisomer is designated cis; when they are on opposite sides, the isomer isdesignated trans. This nomenclature works as long as there are at least twoidentical groups11.

1.12. TautomerizationThe reactive double bond is sensitive not only to intermolecular

reactions (like the one we considered for addition of water to acetylene),but also two intra-molecular reactions (internal reactions). Let us go backto vinyl alcohol, which was formed by addition of water to acetylene. Vinylalcohol still has a reactive double bond. The oxygen atom has a freeelectron pair (actually two pairs) that can displace the double bond:

CC

O

HH

H H

CC

O

HH

H H

CC

O

HH

H

H

acetaldehyde

Fig. 1.32

Acetaldehyde is the product of this internal addition reaction. Thisreaction is so favorable that you cannot isolate vinyl alcohol when wateradds to acetylene. These compounds differ in the positions of the doublebond and one of the hydrogens. Pairs of this kind are called tautomers.

11 At the appropriate time, we will discuss the nomenclature that is used when all fourgroups are different).


1.13. Esters, Amides, Ketones, and Aldehydes Have -BondsThe bonding arrangement illustrated by ethylene and vinyl alcohol

describes not only carbon-carbon double bonds, but carbon-oxygen doublebonds as well. Esters, amides, aldehydes and ketones all contain a -bondbuilt on a -bond framework formed from sp2 hybrid orbitals:

R' R'R'An amide

R' R' R'

An aldehyde

R'R'R'A ketone

An esterRO

R'

RO

R'

RO

R'

H2NOC

H2NOC

H2NOC

C OH

C OH

C OH

ROC

ROC

ROC

C O C OC O

Fig. 1.33

1.14. ElectronegativityWe know that the electrons in a -bond of this type lie at a higher

energy level than those in the -bond because the overlap of atomic porbitals to form the molecular orbital is not as good as the overlap of thesp2 hybrids to for the -bond. It follows that the electrons are more easilydisplaced from their molecular orbital than the electrons. In other words,the electrons are mobile. In each of these compounds, the electrons arepulled toward the oxygen atom because the nucleus of oxygen contains twomore protons and, therefore, has a higher positive charge than carbon. Thiscreates a partial dipole in which the carbonyl carbon carries a partialpositive charge and the oxygen carries a partial negative charge.

C O+

Fig. 1.34


Because of this accumulation of charge, oxygen is said to be moreelectronegative than carbon. As we go across the periodical table, thenumber of protons in the nucleus increases causing the electrons to bedrawn towards the nucleus. Thus, we can predict electronegativity willincrease C


sp3 hybridFig. 1.35

Each of the large lobes of the sp3 hybrid can combine with the orbitalof some other atom. For example, combining the 1s orbital of four differenthydrogen atoms with the sp3 hybrid orbitals of carbon forms methane.

Fig. 1.36

Of course, the combination of atomic orbitals is not limited tohydrogen. Carbon-carbon, carbon-oxygen and carbon-nitrogen -bondsare formed this way. With sp3 hybridization, there are no p orbitals left overto form -bonds. This arrangement of electrons characterizes saturatedcompounds. We expect them to be much less reactive than compoundscontaining -bonds because a -bond must be broken before a reaction canoccur at a saturated carbon and, in general, -bonds lie at lower energylevels than -bonds.


1.16. PhosphateBefore we leave the subject of bonding, we need to discuss the

phosphate group because it plays a very important role in metabolicreactions. We usually write these compounds like this:

O

P

OH

RO OH

O

P

OH

RO O

O

P

O

RO O+ H+ + H+

Fig. 1.37

This implies that phosphates have a double bond. Let us look at thismore carefully. We will use phosphoric acid for this discussion, but theconclusions apply to all phosphates.

Phosphorus is a second row element. It has five valance electrons,which are located at the third quantum level. This has importantconsequences. Since phosphoric acid has four oxygen atoms attached tophosphorus, we need four bonding orbitals to form the -bondframework.

