Biomolecules 9788184882421

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MOHAN P. ARORA M.Sc., M.Phil., Ph D., F E.S.I., F.A.Z., F.A.S.E.A., A.I.C.C E.

K4iJl GJIimalaya GJlublishing GJIouse • Mumbal • Deihl • Bangalore • Hyderabad • Chennal • Emakulam • Nagpur • Pune • Ahmedabad • Lucknow

© No part of this book shall be reproduced, rerpinted or translated for any purpose whatsoever without prior permission of the publisher in WrIting.

ISBN :978-81-84882-42-1 REVISED EDITION : 2009

Published by

: Mrs. Meena Pandey for HIMALAYA PUBLISHING HOUSE, "Ramdoot", Dr. Bhalerao Marg, Girgaon, Mumbai-400 004. Phones: 23860170123863863 Fax: 022-23877178 Email: [email protected] Website: www.himpub.com

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"This page is Intentionally Left Blank"

CONTENTS 1.

ORIGIN OF THE EARTH

1-11

The Nebular Hypothesis, The Planetesimal Hypothesis, The Tidal Hypothesis, Recent Hypothesis, Shape of the Earth, Size of the Earth, Movement of the Earth, The Earth and the Universe, Major Divisions of the Earth, Lithosphere, Hydrosphere, Atmosphere, Geologie Processes, Gradation, Degradation, Aggradations, Volcanism, Diastrophism, Geologic Time, Rate of Sedimentation, Salinity of the Ocean, Law of Intrusion, Geologic Time Scale, Subdivision of Geologic Time, Unconfonnities, Paleogeography.

2.

ORIGIN OF BIOMOLECUlES

12-21

The Biological Fitness of Organic Compounds, The Hierarchy of Molecular Organization of Cells, The Primordial Biomolecules, The Specialization and Differentiation of Biomolecules, The Origin of Biomolecules, The Fitness of Biomolecules, The Dimensions and Shapes of Biomolecules, Biomolecules in Relation to Cell Structure.

3.

ATOMS AND MOLECULES

22-39

Physical and Chemical Properties of Atoms and Molecules, Elements are not Changed in Chemical Reactions, Atoms are the Basic Particles of Elements, Atoms contains Protons, Neutrons, and Electrons, An Atom is Uniquely Identified by its Number of Protons, Protons Plus Neutron Determine Atomic Mass, Electrons, Electrons Detennine the Chemical Behaviour of Atoms, Energy within the Atom, Chemical Bonds: Linking Atoms Together, Covalent Bonds Consist of Shared Pair of Electrons, Orientation of Bonds in Space, Multiple Covalent Bonds, Unequal Sharing of Electrons, Hydrogen Bonds, van der Waals Forces, Ions Form Bonds by Electrical Attraction, Nonpolar Substances have no Attraction for Polar Substances, Significance of Polar Covalent and Hydrogen Bonds in Biological Systems, Eggs by the Dozen. Molecules by the Mole, Molecular Weight and the Mole, pH, Chemical Changes, Bond Energy, Oxidation-Reduction Reactions.

4.

WATER

40-51

Physical Properties and Structure of Water, Hydrogen Bonding, The Kinetics of Hydrogen Bonding, Water is the Principal Solvent in Organisms, Hydrogen-Bonding makes Water Cohesive and Adhesive, Water Clings to Polar Molecules, Water Helps Maintain a Stable Temperature, Hydrophobic Interactions, Effect of Solutes on Water Structure, Solvent Properties of Water, Water Ionizes, Acids are Proton Donors; Bases are Proton Acceptors, pH is a Convenient Measure of Acidity, Buffers Minimize pH Change, An Acid and a Base React to Form a Salt, The Fitness of the Aqueous Environment for Living Organisms.

5.

CHEMICAL BONDS

52-64

The Atom, Atomic Structure, Chemical Bonds, lonic Bonds, Covalent Bonds, How Covalent Bonds are Formed, Polar Covalent Bonds, Hydrophobic Interactions, Weaker Interactions, Hydrogen Bonds, Hydrogen Bonding in Biological System, Hydrogen Bonding in Proteins, Hydrogen Bonding in Carbohydrates, Hydrogen Bonding in Deoxyribonucleic Acids, Co-Ordinate Bonds, Peptide Bonds, Disulphide Bonds, Glycosidic Bonds, Intermolecular Forces, Ionic (Coulombic) Forces, Dipole-clipole Forces, Van der Waals Forces and London - Dispersion Interaction.

6.

CELL AS A UNIT OF LIFE

65-82

Viruses, Prokaryotic Cells, Non Photosynthetic Eubacteria, Photosynthetic Bacteria, Eukaryotic Cells, A Typical Cell, Plasma Membrane, Cell Walls, Golgi Bodies, Lysosomes, Mitochondria, Endoplasmic Reticulum, Ribosomes, Peroxisomes, Crystals and Oil Droplets, Plastids, Origin, Vacuoles, Microtubules, Centrioles, Cilia and Ragella, The Nucleus, Cytoplasm.

7.

BIOMEMBRANES

83-} 0 1

The Functions of Membranes, Delineation and Compartmentalization, Localization and Organization of Function, Regulation of Transport, Detection and Transmission of Signals, Cell-to-Cell Communication, Membrane Structure, Overton and Lang,nur: The Importance of Lipids, Gorter and Grendel: The Lipid Bilayer, Davson and Danielli: The Importance of Proteins, Robertson: The Unit Membrane, Singer and Nicolson: The Ruid Mosaic Model, Unwin and Henderson: The Structure of Membrane Proteins, The Fluid Mosaic Model, The Erythrocyte Membrane, Junctions are Specialized Contacts between Cells, Desmosomes are Points of Attachment between Some Animal Cells, Tight Junctions Seal Off Intercellular Space between Some Animal Cells, Gap Junctions Permit Transfer of Small Molecules and Ions between Animal Cells, Plasmodesmata Allow Movement of Certain Molecules and Ions between Plant Cells, How a Cell's Plasma Membrane Regulates Interactions with Its Environment, The Passage of Water Into and Out of Cells, Maintaining Osmotic Balance, Bulk Passage Into and Out of the Cell, Selective Transport of Substances Across Membranes, The Importance of the Plasma Membrane.

8.

FLOW OF ENERGY IN A CELL

102-123

Synthetic Work, Mechanical Work, Concentration Work, Electrical Work, Heat and Work, Bioluminescent Work, Using Energy : Phototrophs and Chemotrophs, The Row of Energy in the Biosphere, Free Energy, Enzymes: Biological Catalysts, Mechanism of ATP Synthesis, How does Energy Row in Chemical Reactions? Exergonic Reactions Release Energy, Endergonic Reactions Require an Input of Energy, Coupled Reacton Link Exergonic with Endergonic Reactions, Chemical Reactions are Reversible, How do Cells Control their Metabolic Reactions? At Body Temperatures, Spontaneous Reactions Proceed too Slowly to Sustain Life, Cells Regulate the Amount and Activity of their Enzymes, thus Precisely Regulating their Metabolic Reactions, How is Cellular Energy Carried between Coupled Reactions? ATP is the Principal Energy Carrier in Cells, Electron Carrier!! also Transport Energy within Cells, The Energy Currency of the Cell, The ATP Molecule has Three Main Parts, ATP can Denote Energy through the Transfer of a Phosphate Group, ATP Links Exergonic and Endergonic Reactions, The Cell Maintains a Very High Ratio of ATP to ADP, ATP cannot be Stockpiled, Cells Transfer Energy by Redox Reactions, Most Electron Carriers Carry Hydrogen Atoms, Electron Carriers Transfer Energy, Enzymes are Chemical Regulators, Glycolysis, Lactic Acid Fermentation, Alcoholic Fermentation, The Citric Acid Cycle, The Respiratory Chain, Coupling Electron Transport to the Synthesis of ATP, Mitochondria and Chloroplasts are Miniature Batteries, Catabolism of Proteins and Fats, Cellular Respiration of Protein, Cellular Respiration of Fat, Regulation of Cellular Respiration.

9.

CARBOHYDRATES

124-134

Occurrence, Nomenclature, Isomerism, Classification, Monosaccharides, Trioses, Pentoses, Hexoses, Properties of Monosaccharides, Physical Properties, Chemical Properties, Sugar Derivatives, Sugar

Acids, Sugar Phosphates, Deoxy Sugars, Amino Sugars, Oligosaccharides, Disaccharides, Polysaccharides. Starch. Glycogen, Cellulose. Agar, Pectins, Xylan, Inulin, Dextrins, Dextran, Cellulose, Chitin. Hetropolysaccharides. Glycoproteins. Mucopolysaccharides, Functions of Carbohydrates. Storage Substances of Potential Energy. Structural Component, Regulation of Fat Metabolism. Protein-Sparing Function, Role is Gastrointestinal Function.

10.

UPiDS

135-145

Occurrence of Lipids, Simple Lipids, Neutral Fats and Oils, Nomenclature, Physical Properties of the Fatty Acids, Chemical Properties of the Fatty Acids, Reactions of the Unsaturated Acids, Essential Fatty Acids, Waxes, Beeswax, Spermaceti, Sperm Oil, Camauba Wax, Lanolin or Wool Wax, PhYSiological Importance of Waxes, Compound Lipids, Derived Lipids, Caritenoids (Lipochromes), Functions of Lipids.