This suggests that phosphoric acid is hybridized sp3. These orbitalslook just like the ones we considered for carbon except they are largerbecause they lie at a higher energy level.

sp3 at phosphorussp3 at carbon

Fig. 1.38

When we considered carbon, we overlapped these sp3 orbitals with thes orbitals of four hydrogen atoms to give methane. For phosphoric acid, weneed to overlap some sort of orbital from each of the four oxygen atomswith the sp3orbitals of phosphorus to form four bonding MO's and fourantibonding MO's. We are going to use sp3 orbitals for the OH oxygen, butthere are other ways of representing this group.

The O-H group has seven electrons (six from oxygen and one fromhydrogen). One of these sp3 orbitals of oxygen overlaps an s orbital of H,


forming a MO that carries one electron from oxygen and one fromhydrogen. Two other sp3 orbitals carry non-bonding, free electron pairs(lone pairs) from oxygen. The fourth sp3 orbital of oxygen carries a singleelectron, which can be used to bond with phosphorus. This accounts forthe seven electrons in the O-H group:

.O

H

Fig. 1.39

Attachment of the three such OH groups to phosphorus can berepresented as follows:

O

H

OH

O

H

P

Fig. 1.40

This accounts for three of phosphorus's five valance electrons and all ofthe OH electrons. We are left with an sp3 orbital at phosphorus and twovalance electrons that can be used to form a -bond with the fourthoxygen. However, it is not so easy to decide how this oxygen should behybridized. Let us look at two possibilities.

We have to account for eight electrons (two from phosphorus and sixfrom the oxygen). If the oxygen is hybridized sp3, we can form fourbonding and four antibonding orbitals. The eight electrons will fill thebonding orbitals:


HH

H

O

O

O

P

O

= P

HO

O

OHOH

Fig. 1.41

Notice that both the -bond electrons that form the fourth P-O bondcome from phosphorus. This type of bond is sometimes called a coordinatecovalent or dative bond to distinguish it from the usual -bond where eachatom contributes one electron. However, this is just a bookkeeping devise.Once the orbital is filled, we cannot tell where the electrons came from sothere is no real difference between a covalent and a coordinate covalentbond.

Most importantly, this arrangement predicts that phosphoric acid willhave a permanent dipole. It does, but the value is not large enough toexplain full, formal charges of this type.

Alternatively, the oxygen might be hybridized sp2. We still have toaccommodate eight electrons, but the arrangement will be quite different.Combining the sp3 from phosphorus and the sp2 from oxygen creates abonding and an antibonding MO. The bonding orbital can be filled withone electron from phosphorus and one from oxygen giving the following

-bond framework:

HH

H

O

O

O

P

O

Fig. 1.42

This completes the -bond framework and accounts for thirty of thethirty-two valence electrons. We have two electrons left, one from oxygenand one from phosphorus. These electrons can be used to make a -bond.There is an atomic p orbital on oxygen, but we have used up the 3s and 3porbitals on phosphorus. Unlike second row elements, where the d orbitalslie at much higher energy levels than do the 2p orbitals, the 3d orbitals of


third row elements such as phosphorus lie at low enough energy levels to beused in the bonding process. There are five 3d orbitals. They look like this:

x2-y2d

z2d

xyd

d

yzd

dz2-y2z2-x2d

=+x

z

y y

z

x

x

zy

yz

x

y

x

x

z

y

xz

Fig. 1.43

A d-orbital of phosphorus and the remaining p-orbital of oxygen canbe used to form a -bond. Since we know what the -bond orbitals looklike, let us represent them with simple lines for the time being andconcentrate on the -bond. Let us rotate the molecule to make it easier toshow the p-d overlap. The -bond looks like this:

O P

OH

OH

OH= PO

OH

OH

OH

Fig. 1.44

The p-d overlap is superimposed on the sp3 framework of phosphorus.Breaking the -bond has very little (if any) effect on the -bonds. This is in


contrast to the carbonyl where breaking the -bond allows the molecule torehybridize from sp2 to sp3, which is energetically favorable since sp3

orbitals lie at a lower energy level than sp2 orbitals.How do we decide which of these representations is correct? Pure

phosphoric acid can be crystallized. X-ray crystallography gives thefollowing data:

1.52

113

105 107

111

O O

O

1.58

1.57

1.57

O

H

H

H

P

Fig. 1.45

The bond angles do not fit either sp3 (109.5) or sp2 (120) exactly.Still, the molecule is clearly tetrahedral. The discrepancies in bond anglesprobably reflect repulsion between the lone pairs. One of the oxygen atomshas three such pairs. Repulsion of the lone pairs on the other oxygens,which have only two such pairs, will increase one set of angles and decreasethe opposite set. The data shows one of the lengths are shorter than theother three. This is consistent with a double bond. However, theshortening is much less (about 4%) than is typical of a C=O bond (about14%). Finally, the 2p-3d electrons are even more mobile than the 2p-2p electrons in a carbonyl group. From electronegativity, we predict thatphosphoric acid will have a significant dipole.

It appears that the true structure lies somewhere between the two wehave considered, and other possibilities exist. Clearly, the phosphate"double bond" is not the same as a carbonyl double bond. A carbonyl grouptends to react by addition. Phosphates tend to react by substitution. Thisshould be kept in mind when considering phosphate chemistry.

1.17. Rules to RememberThe fundamental principles of bonding we have discussed here leads to

a series of rules that are crucial to electron pushing:

Rule 1.1: A covalent bond must be broken before a new one can be made.Rule 1.2: Double bonds are more reactive than single bonds.


Rule 1.3: The atoms attached to double bonds lie in a plane with bondangles of 120.

Rule 1.4: Rotation about a double bond is restricted.Rule 1.5: When each carbon in a double bond carries two different

groups, A and B, two geometrical isomers are possible.Rule 1.6: -Electrons move toward the most electronegative atom.Rule 1.7: Electronegativity increases: H=P


We should not think of an electron as a particle traveling in an orbitlike the earth traveling around the sun. That gets us into all sorts of trouble.For example, there is a point in a p-orbital where the probability of findingan electron is zero. How, you might ask, does the electron get from one sideof the p orbital to the other if it has to pass through a nodal plane?Remember, an electron is not a simple particle. It also has wave-likeproperties. It is wrong to think of it as discrete particle moving around inan orbital. We should think of an electron as a negative cloud with aspecific shape that depends on its energy.

The principles we have developed apply to all compounds. Our mentalpicture of bonding will enable us to understand why the three-dimensionalshapes of molecules are different and why different kinds of reactionsoccur.


2SOME STEREOCHEMISTRY

The purpose of this chapter is to make sure we have sufficientinformation to deal with the three dimensional aspects of organic chemistrythat are important in enzymatic reactions. We are going to discuss theconcepts of chirality and prochirality and the conventions that have beendeveloped for dealing with these three-dimensional problems.

2.1. Dealing With Asymmetric Atoms in Three DimensionsLet us start with a carbon that has four different groups attached to it.

From our discussion in Chapter 1, we know that such a carbon is hybridizedsp3 with the four bonds pointing towards the apices of a tetrahedron. We canrepresent the groups attached to an asymmetric carbon as follows:

b

a

c

d

Fig. 2.1

Now tilt the tetrahedron backwards so the a and d groups point towardsus and the b and c groups point away:

b

c

b

a

cda

d=

Fig. 2.2

Next add back the central carbon while maintaining the 3D arrangementof the groups attached to it using solid "wedge bonds" to indicate groups that

Chapter 2. Some Stereochemistry 35

point towards us and "broken wedge bonds" to indicate groups that pointaway from us:

b

C

c

a

b

c

=a d d

Fig. 2.3

Finally, flatten the 3D structure into the plane of the paper:

Ca

b

d a

b

c

d=

c

Fig. 2.4

Let us go back to the tetrahedron. Arbitrarily, we have drawn the a andd groups at the front, which will point towards us in the 3D structure, butwe could have chosen any two of the four groups as long and we represent thestructure in three dimensions. However, we get into trouble when we go to twodimensions. Let us illustrate this with glyceraldehyde.