11.

AMINO ACIDS

146-159

Digestion of Proteins. Activation of PepSinogen in the Stomach, Activation of Proteolytic Enzymes in the Intestines. Activation of Enzymes versus Enzyme Synthesis, Absorption of Amino Acids, Classification. Abbreviations. Neutral Aliphatic Amino Acids. Aromatic Amino Acids. Basic Amino Acids, Dicarboxylic Amino Acids. Imino Acids. Unusual Amino Acids, Ionization of Amino Acids, Stereoisomers, 20 Standard Amino Acids, Ionization, Reactions, Reactions of the Carboxyl Group, Reactions of Amino Group, Reactions of the R Group, Properties of Amino Acids and Proteins, Ionic Properties.

12.

PROTEINS

160-172

Essential Amino Acids, Non-Essential Amino Acids, Structure of Proteins, Primary Structure, Secondary Structure, Tertiary Structure, Quaternary Structure, Classification of Proteins, Simple Proteins. Conjugated Proteins, Derived Proteins, Properties of Proteins. Colour and Taste, Molecular Weight, Colloidal Nature. Denaturation. Amphoteric Nature. Solubility. Hydrolysis, Reaction Involving COOH Group, Reaction Involving NH2 Group, Reaction Involving both COOH and NH2 Groups, Reactions Involving-R Group or Side Chain, Biological Importance of Proteins.

13.

ENZVMES

173-190

Distribution of Enzymes, Intracellular Localization of Enzymes, Extracellular Localization of Enzymes, Nomenclature. Extraction and Purification of Enzymes. Units of Enzyme Activity, Chemical Nature of Enzymes. Simple Enzymes. Conjugated Enzymes, Classification of Enzymes. Oxido-reductase, Transferases, Hydrolases, Endopeptidases, General Properties of Enzymes, Unaltered State, Quantity, Proteinaeous, Rate of Reaction, Reversibility. Denaturation, Activation, Enzyme Specificity, Factors Affecting the Enzyme Activity, Contact between the Enzyme and Substrate, Temperature, pH, Enzyme Substrate Concentration, Concentration of the Products, Time, Oxidation State of the Enzyme, Radiation, Activators, Inhibition, Mechanism of Enzyme Catalysis, Enzyme Catalyzed Reaction, The Enzyme Substrate Complex, The Michalies-Menten Constant, Turnover Number of Molecular Activity of an Enzyme, Co-factors in Enzyme Action, Ion as Co-factors, Coenzymes and Vitamins. Enzymatic Pathway. Structure of Enzymes. Some Features of Active Sites, Inhibition of Enzymes. Reversible Inhibition. Competitive Inhibition. Non Competitive Inhibition. Irreversible Inhibition. Allosteric inhibitions. Lock and Key Theory. The Induced Fit Hypothesis. Enzyme Deficiencies, Enzymes and Human Diseases.

14.

NUCLEIC ACID (DNA)

191-212

Location of Nucleic Acid, Chemical Background of Nucleic Acids, Bases, Pentose Sugar, PhosphoriC Acid, Deoxyribonucleic Acid (DNA), DNA Contents, Structure of DNA, Molecular Weight of DNA, Molecular Structure of DNA. Nitrogenous Bases. Chargaffs Rules, Polarity of DNA, Unusual Bases in DNA, Sugar, PhosphoriC Acid. Nucleosides. Nucleotides. The Primary Structure of ~NA, The Secondary Structure of pNA. The Tertiary Structure of DNA, Double

Stranded Linear DNA Double Stranded Cyclic DNA, The Structure of the DNA in Eukaryotic Chromosome. Forms of DNA Z-DNA or Left Handed DNA, B-DNA or Right Handed DNA, A-DNA C-DNA D-DNA E-DNA, P-DNA Repetitive DNA or Satellite DNA, Structural Variation in DNA Double Stranded DNA Single stranded DNA Circular DNA, Non-Chromosomal DNA, Mitochondria. Yolk-and Chloroplast DNA, Centriolar DNA, Nucleolar DNA, Catalytic Function, Watson & Crick's Model for Replication of DNA (Autocatalytic nature), Delbruck suggested theoretically three modes of replication of DNA, Replication of DNA in Bacteria or Prokaryote Cells, Enzymes taking part in Replication, Biological Significance of DNA, The Meselson and Stahl Experiment, Cairn's Autoradiography Experiment, Taylor's Experiment on Vicia faba Root Tips, Identification of the Genetic Material, DNA as the Genetic Material.

15.

RIBONUCLEIC ACID (RNA)

213-227

Structure of RNA, Primary Structure of RNA, Secondary and Tertiary Structures of RNA, Types of Ribonucleic Acids, Transfer Ribonucleic acid (t-RNA), Messenger Ribonucleic Acid (mRNA), Ribosomal Ribonucleic Acid (r-RNA), Viral RNA (v-RNA), Mitochondrial RNAs, Structural variations of RNA, Replication of Ribonucleic Acid, Biosynthesis of Viral RNA, Biosynthesis of m-Rl\lA, The notion of a Messenger, Biosynthesis of Transfer RNA, t-RNA Synthesis, Virus Nucleic Acids, Viroids, Prions.

16.

BIOSYNTHESIS OF PROTEIN

228--249

Flow of Information, Central Dogma Reverse. Restatement of Central Dogma, Open Reading Frames (ORFs), Transcription. Core Enzyme. Sigma Factor (s) having molecular wt. 90,000, Intron and Exons. Post TranSCriptional Control in Eukaryotes, Processing of the Primary Transcript, Transport of the Processed Transcript Out of the Nucleus, Selecting Which mRNAs are Translated, Selectively Degrading mRNA Transcripts, Post Transcriptional Control in Prokaryotes, Translation, Amino Acids have no Specific Affinity for RNA, Amino Acids are Aligned on RNA Templates by Means of Adaptors, Specific Enzymes Recognize Specific Amino Acids, The Adaptor Molecules are themselves RNA Molecules, Charging tRNA, Activation of Amino Acids, lsoacceptors, Peptide Bond Formation Occurs on Ribosomes, tRNA Binding Sites, Heterogeneity of Messenger RNA, Attachment of tRNA to mRNA, Polyribosomes, Formation of Polypeptide Chain, Initiation of Protein Synthesis, The Direction of mRNA Reading is 5'~3', Elongation of Polypeptide, Chain Elongation Requires GTP, Movement of mRNA Across the Ribosomal Surface, Termination of Polypeptide Chain, Gene Mutation, The Role of ER, Protein Folding and Processing, Chaperones and Protein Folding, Protein Cleavage, Glycosylation.

17.

BIOSYNTHESIS OF PuRINE AND PYRIMIDINE

250-257

Folate in Nutrition. Utilization of Folic Acid, Structural Relationships between Folate and its Analogues, Methyl-trap Hypothesis. Purine Nucleotide Synthesis, Rate-limiting Reaction in Purine Synthesis, Ribose-phosphate Intermediates, IMP: The precursor of AMP and GMP, Pyrimidine Nucleotide Synthesis, Rate-limiting Steps in Formation of Carbamoyl Phosphate, Clinical Comment, Orotic Aciduria, Salvage Synthesis, Deoxyribonucleotide Synthesis, Thymidylate Synthesis, Analogues of Pyrimidines and Purines, Pyrimidine Analogues, Purine Analogues, Conversion of Analogues to Analogue Nucleotides, Biosynthesis of Nucleotide-containing Coenzymes, Coenzyme A, FAD and FMN, NAD, NADP, TPP and Pyridoxal Phosphate, Biotin and Lipoic Acid.

18.

AMINO ACID METABOLISM

258-274

Fate of the Nitrogen Atoms, Removal of Nitrogen Atoms, Amino Group Transfer, Formation of Ammonium lon, Transport of NH3 to the Liver, Alanine Cycle, Glutamine Cycle, Detoxification in the Brain, Detoxification in the Liver, Urea Synthesis, Synthesis of Carbamoyl Phosphate, Formation of Arginine via Citrulline, Complete Cycle, Compartmentation of Urea Synthesis, Disposal of Excess Alanine, Disposal of Excess Glutamate, Amino Acids as used for Production of Nonproteinic Nitrogenous Compounds. Biosynthesis and Breakdown of Creatine, Production

and Breakdown of Mediators, Production of Histamine, Production of Serotonin, Production of yAminobutyric Acid (GABA), Production of Catecholamines, Production of Taurine, Production of Certain Vitamins and Coenzymes from Amino Acids, Amino Acid Metabolish Control in the Organism, Amino Acid as Medicinal Preparations, Protein Hydrolysate Preparations (Preparations for Parenteral Nutrition), Preparations from Individual Amino Acids, Inborn Errors Concerning the Un~a Cycle Enzymes, Liver Disease and Urea Cycle, Reye's Syndrome.

19.