Suppose the groups are arranged as follows:

CH2OH

HHO

CHO

Fig. 2.5

We can represent this as:

= H= HHO=

CHOCHO

HOHO

HHCH2OH

CH2OH

C

CHO

HO

CH2OH

CHO

CH2OH

Fig. 2.6


Now suppose we rotate the tetrahedron counterclockwise about a verticalaxis passing through the asymmetric carbon and the carbonyl carbon:

rotateCHO

HO H

CH2OH

H

OH

CHO

CH2OH

CH==

OH

CHO

CH2OH

OHCH2OH

CHO

OH

H

CHO

CH2OHH

Fig. 2.7

There is still no confusion as long as we keep thinking in three dimensions.These structures are simply different orientations of a single compound. Butlook what happens when we flatten the 3D structure into two dimensions:

CH =CH2OH

CHO

OH

H

CHO

CH2OH

OH

Fig. 2.8

Now compare the two-dimensional structures we have generated:

H

CHO

CH2OH

OHCH2OH

HO

CHO

H

Fig. 2.9

They do not look alike even though they were derived from a singlecompound. Of course, we could go back to the 3D structures, but only if weobey the same rules we used in arriving at the two. Furthermore, if we weregiven one of these 2D structures and asked to draw the 3D structure, we couldnot do it unless we knew the procedure that was used to arrive at the 2Dstructure. Clearly, we need some rules to follow if we are to avoid confusion.


2.2. The Fisher ConventionTwo conventions have been devised for dealing with 3D structures of this

kind, i.e. carbon with four different groups attached to it. Emil Fischer, afamous German organic chemist, developed a convention to deal withcarbohydrates. This very useful convention applies not only to carbohydrates,but to amino acids and fatty acids.

Rule 2.1: Convert the 2D structure to 3D by pulling the groups attached tothe horizontal bonds towards you and pushing the groups attachedto the vertical bonds away from you. Use "wedge bonds" to indicatethe 3D arrangement.

Rule 2.2: If necessary, switch the OH group with one of the other groups sothat the OH group is attached to a "horizontal bond" and pointstowards you. Now switch another pair so the H is attached to theother "horizontal bond" and points towards you. The aldehydeshould now be on top and the CH2OH on the bottom with bothpointing away from you

Rule2.3. If the hydroxyl group attached to the penultimate (next to last)carbon is on the right, the compound belongs to the D series.

Let us apply these rules to the structures:

CH2OH

HO

CHO

H H

CHO

CH2OH

OH

original rotomer

Fig. 2.10

We already know these structures were derived from a single compoundby counterclockwise rotation about the vertical axis. Converting the firststructure from 2D to 3D by Rule 2.1 gives:

C H=

CH2OH

HO

CHO

CH2OH

HO

CHO

H

Fig. 2.11


The groups in the 3D structure are oriented properly according to Rule2.2. Now consider the second 2D structure:

CH CH2OH

CHO

OH

H

CHO

CH2OH

OH

Fig. 2.12

The groups in the 3D structure are not oriented properly so we must carryout two switches (the first switch inverts the configuration; the second switchrestores the original configuration). The CHO is in the correct positionaccording to Rule 2.2 so we do not want to move it. If we switch the OH andthe CH2OH we get:

CH OHC

OH

CHO

CH2OH

CH2OH

CHO

H

Fig. 2.13

The groups are now oriented according to Rule 2.2, but we have invertedthe structures so it no longer represents the original structure. Furthermore,we cannot carry out a second switch, which would restore the originalconfiguration, without violating Rule 2.2. Since switching the OH and theCH2OH did not work, let us go back and switch the H and OH instead:

C

H

HO CH2OHC

OH

CHO

CH2OH

CHO

H

Fig. 2.14


Now we can switch the H and the CH2OH to give:

C HC

CHO

HO

CH2OH

CH2OH

CHO

HO

H

Fig. 2.15

In two dimensions, the original structure and its rotomer look different:

=

CH2OH

HO

CHO

H H

CHO

CH2OH

OH

original rotomer

Fig. 2.16

However, by applying the rules of the Fisher convention, we have shownthat these two structures represent the same compound.