METABOUSM OF INDIVIDUAL AMINO ACIDS

275-288

Glycine, Inborns Errors of Glycine Metabolism, Serine, Cysteine, Inborn Errors of Sulfur Containing Amino Acids, Inborn Errors of Glutathione Metabolism, Aspartic Acid and Asparagine, Glutamic Acid and Related Amino Acids, Special Role of Glutamate in the Brain Tissue, Alanine, AlanineGlucose Cycle, Proline and Hydroxyproline, Arginine, Lysine, Threonine, Tryptophan, Synthesis of NAD+, Synthesis of Serotonin, Synthesis of Melatonin, Hartnup Disease, Histidine, Valine, Leucine, Isoleucine, Phenylalanine and Tyrosine, Main Catabolic Pathway of Phenylalanine(fyrosine, Formation of Catecholamines, Thyroid Hormones and Melanin from Tyrosine, Inborn Errors of Metabolism, Methionine, Methyla~()n.

20.

METABOUSM OF CARBOHYDRATES

289-314

Energy Requirements, Source of Energy, General Metabolism, Basal Metabolic Rate, Intermediate Metabolism, Carbohydrate Metabolism, Cori Cycle, Glycogenesis, Glycogenolysis, Reactions of Glycogenolysis, Control of Glycogen Metabolism, Gluconeogenesis, Mechanism for Gluconeogenesis, Noncarbohydrate Sources for Ghlconeogenesis, Energy Liberation, Glycolysis, Phosphorylation of Glucose, Glucose-6-phosphate Conversion to Fructose-6-phosphate, Phosphorylation of Fructose-6-phosphate to Fructose- 1, 6- Diphosphate, Cleavage of Fructose1, 6-diphosphate to Triose Phosphate, Oxidation of Glyceraldehyde-3-phosphate and Formation of ATP, Conversion of 3-phosphoglyceric Acid to Phosphopyruvic Acid, Phosphopyruvic Acid to Pyruvic Acid & ATP Formation, Pyruvic Acid to Lactic Acid, Glycolysis in Other Animals, Glycerol Phosphate in Insects, Kreb's Cycle, Condensation, Dehydration, Hydration I, OxidationI-or Dehydrogenation I, Decarboxylation I, Oxidation II or Dehydrogenation II and Decarboxylation II (Oxidative decarobxylase), Phosphorylation, Oxidation III or Dehydrogenation III, Hydration II, Oxidation IV or Dehydrogenation, Electron Transport Chain, NADH-dehydrogenase or NADHCOQ-Reductase Complex, Succinate-Dehydrogenase or Succinate-CoQ-Reducates Complex, Coenzyme-Q or Ubiquinone, CoQ-cytchrome c-reductase Complex, Cytochrome c, Cytochrome c-oxidase Co.mplex, Electron Transport Shuttles, Malate-asparate Shuttle, Glycerol-phosphate Shuttle, Oxidative Phosphorylation Mechanism, Chemical Coupling Hypothesis, Chemiosmotic Coupling Hypothesis, Conformation Coupling Hypothesis, Regulation of Carbohydrate Metabolism.

21.

METABOUSM OF LIPID

315-339

Acyl CoA Synthase,' Acylcamitine Formation, ~-Oxidation Sequence, Other Types of Fatty Acid Oxidation, Energy Yield of Fatty acid ~-oxidation, Regulation of Substrate Utilization, Metabolism of Structural Lipids, Phospholipids, Sphingolipids, Metabolism of Triacylglycerols, Cholesterol Metabolish, Fatty Acid Synthesis, Source of acetyl CoA, Carboxylation of Acetyl CoA, Fatty Acid Synthase, Regulation of De novo BiosyntheSis, Differences between Synthesis and Oxidation, Fatty Acid Chain Elongation, Fatty Acid Desaturation, Ketone Bodies, Ketone Body Synthesis, Ketone Body Oxidation, Clinical Comment.

,22.

CATABOUSM OF NUCLEOTIDES

340-344

Purine Catabolism, Lesch-Nyhan Syndrome, Adenosine Deaminase Deficiency, Purine Nucleoside Phosphorylase Deficiency, Pyrimidine Catabolism, Inhibitors of Purine, Pyrimidine Nucleotide Biosynthesis, .Glutamine Analogs, Folic Acid Analogs, L-asparaginase, Purine Anologs, Pyrimidine Analogs, Hyperuricemia and Gout, Gouty Arthritis.

23.

VITAMINS

345-361

Fat Soluble Vitamins, Vitamin A, Vitamin D, Vitamin E, Vitamin K, Water Soluble Vitamins, Vitamin C, Vitamin B Complex, Thiamine (Vitamin B1), Riboflavin (Vitamin B2), Nicotinic Acid, Pantothenic Acid, Pyridoxine, Biotin, Folic Acid, Cobalamin (Vitamin B I2), Inositol, Choline, PAmino Benzoic Acid.

24.

HORMONES

362-433

Foundations of Endocrinology, Historical Beginings, Methods of Study, Chemical Nature of Hormones, Functional Classification, Developmental Hormones, Hormone Metabolism and Excretion, Regulation of Hormone Secretion, Feed Back Control, How Hormones Work, Principles of Hormone Action, Specificity of Hormone Action, Time for Hormone Action, Dose-ReSponse Effects, Control of Hormone Concentration, Neuroendocrine Reflex Arcs, Feedback.. Hormone that Enter the Cell, Hormone that Do Not Enter Cells, Second Messengers in Action: How Epinephrine Works, Invertebrate Hormones, Properties of Hormones, Functions of Hormones, Endocrine Glands in Vertebrates, The Pituitary Gland, Location, Development, Structure, Cell Types in Adenohypophsis, Hormones of the Pituitary Gland, Hormones of the Adenohypophysis, Secretion of Growth Hormone, Function, Other Functions, Secretion, Prolactin-Releasing Factors, Episodic and Sleep Related Secretion, Other Stimuli, Effects of other Hormones, Effects of Pharmacologic Agents, Secretion, TRH, Somatostatin, Neural Control, Temperature and Stress, Effects of Cortisol and Estrogens, Hormones of the Neurohypophysis, Thyroid Gland, Synthesis of Thyroxine, Antithyroid Drugs, Physiological Effects, Thyroid Gland and Metamorphosis, Regulation of Thyroid Secretion, Hypothroidism, Thyrocalcitonin, Parathyroid Gland, Control of Parathromone, Functions, Hypoparthyroidism, Hyperparathyroidism, Adrenal or Supra Renal Gland, Adrenal Cortex, Hormones of the Adrenal Cortex, Mechanism of Action of Aldosterone, The Control of Aldosterone Secretion, Biological Functions, Control of Gluco-corticoid Secretion, Adrenal Medulla, Biological Functions, Nor-epinephrine, Control of Hormones, Pancreas, . Glucagon, Insulin, Effects of Insulin Lack, Action of Insulin, Glucagon, Biochemistry, Action of Glucagon, Somatostatin, Pancreatic Polypeptide, Regulation of Insulin and Glucagon Secretion, The Pineal Body, Thymus, Gastro-Intestinal Hormones, The Gonads, Testes, Functions of Androgen, Castration, Ovary, Functions of Estrogens, Functions of Progesterone, Placental Hormones, Progesterone, Estrogens, Human chorionic gonadotropin (HCG) , Prostaglandins, Composition, Sources, Actions, Parahormbnes, Angiotensin, Erythropoietin, Bradykinin, Neurosecretions, Inter-relation between the Endocrine System and Nervous System, Disorders and Medical Terminology, Hyperpituitarism, Hypo-pituitarism, Dispituitarism, Adrenal Gland, Gonads, Thyroid Gland.

25.

INBORN ERRORS OF METABOUSM

434-445

Alkaptonuria, Phenylketonuria, Tyrosinosis, Sickle Celled Anaemia, Albini.sm, Tay-Sachs Disease, Histidinemia, Galactosemia, Goucher's Disease (GD), Glycogenosis, Christmas Disease, Huntington's Disease, Glucose-6-Phosphate Dehydrogenase Deficiency, Pentosuria, Essential Fructosuria, Mucopolysaccharidosis -(Hurler's Syndrome), Hunter's Syndrome, Maple-Syrup Urine Dise3$¢, Cystinuria, Hartnup Disease, Niemann-Pick Disease, Fabry's Disease, Refsum's Diseas~, Tyrosinemia, Citrullinuria, Glycinuria, Homocystinuria, Hyperlipemias, Analbuminemia, LaschNyhan Syndrome, Kwashiorkor: Protein Starvation. INDEX

446-466

1 ORIGIN OF THE EARTH Man has been speculating on the origin of the earth for thousands of years. His early speculations took the form of myth-making. Later, the ancient philosophers developed theories based on observation and reasoning. Truly scientific hypotheses had to wait until man had assembled adequate knowledge of the earth and its actual relation to the heavenly bodies. None of the numerous hypotheses regarding the origin of the earth has been generally accepted because each hypothesis fails to account for certain facts which must be explained by any fully acceptable theory. Confronted by confusions and contradictions, as well as by the complex mathematics involved, the average geologist is tempted to consider this matter as one that does not directly concern his subject and to relegate it to the astronomers . Nevertheless, certain elaborate hypotheses of earth origin have played a significant role in the development of geologic theory, and for that reason are worthy of at least a brief consideration. Whether the universe exploded at a single Fig. 1.1. Cloud of dust from which the solar system evolved may have developed this instant from' one mass of intricate pattern of turbulence.

Biomolecules

2

gaseous material, which has been expanding through space evel'since (the "big-bang" hypothesis), or whether it represents alternating transfonnations of energy into mass and mass into energy (the "Steady-state" hypothesis) 4s still subject to dispute. The universe may be several times as old as the solar system perhaps 15-20 billion years - but it seems likely that all the planets originated at one time. Once the earth had formed, th~ oceans and the atmosphere may have developed later, the prevailing belief being that they resulted from the expulsion of volatile matter (planetary outgassing), as by volcanism. THE NEBUlAR HYPOTIlESIS

I

Copernicus, Kepler and Newton followed one another at Intervals of about a century; another century later, based upon their profound discoveries, the first scientific hypothesis of earth origin was presented (in 1755) by the German metaphysician Immanuel Kant. This hypothesis was further developed in 1796 by the French mathematician and astronomer Pierre Simon de Laplace, and was widely accepted during the nineteenth century. The nebular hypothesis of Kant and Laplace is referred to as a star hypothesis because it explains the origin of the planets by the evolution of a single star, our sun, without (a) the interference of any outside body. The hypotheSis postulates that ' a greatly diffused spherical cloud of gas, a nebula, extended outward at latest to the present distance of the outennost planet. ·This cloud rotated slowly; as it cooled and contracted its velocity around the sun increased, just as a whirling dancer spins faster as he draws his (b) arms together. The gaseous mass gradually became a disk around the sun's equator. At critical points during this rotation, rings of fiery gas were assumed to have been thrown off by centrifugal force. Each ring then broke up and gathered into a sphere producing a planet, which began to revolve around the sun in the same path as the former ring. A (c) comparable process, operating in a minor way in the case of several of the planets, accounts for the fonnation of the satellites such as our moon. The earth-planet liquefied as it cooled, then with further cooling acquired a solid crust. The main body of the gas meanwhile condensed (d) further to become the sun. Among the more serious objections to the Fig. 1.2. Stages during the condensation of the solar nebula nebular hypothesis are the following. into the solar planetary system. (a) Fragmentation 1. The planets possess 98 per cent of the of an interstellar cloud. (b) Contraction and • rotational energy of the solar system, flattening of the solar neubla. (c) Condensation of nebular material into meteorites and protoplanetary whereas the sun has about 99.87 per bodies. (d) Solidification of planets. cent of the mass.

I j

I

c; J

Origin of the Earth

3

2. The heavy elements in the earth can originate only at temperatures far higher than those prevailing on the sun . 3. Some of the satellites revolve in a retrograde (backward) direction, and one of them revolves faster than its planet rotates. 4. The mechanism of ring-formation does not correspond to the rotational velocity of a solar nebula as postulated . THE PlANETESIMAL HYPOTHESIS

Finding nothing in outer space that corresponded to the netbula postulated by Kant and Laplace, and impressed by the recurrence of largescale glaciation throughout geologic history, which seemed to contradict the idea of a steadily cooling earth, Thomas C. Chamberlin and Forest R. Moulton , an American geologistastronomer team, presented the Solardisruption, or planetesimal hypothesis in 1895. Geologists favoured this conception of earth origin until about the time of World War II. The planetesimal hypothesis of Chamberlin and Moulton is referred Fig. 1.3. Evolution of the Earth and the planets Mercury. Venus and to as a two-star hypothesis because Mars . it explains the origin of the planets by a near collision between the sun and another star, which disrupted it. The hypothesis supposes that the world began with a sun nearly the same as the present one, ejecting material by internal explosions, such as those seen today in the so-called solar prominences. The close passage of another star developed huge tidal bulges on two opposite sides of the sun, and the explosive force within drove them out in varying directions as great jets, or bolts of gas. The attraction of the passing star drew these bolts into elliptical orbits around the sun. Five large bolts on the tidal bulge facing the sun broke up to form the major planets, and five smaller ones on the opposite side became the minor planets and asteroids, nearer the sun . Some of the hot gas in the smaller bolts is presumed to have condensed rapidly into liquid and then cooled to solid particle~, :called planetesimals and resembling meteorites, revolving in swarms aroLlAd the sun in a pattern like that of the many spiral nebulae seen in the sky. As their paths intersected, the nucleus of each planet was assembled , and these few grew by accretion to their present size, moving in more circular orbits as a result of the encounter. Thus the earth was solid from the start, except for a temporarily molten surface, which could result from heat generated by the rapid infall of planetesimals. The craters on the moon seem to show the effect of similar impacts. Among the more serious objections to the planetesimal hypothesis are the following :

Biomolecules

4

/

1. The concentric density layering of the earth indicates that it must · once have been a molten body, which segregated into unlike zones, as slag does from molten metal. 2. The mechanism of solar eruption is inadequate to explain the tidal bulges. 3 . A very rare and perhaps unique even in . the universe is requirea. 4. The collision of planetesimals would more likely destroy than preserve them. 5. The amount of salt in the oceans, derived mostly from weathering of rock, is not enough for a slowly growing planet to have provided, considering all the time available for the process of erosion and deposition. THE TIDAL HYPOTHESIS

Unconvinced that the earth has always Fig. 1.4. Stage in the early euolution of the earth. Beginning been solid, ' James Jeans a Harold leffries, with the sphere that developed by accretion in the first English astronomer and geophysicist, few million years following the condensation of the respectively, proposed in 1917 a gaseous solar nebula. two-star hypothesis. They theorized that another star closely approach the sun, producing tidal bulges, from which streamed an enormous cigar-shaped filament of solar gases. This incandescent filament was put into revolution around the sun . It then broke into segments which contracted into rotating spheres, the planets. Cooling from gas to liquid one of these planets, Earth, gradually solidified to its present condition. The objections to the tidal hypothesis are like many of those made to the planetesirpal hypothesis, excluding the ones that ~re dependent upon an aggregation of originally . solid material. The 'hypothesis explains most of the peculiarities of planet and satellite motion by complicated devices. RECENT HYPOTHESIS

Fig. 1.5. Stage in the early euolution of the earth. Later; when some tens of millions of year haue passed, a combination of grauitational compression, radioactiue decay and impact heating produces melting and differentiation; heauy: co~;md mantle materials sink inward and light crusta'l materials float outward.

With the development of new mathematical techniques and the discovery of new facts about the universe, the formulation of systems of earth origin is once again . becoming fashionable among cosmogonists. Among those who have proposed new hypotheses are Ber/age (1940), Alfuen (1942), Von Welzsacker (1945), Whipple (1947), ter Haar (1948) and Hoyle (1950). The general trend is apparently towards a modified version

5

Origin of the Earth

of the ancient nebular hypothesis, involving a vastly greater amount of very hot gaseous material flying from the sun by centrifugal force, a small portion of which condensed to form the planets. The dust-cloud hypothesis of Fred L. Whipple suggests for the beginning of the earth a cloud of cosmic dust, similar to the many dark nebulae. This dust was driven together by the force of light on the outside, which, as we know, is able to propel the substance of comets. Moving streams of particles then separated into the planets and their satellites. The nova hypothesis of Fred Hoyle and older binary-star hypothesis (1936) of Henry N. Russel and R.A. Lyttleton involve double stars or binaries, which are pairs revolving about a Fig. 1.6. Stage in the early evolution of the earth. The Archaean era, common center (about half the stars are between 3.7 and 2.2 billion years ago. As atmospheric water probably of this type). Hoyle's view is and gases condense, the oceans appear; the earliest that one member of the pair exploded continents arise and valcanic activity is intense. to produce the material of the planets. Russell and Lyttleton favoured tidal disruption of the companion star by the sun. SHAPE OF THE EARTH

Fig. 1 7. The Proterozoic era follows the Archaean era. It harbors a few scent traces of plant and animal life. It ends some qOOJnillion years ago as the Paleozoic era, with its rich fossil record, begins. During that era's 1.5 billlon·year span crustal cooling and thickening .

The earth is a geoid, a triaxial ellipsoid, nearly spherical but slightly flattened at the poles. The diameter from pole to pole is about 27 miles less than the average distance through the equator, which has two unequal axes. Curious small bl· 1ges make four "corners" at Ireland, off Peru, south of Africa, and near New Guinea. The vertical difference between the highest mountain peak (Everest, over 29,000 feet) and the greatest known depth of ocean (near the Philippine del island of Mindanao, 36,560 feet) is only a little more than 12 miles. The spherical shape of the earth accounts for the following phenomena noted in antiquity: the apparent submergence of ships putting out to sea; the circular shadow of the earth cast upon the moon during lunar eclipses; and the

Biomolecules

6

changing elevation of the north Star in relation to the place from which it is observed. Now the curvature of the earth can be clearly seen in high-altitude photographs. SIZE OF THE EARTH

The earth has a polar diameter of about 7,900 miles and an equatorial diameter of about 7,927 miles. Its circumference, around the equator, is approximately 24,900 miles. The area of the earth's surface is about 197 million square miles, of which about 71 per cent is covered by oceans. The volume of the earth is a little more than 250 billion cubic miles, and its mass has been estimated at about 6,600 quintillion (6,600,000,000,000,000,000,000) tons. MOVEMENT OF THE EARTH

The earth moves in a number of directions at the same time. It rotates on its axis (which extends through the poles), and the duration of one rotation (relative to the sun) has become one of our basic measurements of time, the solar day. The velocity of rotation at the equator is about 1,037 miles per hour; at the poles it is, of course, zero. Besides its rotation, the earth has a light wobbling motion (called precession), which causes its axis to describe, a cone-shaped figure once in about 25,800 years. The earth also revolves about the sun at an average velocity of about 18.5 miles per second, completing one revolution in a little more than 365 1/ 4 days, thus providing us with another basic measurement of time, the solar year. The earth's orbit is elliptical, its distance from the sun varying from about 91.5 to about 94.5 million miles. The axis of the earth is inclined at an angle of 23 .5 degrees, so that each hemisphere is tilted toward the sun during half of the year and ' away from it during the other half, thus producing the seasons. The entire solar system is speeding through space toward the star Vega at a velocity of 12 miles per second. Our galaxy, furthermore, apparently rotates on a axis and has an independent movement away from other galaxies. THE EARTH AND THE UNIVERSE

The universe, as we know it, is composed of innumerable galaxies, or star clusters. The galaxy in which we live, the Milky Way, is a disc-shaped revolving cluster of some 30 billion stars of varying size and brightness. One of these stars, of modest size and brightness and nearer the edge than the center of the galaxy, is our sun.