2.3. Enantiomers

Rule 2.4: The number of isomers is given by 2n where n=the number ofasymmetric carbons.

Rule 2.5: Enantiomers are mirror images.

Glyceraldehyde has one asymmetric carbon so there must be two isomers.We can write the 2D structures as:

HO H OHH

CHO

CH2OH

CHO

CH2OH

Fig. 2.17

Imagine that the structures are lying in your hands with the OH groupspointing towards your thumbs, the H towards your little fingers, the aldehyde


towards your remaining fingers and the carbinol towards your wrists. Nowrotate your hands so the palms face each other:

CHO

CH2

HHO

OH

CHO

CH2

OHH

OH

CHO

CH2

H

OH

OH

CHO

CH2

H

OH

OH

L-glyceraldehyde D-glyceraldehyde

Fig. 2.18

Like groups face each other. Isomers of this kind are called enantiomers.The 3D structures look like this1:

L-glyceraldehyde D-glyceraldehyde

C HHO

CH2OH

CHO

C OHH

CH2OH

CHO

Fig. 2.19

Let us reorient these structures while maintaining the 3D relationship ofthe groups2:

1 We used Rule 2.3 assign the structures to the proper series.2 Unless you are used to thinking in three dimensions, the best way to carry out this kind ofmanipulation is by inspection of molecular models.


H

OH

CH2OH

HO

OHC

H

HOCH2CHO

Fig. 2.20

Since mirror images are related in the same way as our right and lefthands, the central carbon is said to be chiral (Greek for hand)3. Thus,enantiomers have a chiral center.

All the properties of enantiomers are identical except for the effect onplane polarized light. When Fischer measured the specific rotation4 of theseglyceraldehyde isomers, he found that one of them rotated polarized light 8.7to the right; the other rotated the light 8.7 to the left. This is characteristicof enantiomers. They rotate plane polarized light to an equal degree, but inopposite directions. Compounds that do this are said to be optically active.Unfortunately, the sign of rotation does not tell us anything about theabsolute configuration of the molecule so Fischer had to make an arbitrarychoice. He assigned the configurations as follows:

CH OH CHO H

D

CH2OH

CHO CHO

CH2OH

- 8.7 (c =2 in H2O)[ ]25

D-glyderaldehyde L-glyceraldehyde

+ 8.7 (c =2 in H2O)= D[ ]

25=

Fig. 2.21

2.4. D and L Configurations are Based on D-Glyceraldehyde

Rule 2.6: The notations D or L refer to absolute configurations and are basedon D-glyceraldehyde.

3Chirality refers to mirror images. Mirror images can occur without an asymmetric carbon.4Specific rotation is measured under defined conditions. The light source (often the D-lineof a sodium lamp), the concentration (usually in grams/100 ml), the solvent, and the lengthof the light path through the solution (expressed in decimeters) must be specified. We neednot worry about these experimental details here, but the specific rotation measured at 25using the D-line of sodium as a light source is given as [ ]D

25.


It is important to understand that the original designations D and L werearbitrary. The fact that D-glyceraldehyde is dextrorotatory was fortuitous. Drefers to the absolute configuration. The sign of rotation (dextrorotatory = d= +) has no necessary correlation to D or L. Fisher made a lucky guess. As itturns out, Rule 2.3 gives the correct absolute configuration for d-glyceraldehyde.