Table 1.1. Characteristic of the Sun, Moon and Planets Name

Sun Mercury Venus Earth Mars Jupiter Saturn Uranus Neptune Pluto Moon

Diameter (m.les) (earth-l)

864,000 3,000 7,600 7,920 4,200 87,000 71,500 29,400 28,000 7,600 2,160

109 0.38 0 .96 1.00 0.53 10.9 9 .0 3.7 3.5 0.96 0.27

Volume (earth-l)

Density (water-l)

Mass (earth-l)

1,300.000 0 .055 0.88 1.00 0.15 1,325 730 50 43 0 .88 0.02

1.4 3.8 4.9 5.5 4.0 1.3 0.7 1.3 1.6 5.8 3.3

332.000 0.056 0.08 1.00 0.11 318 95 15 17 0.93 0.012

Distance From Sun (million-miles)

36 67 93 142 483 886 1,780 2,790 3,670 93

7

Origin of the Earth

The sun is the center of the solar system, in which the largest solid bodies that revolve about the sun are called planets. The nine planets, in the order of.their proximity to the sun, are Mercury, Venus, Earth, Mars, Jupiter, Saturn, Uranus, Neptune, and Pluto. Table 1.1 gives a simplified comparison of the sun, moon and planets. An unknown number of smaller bodies also revolve about the sun; if they are large enough to be seen (telescopically) and named, they are called asteroids or planetoids. Many very small bodies called meteorites enter the earth's atmosphere daily. (The earth' has been struck by meteorites or asteroids large enough to make huge craters.) The huge but tenuous objects called comets also belong to the solar system. Most of the planets have one or more satellites about the earth at an average distance of a little less than 240 thousand miles. Because the moon makes one rotation about its axis in the same time (about 271/3 days) that it makes one revolution about the earth, it tums the same fac~ to the earth at all tllnes. (But it wobbles sufficiently to enable us to see 59 per cent of its surface at one time or another). The most noteworthy physical effect of the moon upon the earth is the creation of tides. MAJOR DIVISIONS OF THE EARTH

Three zones, corresponding to the three states of matter (solid, liquid, gas), constitute the globe we know as the earth. The solid central zone is the lithosphere. Cradled in the ocean basins and distributed across the surface of the land is the zone of water, the hydrosphere. Surrounding them both is a gaseous envelope, the atmosphere. The boundaries between the three zones are not perfectly sharp; there is some mingling, of air and water, of air and rock, of water and rock. Uthosphere . The nature of the lithosphere-the earth's crust and interior. Here we shall mention some important processes and relief features of the lithosphere.

Process within the Lithosphere Within the lithosphere act the physical and chemical forces that produce volcanism and elevate portions of the crust, forming land masses and making possible the many forms of life on land. Within the lithosphere, too, resides most of the force of the earth's gravity which holds the hydrosphere and atmosphere captive. Fluids and gases-mostly water and air-occupy pores and larger cavities within the outer portions of the crust. In regions where freezing occurs, small quantities of water in fractures and other open spaces in rocks alter their contours by the process of weathering. Mountains are slowly leveled and other features are altered by erosion, the wearing effect of wind, moving water, and ice (glaciers).

Relief features These are the elements of topography that give height and depth to the surface of the lithosphere.

First-order relief features The continents and ocean basins constitute the major irregularities in the almost smooth form of the lithosphere. The continents are the exposed areas of the true continental masses, the hearts of which consist of oval or shield-shaped areas of very ancient and greatly altered rock; these areas are known as shields. The continents extend as continental shelves beneath the shallower fringes of the oceans; the outer margins of the continents (continental slopes) drop steeply down into the actual ocean basins. Intermittently the continental shelves have been slowly laid bare and then slowly covered again by water; thus the shorelines of the continents have changed throughout the geologic past. The continents and ocean basins are presumed to be kept in balance by a principle called isostasy, whereby the lighter granitic layer of the continents "floats" in the heavier basaltic layer which constitutes the ocean basins and underlies the continents.

Biomoiecules

8 Lithification

. Sediments ··.·

Fig. 1.B. Diagrammatic representation of transitional events in the rock cycle.

One of the chief current problems in geology is that continental drift. It is concerned with whether the continents and ocean basins have always been permanently situated relative to one another, or if they have moved (and may still be moving) . . Second-order relief features Plains, plateaus, and mountains constitute the next most conspicuous aspects of topography. Both plains and plateaus are underlain by flat-lying rocks, the difference between them being merely a matter of elevation. Mountains may possess any of several kinds of structure. Second-order relief features exist within the ocean basins just as they do on land; abrupt, narrow mountain chains called ridges and long, narrow troughs called deeps contribute to the ruggedness of the sea floor. Especially interesting are the submarine canyons of the continental shelves, which rival any canyons found on dry land and which have never been adequately explained. Third-order relief features These are features of the land-escape of lesser importance than those mentioned above. They include a great variety of scenic effects such as glacial troughs (U-shaped valleys), sand dimes, lava flows, and fault scarps, due, in corresponding order, to erosion, deposition, volcanism, and earth movements. Hydrosphere A sphere of water, containing absorbed air and carrying particles of rock as sediment, surrounds the earth. Most of it lies within the ocean basins. (The extent and shape of' the .seas have varied greatly with the rising and Sinking of portions of the continents). Water also appears on the surface of the land in the form of lakes and running streams, which are important agents of erosion and

Origin of the Earth

9

transportation, and sites of deposition. A relatively small amount of the earth's water penetrates into the lithosphere; nevertheless, such ground water is of immense importance.

Atmosphere The envelope of air that embrace our planet contains absorbed water and small quantities of rock as dust, which may act as centers for the condensation of water vapour as clouds or fog. The chief component (78 per cent) of the atmosphere is nitrogen, but this gas is almost inert, as are the very small amounts of argon, neon, helium, krypton, xenon and other rare gases. The gases significant to man, and significant geologically, are oxygen (21 per cent) and carbon dioxide. Water vaopur (measured as humidity) is present in the air in amounts that vary with place and time. Energized by the heat of the sun, the hydrologic cycle involves the evaporation of water, mainly from the oceans, its circulation by air currents over the continents, its precipitation as rain or snow, and the ultimate return of most of it to the sea under the influence of gravity. GEOLOGIC PROCESSES

The geologic process operating upon and within the crust of the earth may be grouped under three main headings: gradation, volcanism, and diastrophism. Gradation This term encompasses the opposing processes of degradation and aggradations. Weathering, the decomposition and disintegration of rock, can be considered part of the process of gradation. Degradation This process, also called erosion, is the wearing down of rocks by water, ice, and wind. Aggradations This process, also called deposition, is the building up of rock layers by the accumulation of sediment, which is deposited by the action of water, ice and wind.

Volcanism This term refers to all movements of molten rock and the formation of solid rock from a molten state both within the lithosphere and on the surface. Diastrophism Included in this term are all movement of the solid parts of the earth, resulting in its displacement (faulting) or deformation (folding). The complex processes of metamorphism belong in large part at least under this heading, although volcanism is also involved. GEOLOGIC TIME

Although the evolutionary changes in life, as revealed by fossils, leave in the rocks their record of the passage of time, the comparisons are relative only. Neverthe1ess, the resolving power of biochronology is so great that it is able, for instance, to distinguish no less than 58 world - wide zones of ammonites in rocks of Jurassic age. The basic problem still remains, however, of determining the absolute ages, in years, of rocks in -the earth's crust and of accurately gauging the immensity of geologic time. Rate of Sedimentation Elaborate attempts have been made to determine the total age of the earth, or the duration of certain eras, by the rate of sedimentation. Time estimates are arrived at by measuring the thickness of layers of rock and multiplying these figures by the rate at which the layers are estimated to have been deposited. However, the process of sedimentation is so erratic that it is not a reliable indicator. The rate of erosion also varies, being more rapid today than at many times in the past. Salinity of the Ocean The rate at which salt is being added to the seas as the land is worn away has been used in an effort to calculate the length of geologic time. The varying rate of erosion and the lack of knowledge about several factors involved therein prevent this method from being useful.

Biomolecules

10

Law of Intrusion The age of sedimentary rocks can be determined on the basis of their spatial relationship to associated igneous rocks. The law of intrusion states that an igneous body that cuts across another rock is the younger of the two. (For extrusive rocks, the igneous body that rests upon another rock is the younger). Rocks that are metamorphosed must have been in position before any igneous intrusion that may have caused the metamorphism. Geologic Time Scale The construction of the geologic time scale or chart is one of the most significant achievements of geologists. This gargantuan calendar of the past is as fundamental to historical geology, and as indispensable, as the alphabet is to reading or the multiplication table to arithmetic. The oldest time intervals are shown, as is customary, at the bottom in the order in which the rocks of those ages normally occur in the earth. The time chart must be memorized, and it should be learned from the bottom up. Subdivision of Geologic Time _ The largest segments of geologic time are called eras. These are separated from one another by major orogenic (mountain-making) events known as revolutions, which were probably world-wide in scope and profoundly affected life on the earth. The eras are named according to the characteristic type of life that flourished, and the names are terms that end in zoic (from Greek zoe, life): Archeozoic (era of ancient life), Proterozoic (era of former life), Paleozoic (era of fold life), Mesozoic (era of middle life), and Cenozoic (era of recent life). The rock units of an era are spoken of as a group. Eras are divided into periods, which are separated from one another by less extensive local orogenies called disturbances; in some places these have been accompanied by uplift of the land and consequent withdrawal of the sea from the low parts of the continents most subject to repeated inundations_ Many of the period names refer to districts where rocks of that age were first studied, such as Devonian from Devonshire in England and Cretaceous in reference to the chalk (Latin creta) deposits on both sides of the English Channel. Others, such as the names Ordovician and Silurian, are less obvious. These are derived from the names of two Celtic tribes that during Roman times inhabited the parts of Wales in which these formations were first studied. An exception to this locality nomenclature is the division of the Cenozoic era into the Tertiary ("third') and Quaternary ("fourth') periods. The rock units that belong to a given period constitute a system. The periods are subdivided into epochs, which are often separated by retreats- of the sea on a regional scale. The more recent epochs, about which more geologic detail is known, are those of the Cenozoic era. The epochs are designated in two ways: (1) periods of the earlier eras (Paleozoic and Mesozoic) are subdivided into Early, Middle and Late ( as Early Devonian); and (2) the Cenozoic era is divided as follows: the Tertiary period into the following epochs: Paleocene (paleo-,old), Eocene (eo-, dawn), Oligocene (oligo-, a few), Miocene (mio-, less), and Pliocene (plio, more); and the Quaternary period into the Pleistocene (pleisto-, more) and Recent epochs. The rock units of a given epoch are spoken of as a series, and may be further designated as Lower, Middle, and Upperor geographic names may be used. Epochs are subdivided into ages. The rock units that belong to a given age constitute a stage, which may be subdivided into sub-stages. Still smaller units, zones, are named according to their characteristics assemblage of fossils. In as much as organic evolution is affected by changes in the environment, divisions of time correspond, at least in a general way, to the transformation in the structure and topography of the earth. Thus geologic time is based ultimately upon the interplay of physical and biologic changes. UNCONFORMITIES

Gaps ,in the geologic record are known as unconformities. Certain layers may have been removed by erosion before later ones were added, or there may have been no deposition at all during

Origin

11

0/ the Earth

an interval of time. If the beds are parallel above and below the unconformity, the structure is called a disconfonnity. If the upper rocks lie upon an eroded surface of a different major type of rock (as sedimentary on igneous or metamorphic), the relationship is called nonconfonnity. The most conspicuous kind of unconformity is the angular unconformity, where two groups of sedimentary rocks with different angles of dip are in depositional contact, indicating the that the lower sequence of strata was tilted or folded before being eroded and covered by the upper sequence. PALEOGEOGRAPHY

Because the land is constantly warping and the sea level is fluctuating, the present geography of the surface of our planet is a transient thing, its history brief in terms of geology. The study of ancient lands and seas is paleogeography. On paleogeographic maps are reconstructed the patterns of continents and oceans, island and seaways of the past, as determined from the evidence of the rocks and their fossils. Paleoclimatology, the study of ancient climates, examines the distribution of soils, sand dunes, salt deposits, coral reefs and other fossils, and glacial deposits. The roots of vanished mountains, with their folded, faulted, intruded, and metamorphosed structures, bear testimony to 'the ranges that once existed. Fossils of water-dwelling organisms indicate the former distribution of water and whether it was salt or fresh, deep or shallow, Fossils of land organisms tell much about the former climate, altitude, and topography. The agents of erosion, transportation, and deposition are identified by the nature of the sediment and the forms it has assumed. The source of sediment is indicated by the direction in which it becomes coarser and thicker. In these ways the story of the history of the earth is patiently pieced together by thousands of geologists.

2 ORIGIN OF BIOMOLECULES THE BIOLOGICAL FITNESS OF ORGANIC COMPOUNDS

The fact that the elementary composition of living matter is very different from that of the lithosphere and atmosphere suggests that some chemical elements are more "fit" than others to make up the molecules of living organisms. Only 22 or the 100 chemical elements found in the earth's curst are essential components in living organisms, and of these, only 16 are found in all classes of e>rganisms. Moreover, the distribution of these chemical elements in living organisms is not in proportion to their occurrence in the earth's crust. The four most abundant elements in living organisms are hydrogen, oxygen, carbon, and nitrogen; they make up about 99 percent of the mass of most cells, whereas the four most abundant elements in the earth's crust are oxygen, silicon, aluminium, and sodium. Actually, carbon, hydrogen, and nitrogen are far more abundant in living matter than in the earth's crust. We may therefore pressure that compounds of these processes that collectively constitute the living state. Table 2.1. The relative abundance of some chemical elements in the earth's crust

Element 0 Si AI

Na Ca Fe Mg

P C N

Atoms percent 62.5 21.2 6.47 2.64 1.94 1.92 1.84 1.42 0.80 0.0001

The four elements carbon, hydrogen, nitrogen, and oxygen possess a common property: they readily form covlent bonds by electron-pair sharing. Hydrogen needs one electron, oxygn two, nitrogen three, and carbon four to complete their outer electron shells and thus form stable covalent bonds. All four elements can readily react with each other to fill their outer shells. Furthermore, three of these elements (C, N, and 0) can share either one or two electron pairs to yield either single or double bonds, a capacity which endows the with considerable versatility of chemical bonding. Carbon is also capable of forming triple bonds with other carbon or nitrogen atoms. This type of bonding occurs only rarely in nature. 12

13

Origin of Biomolecules

Carbon, nitrogen, hydrogen, and oxygen are uniquely fit in another way: they are the lightest elements capable of forming covalent bonds. Since the strength of a covalent bond is inversely related to the atomic weights of the bonded atoms, it appears that living organisms have selected those elements capable of forming the strongest covalent bonds. Table 2.2. The relative abundance of some chemical elements in the human body

Element H 0 C N Na Oi P S K CI

Atoms percent

60.3 25.5 10.5 2.42 0.73 0.226 0.134 0.132 0.036 0.032

Particularly significant is the capacity of carbon atoms to interact with each other to form stable, covalent carbon-carbon bonds. Since carbon atoms may either accept or donate four electrons to complete an outer octel, each carbon atom can form covalent bonds with four carbon atoms. In this way, covalently-linked carbon atoms may constitute the backbones for an immense variety of different organic molecules. Moreover, since carbon atoms readily form covalent bonds with oxygen, hydrogen, and nitrogen, as well as with sulfur, a large number of different kinds of functional groups can be introduced into the structure of organic molecules. Moreover, since carbon atoms readily form covalent bonds with oxygen, hydrogen, and nitrogen, as well as with sulfur, a large number of different kinds of functional groups can be introduced into the structure of organic molecules. Organic compounds of carbon have yet another distinctive feature. Because of the tetrahedral configuration of the shared electron pairs around each carbon atom, different types of organic molecules possess different threedimensional structures. No other chemical element can form stable molecules of such widely different sizes and shapes, nor with.such a variety of functional groups. Silicon is the only other element that po~-this capacity to combine with itself by electron-pair sharing. Although it is far more abundant in the lithosphere, silicon is evidently inferior to crbon for the purposes of living organisms. Perhaps the major reason is that silicon-silicon bonds are unstable in the presence of oxygen, leading to the formation of silicates and insoluble silicon dioxide polymers, such as quartz. One other point: The organic compounds of carbon found in living organisms are highly reduced, or hydrogenated, whereas the carbon of the earth's curst is largely present in oxidized form as bicarbonates or carbonates. Because oxygen is very abundant in the atmosphere, carbon and hydrogen normally tend to .become oxidized to carbon dioxide and water, compounds which are stable and energy-poor The reduced orgatiic molecuels found in living matter are energy-rich since, to make thtml ffom CO2 and water, living organisms must expend free energy. We may therefore conclude that organic carbon compounds must be especially well suited for the purposes of living organisms, since they were selected despite the relative sparseness of carbon in the lithosphere and despite the fact energy must be expended to reduce inorganic carbon.