2.5. The R/S ConventionThe Fisher convention is simple and very useful. We will use it frequently.

However, it requires that we relate the compound of interest to D-glyceraldehyde, and sometimes this is not possible. The "R/S" convention wasdeveloped to deal with this problem. It is more complicated than the Fischerconvention, but it is more general. It is widely used so we need to be familiarwith it. Basically, one assigns a priority to the groups attached to a chiralcarbon. The higher the atomic number of the atom attached to the chiralcarbon, the higher the priority. Thus, O>C>H etc.

The R/S convention, like the Fischer convention, is a way of representinga 3D structure unambiguously. From x-ray crystallography, we now that D-glyceraldehyde has the following absolute configuration5:

C OHH

CHO

CH2OH

=

CH2OH

OHH

CHO

Fig. 2.22

2.6.1. The Sequence RuleThe sequence rule is actually a series of rules. The ones we need here

provide a way of ordering the groups attached to a chiral center.

5Determination of absolute configuration of glyceraldehyde by x-ray diffraction is not asstraightforward as it might seem. It took 75 years to solve the problem. For a brief discussion,see Eliel, E.L. (1962). Stereochemistry of Carbon Compounds. McGraw Hill, New York,San Francisco, Toronto, and London. pp.95-97.


Rule 2.7: Each atom attached to the chiral carbon is assigned a priority thatincreases with atomic number (S>P>O>N>C>H). If anambiguity occurs (two atoms have the same priority), then onemoves to the next atom until the ambiguity is resolved.

Rule 2.8: The group with highest priority is designated a. The group withthe lowest priority is designated d.

Rule 2.9:After assigning the priority, the groups (a,b,c and d) arerearranged, if necessary, so that the group with the lowest priority(d) is at the bottom of the 2D structure. If two groups are switchedto put d at the bottom, then the other two groups must also beswitched.

Let us apply these rules to the Fisher projection of D-glyceraldehyde.

H OH

CHO

CH2OH

ad

b

c

Fig. 2.23

There are three types of atoms attached to the chiral carbon: O, C and H.Of these, O has the highest priority and assigned the designation a; H hasthe lowest priority and is assigned d. The aldehyde and carbinol groupsmake C-C bonds with the chiral carbon, and we must decide which has thehigher priority. To do this, one analyzes each group atom by atom until apriority emerges. Since the CH2OH has two hydrogens and the CHO onlyone, the CHO has higher priority. Therefore, the CHO group is assigned thepriority b leaving c for the CH2OH group. This works in this simple case,but there is a more general procedure for dealing with multiple bondsunambiguously.

2.6.2. Assigning Priorities to Atoms With Multiple Bonds

Rule 2.10: To assign priority to a group with a double bond (e.g. a carbonyl),expand the double bond with replica atoms:


replica atom replica atom

C OC O

(O) (C)

H

H

Fig. 2.24

The expanded aldehyde carbon has two oxygen atoms and a hydrogenatom attached to it while a CH2OH has one oxygen and two hydrogens.Clearly, the CHO takes precedence over the CH2OH.

The priorities in glyceraldehyde are OH >CHO > CH2OH >H, and wecan write:

H OH

CHO

CH2OH

ad

b

c

Fig. 2.25

Now we need to apply Rule 2.9. First, rearrange the groups so that thelowest priority is at the bottom and pointing behind the plane of the paper.To do this we must make two changes. The first inverts the configuration; thesecond restores it.

ad

a

b=

b

c d

c

Fig. 2.26

2.6.3 The Chirality RuleAt this point, we need to restore our 3D representation. To do this we

need to apply the Chirality Rule:


Rule 2.11: The 3D structure is viewed with the group of lowest prioritypointing away from the viewer and the other three groupspointing toward the viewer like the spokes on a steering wheel.Starting with the group of highest priority the circular path a tob to c is traced. If the groups read clockwise the compound isdesignated R6. If they read counterclockwise, the compounds isdesignated S

First, as shown in Fig. 2.27, we tilt the structure so the a,b,c groups pointtowards us and the d group points away:

a

d

a

c

d

b

=bc

Fig. 2.27

Now read the direction of the a>b>c priorities:

a

chemical basis of metabolism

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