The Hierarchy of Molecular' Organization of Cells The biomolecules of living organisms are ordered into a hierarchy of increasing molecular complexity. All organic biomolecules are ultimately derived from very simple, low-molecular-weight precursors obtained from the environment, namely, carbon dioxide, water, and atmosphere nitrogen. These precursors are converted by living matter, via sequences of metabolic intermediates of increasing

14

Biomolecules

The cell

molecular size, into the buildingblock biomolecules, organic

i

compounds of intermediate molecular weight. These building blocks are then Nucleus linked to each other covalently to form Organelles Mitochondria the macromolecules of the cell, Chloroplasts which have relatively high molecular weights. Thus the amino acids are the building blocks of the proteins, the Enzyme complexes mononucleotides are the building Supramolecular Ribosomes blocks of the nucleic acids, the assemblies Contractile systems monosaccharides are the building blocks of the polysaccharides, and fatty acids are building blocks of most lipids. At the next higher level of Macromolecules lipids PolyNucleic Proteins organization, macromolecules ,of saccharides acids different classes associate with each other to form supramolecular complexes, such as lipoproteins, MonoSimple Fatty Amino which are complexes of lipids and Building blocks sugars nucleotides acids acids, proteins, and ribosomes, which, in glycerol tum, are complexes of nucleic acid and proteins. However, there is now a distinctive difference in the manner in which the components are ass- Intermediates Ribose, a-Keto Phospho- Acetate, pyruvate, malonate carbamyl acids embled. In supramolecular complexes, phosphate malate the component macromolecules are not covalently bonded to each other. For example, the nucleic acid and protein components of ribosomes are Precursors from CO2 not attached to each other covalently; the environment H 20 rather they are "stuck together" by N2 weak noncovalent forces, such as ionic Fig. 2.1. The hierarchy of molecular organization In cel/s. interactions, hydroge bonding, hydrophobic interactions, and van der Waals interactions. Nevertheless the rroncovalent association of macromolecules into supramolecular complexes is very specific and, usually, very stable, the result of the precise geometrical "fit" or complementarity between and among the component parts. There are good reasons for cells to employ noncovalent rather than covalent interactions to form supramolecular complexes from macromolecular components. At the highest level of organization in the hierarchy of cell structure, various supramolecular complexes are further assembled into organelles, such as nuclei, mitochondria, and chloroplasts, and into other bodies and inclUSions, such as lysosomes, microbodies, and vacuoles. Here again, so far as is known, the various components are associated by noncovalent interactions. The four major types of biomacromolecules also have identical functions in all species of cells. The nucleic acids universally function to store and transmit genetic information. The proteins are the direct products and effectors of gene action, and into them the genetic informaion is incorporated. Most proteins have specific catalytic activitY and function as enzymes; others serve as structural elements. Many other biological functions are served by proteins, which are the most versatile of all biomolecules. The polysaccharides have two major functions: some, such as starch, serve as storage forms of

i

/ i

i i

i i

~

i

i

i

r

~V

i

15

Origin of Biomolecules

energy-yielding fuels for cell activity; and some, such as cellulose, serve as extracellular structural elements. The lipids serve two chief roles: as major structural components of membranes, and as a storage form of energy-rich fuel.

Table 2.3. Molecular components of an E.coli cells Percent total weight Water Proteins Nucleic acids DNA RNA Carbohydrates Lipids Building-block molecules and intermediates Inorganic ions

70 15 1 6 3 2

2 1

Number of each kind

-3,000 1 -1.000 -50 -40

-500 12

There is an important and fundamental difference between the nucleic acids and proteins on the one hand and the polysaccharides and lipids on the other. nucleic acids and proteins are infonnational macromolecules by virtue of their structure. Each nucleic acid molecules contains four or more types of mononucleotides arranged in a specific information-rich sequence. Similarly, each protein molecule contains a specific informationrich sequence of some 20 different amino acids. On the other hand, the various polysaccharides do not bear information; their recurring building blocks either are all identical, as is the case in starch, a polymer of D-glucose, or consist of only two types of sugar building blocks, which merely alternate. Similarly, lipids are noninformational, since their fatty acid components are also units that have two carbon atoms.

Noninformational

Informational

~--------~I~I----------~

I A

I Arg

I

T

I Ala

Met

G\c I Glc

I G I

I Trp I

I C I G

I Asn I Glu

C I G I G I

Phe I lie I lie I

A.

I A I

I Tyr I

A I

Tyr I

T

I T I G I

I

G\c I Glc I Glc

I

Glc

I

Glc

I

Glc

I

t

G\c

I

I

g

~he

Nucleotide sequence in a DNA molecule

I

Glc

Glc I Glc

I

~

I

Ala

I Lys I Lys

b

I

G\c I G\c

Amino acid sequence

p;~t:in

Repeating glucose units

pol~~charide

molecule

Fig. 2.2. Informational and nonlnfor-

mational macromolecules of the cell

constructed from repeating, identical

The Primordial Biomolecules It is seen that the immensely large number of different proteins and nucleic acids in living matter are made from a small number of different building-block molecules, which are identical in all species of living organisms. Recent studies of the chemical composition of the simplest cells, among them the Mycoplasma, suggest that the first cells to have arisen on earth may have been built from only some 30 different organic molecules. This set of 30 primordial biomolecules includes 20 amino acids, five nitrogenousaromatic bases, a fatty acid, two sugars, the alcohol glycerol, and the amine choline. In fact, this list may be shortened to 25, since recent research on the genetic code suggess that the first living cells required only 16 amino acids rather than the 20 known to be present in proteins today. Whatever their precise number, the primordial biomolecules may be regarded as the ancestors of all other biomolecules; they are the first alphabet of living matter. Although many of the primordial biomolecules appear at first glance to be chemically umelated to each other, they are in fact, related through the enzymatic reactions of metabolism, which we have seen proceed through consecutive reactions having common intermediates. For example, although the sugar glucose and the fatty acid palmitic acid and the amino acid alanine appear to be wholly different molecules, it has been found as a result of isotope-tracer and metabolic studies that all the carbon atoms of glucose can be used by living cells to form the carbon skeleton .of alanine and that four of

Biomolecu/es

16

the six carbon atoms of glucose can be converted into the carbon skeleton of palmitic acid via the intermediate acetic acid. Many other metabolic interconversions of the primordial biomolecules are known to take place. It therefore appears likely that the primordial biomolecules were particularly suited to be the components of living matter not only because of their intrinsic struchtres and properties but also because feasible chemical pathways existed for their enzymatic interconversion.

The Specialization and Differentiation of Biomolecules As living organisms evolved into more highly differentiated and complex forms, new biomolecules of greater complexity and variety also evolved. These more specialized and differentiated biomolecules are structural and metabolic derivatives of the 30 primordial biomolecules. For example, over 150 different biologically occurring amino acids are known today. Nearly all of these are derived from the basic 20 amino acids used for the construction of proteins. Similarly, dozens of different nucleotides and nucleotide derivatives are known, all descendants of the five major nitrogenous bases found in nucleic acids. Over 70 simple sugars derive from glucose, and from these a large variety of polysaccharides are formed in different organisms. There are many different fatty acids, which are all descended from palmitic acid.

Table 2.4. Some specialized derivatives of the primordial biomolecules Arginine Ornithine Citrulline Proline. 3-Hydroxyproline 4-Hydroxyproline 4-Hydroxymethylproline 4-Methyleneproline 4-Ketoproline Laucine ~-Hydroxyleucine

B-Hydroxyleucine y-B-Dihydroxyleucine y-Hydroxyleucine N-Methylleucine Guanine I-Methylguanine 2-Methylguanine 2-Dimethylguanine 2-o-Methylguanine 7 -Methylguanine

D-Glucose D-Mannose D-Fructose D-Galactose N-Acetylglucosamine D-Glucuronic acid D-Glucose 6-phosphate Ascorbic acid Inositol Sucrose Maltose Lactose Palmitic acid Oleic acid Stearic acid Lauric acid Palmitoleic acid Palmitaldehyde Stearaldehyde

Many specialized biomolecules known today are extremely complex and appear to bear little resemblance to the 30 primordial biomolecules. Among these are igments, odor-bearing essential oils, waxes, hormones, molecules such as antibiotics and alkaloids that are toxit to some organisms, and various structural molecules, such as lignin of wood. Nevertheless, recent research on the biogenesis of many of these substances shows that they can be classed into a few different types, all of which are ultimately derived from the primordial biomolecules or their breakdown products. Among these are the acetogenins, so named becaused they are formed by head-ta-tail condensations of acetic acid, a degradation product of both glucose and fatty acids, and the terpenes, which are built of 5-carbon isoprene units, which also derive from acetic acid. Most of the alkaloids are in tum derived from the promoridal amino acids.

17

Origin of Biomolecules THE ORIGIN OF BIOMOLECULES

We have seen that organic compounds occur only in traces in the earth's crust. How is it, then, that the first living organisms, which were presumably formed from chemical components in the primitive sea, acquired the primordial biomolecules? Recent research suggests that early in the history of the earth, conditions favoured the existence of many different organic compounds in relatively high concentration in the surface waters of the ocean and that the first living cells arose in this warm "soup" of organic compounds. It is now generally believed that the earth is approximately 4.6 billion (4.6 x 109) years old. It first arose from condensation of interstellar gases and dust, which culminated in the formation of a firm, solid mantle, consisting largely of iron and magnesium silicate and surrounding a central core of partially molten iron and nickel. Geological evidence indicates that the otuer crust of the earth stabilized about 4.0 to 4.5 billion years ago. Living organism probably first arose earlier than 3.5 billion years ago; in fact, there is fossil evidence that bacteria similar to those known today already existed about 3.1 billion years ago. In the late 1920s. AI. Oparin, a biochemist in the Soviet Union, suggested that chemical and physical processes occurring in the primitive atmosphere could have led to the spontaneous formation of simple organic compounds, such as amino acids and sugar, from methane, ammonia, and water vapour which he postulated to be present in the primitive atmosphere. According to his theory, these gases were activated by the radiant energy of sunlight or by lightning discharges to react with each other. The simple organic products so formed condensed and dissolved in the primitive ocean, which gradually became enriched in a large variety of organic compounds. Oparin postulated that the first living cell arose spontaneously from this warm, concentrated solution of organic compounds. Oparin's hypothesis was not immediately accepted for lack of evidence. There has been continuing disagreement about the constitution of the atmosphere during the period when life is believed to have begun, particularly whethr it containeq the relatively reduced gases methane and ammonia. The most recent research indicates that the atmosphere 3.0 to 3.5 billion years ago was rich in nitrogen, hydrogen, carbon monoxide and carbon dioxide, but that its content of the reduced compounds methane and ammonia was probably not very high; free oxygen was not present. The temperature of the earth's crust and atmosphere during this period, but well before 100°C. Most of the surface of the earth was covered with water. water } ammonia That the gaseous components thought to be present methane in the primitie atmosphere can be precursors of organic hydrogen compounds is now well supported by laboratory studies. Among the early experiments on the abiotic origin of organic molecules were those carried out in 1953 by Miller. He subjected gas mixtures of methane, ammonia, water, and hydrogen, then believed to be predominant in the primitive atmosphere, in a closed flask at BOoC to electrical sparking across a pair of electrodes, to simulate lighting, for periods of a week or more. Then he collected and analyzed the contents of the system. The gas phase was found to contain carbon monoxide, carbon dioxide, and nitrogen, which were evidently formed from the gases initially introduced. In the chilled condenseate, he found boiling water significantly large amounts of water-soluble organic substances, which he separated by chromatographic Fig. 2.:3. Miller's experiment.

--~·t

t

Biomolecules

18

methods. Among the compounds Miller identified were Table 2.5. Some organic compounds a number of a-amino acids, including some present in generated by spark discharges under primitive-atmosphere conditions proteins, such as glycine, alanine, aspartic acid, and glutamic acid. He also found several simple organic acids Glycine known .to occur in living organisms, such as formic, Alanine acetic, protionic, lactic, and succinic acids (Table 2.5). Sarcosine Miller postulated that the various organic ~-Alanine compounds formed in these experiments arose by the a-Aminobutyric acid sequence of reactions shown in Table 2.6. In the first N-Methylalanine reaction, hydrogen cyanide was formed from methane Aspartic acid and ammonia. Methane also was converted by the Glutamic acid electrical discharge into ethylene and other h>drocarbons. Iminodiacetic acid The hydogen cyanide reacted with the ethylene to form Iminoacetopropionic acid a nitrile (reaction 2), which then underwent hydrolysis to propionic acid (reaction 3). Similarly, aFormic acid hydroxynitriles reactd with ammonia to form aAcetic acid aminonitriles (reaction 4), which then underwent Glycolic acid hydrolysis to form a-amino acids, such as alanine Lactic acid (reaction 5). Miller's experiments were carried out in a a-Hydroxybutyric acid system rich in the reduced compounds methane and Succinic acid ammonia. In more recent experiments in which mixtures Urea containing nitrogen, hydrogen, carbon monoxide, and Methylurea carbon dioxide, but no methane or ammonia, were exposed to radiant energy, amino acids and other organic molecules were again formed, also via hydrogen cyanide, showing that preformed ammonia and methane are not essential for the abiotic formation of organic molecules. Table 2.6. Chemical reactions in spark discharges CH 4 + NH2 C2H 4 + HCN CH2CH2CN CH2CHOHCN

> HCN + 3H 2 >CH2CH2CN

(1)

(2)

A nil'ile

HO 2 > CH3CH2COOH Propionic acid

> CH 3CHNH2CN

(3) (4)

An aminonitrile (5)

Many different forms of irradiation have yielded. organic compounds from such gas mixtures, including visible light ultraviolet light, x-rays, gamma radiation, sparking and silent electrical discharges, ultrasonic irradiation, and high-energy a- and ~-particles. Several hundred different organic compounds havebeen formed in such experiments, including representatives of all the important types of molecules found in cells, as well as many not found in cells. All the common amino acids presnt in proteins, the nitrogenous bases adenine, guanine, cytosine, uracil, and thymine, which serve as the building blocks of nucleic acids, and many biologically occurring organic acids and sugars have been detected. In view of these results, it now appears quite plausible that the primordial ocean was rich in dissolved organic compounds and that these may have included many or all of the basic building-block molecules we yecognize in living cells today.

Origin of Biomolecules

19

We may now ask: Why do abiotically formed organic molecules no longer exist in the surfae waters of the earth, and why must present-day organisms synthesize their own biomolecules? The atmosphere and surface of the earth apparently underwent substantial changes during and after the period when life began, with the result that the rate of formation of new organic molecules declined. For one thing, the surface of the earth gradually cooled. Hydrogen and carbon monoxide were gradually lost from the atmosphere and were replaced by oxygen and carbon dioxide. In all probability, the concentration of abiotic organic compounds in the oceans was 'also reduced as a result of he metabolic activity of primitive organisms, whch used them as fuel and as building blocks for forming new cells. As the abiotic organic compounds in the sea became increasingly scarce, those organisms that were capable of internally synthesizing their own biomoleculs from simpler precursors such as CO 2 , HP, ammonia, and nitrogen had a selective advantage. Presumably it was under these conditions that the first photosynthetic cells arose, capable of utilizing light energy to convert atmosphere CO 2 to glucose and other cell components.

The Fitness of Biomolecules Why should living organisms have selected the specific types of organic molecules which they now possess? Why should 20 a.-amino acids be the bUilding blocks of proteins? Why not only 10? Why not 40? Why are they all a.-amino acids? Could we not equally well construct large "protein" molecules from amino acids having their amino groups in the ~-positions? Why should be purines adenine and guanine and the pyrimidines cytosine and thymine have been selected, out of the dozens of purine and pyrimidine derivative known, to be the essential building blocks of DNA in all species? Current evidence supports the concept that the biomolecules we know today were selected from a much larger number of available compounds because of their special fitness, which gave cells containing them superior survival value. Since several hundred orgaic compounds have been isolated during experiments on the abiotic origin of organic molecules, such as those described above, and only some 30 different compounds may have been required to form the first cells, it appears very likelythat a process of selection took place. Evidence for the fitness concept is inherent inother facts already mentioned. Over 150 different amino acids have been found to occur biologically, yet all proteins in all species are built from the same set of 20 primordial amino acids. If any of the other amino acids had proved to be more "fit" as protein building blocks than the primordial amino acids, there was ample evolutionary time available for organisms to have acquired the ability to 'use them. This would particularly be the case with bacteria, which, because of their short generation time, could have selected the fittest amino acids for protein structure faster than other kinds of organisms. Yet the proteins of bacteria today are madeof he same building-block amino acids as those in all other species. Since the newer, more highly specialized aminoa cids 'are more comples than the primordial amino acids from which they were derived, we may also conclude that a given biomolecule is no more coplex than it needs to be to fulfill its function. It may sometimes appear that a given biomolecule is less suited for a gien function than some other substance could be. However, we must remember that a biomolecule usually has more than one biologically important funtion or property, and its biological fitness must be assessed on the basis of all its functions. For example, although chlorophylls are the universal light-trapping pigments of photosynthetic cells, they are not particularly effective light absorbers at thsoe wavelengths of sunlight that are transmitted to algal cells through the turbid water of a pond. The chlorophyll molecule possesses other important properties whose advantages greatly outweight this particular disadvantage. We conclude that each biomolecule possesses properties that are optimal for all its biological functions considered together. The Dimensions and Shapes of Biomolecules In the next we shall examine the structures and properties of the major classes of biomolecules. In particular, we shall take special notice of the sizes and shapes of biomolecules, since these attributes

Blomolecules

20

are of great significance in biochemistry and molecular biology. We have already seen that the precision of the complementary fit of the substrate to the active site of an enzyme is so great that it makes possible the selectivity of enzymatic catalysis and the absence of by-products. For example, the enzyme trypsin hydrolyzes only those peptide bonds of proteins in which the amino acids lysine or arginine contribute the carbonyl group. We now know that a change of but a fraction of an angstron unit in some critical dimension of a substrate molecule can cause it to become inactive as a substrate. It is therefore essential to become familiar not only with the dimensions of biomolecules, but also the dimensions of cells and their components. The basic units of mass and length shown below.

Units of mass and length

Mass 1 dalton

mass of one hydrogen atom 1.67 x 10-24 gram

Length 1 meter (m) = 1000 millimeters (mm) 1 micron (11) = 1 micrometer (11m) 1000 millimicrons (mil) 1000 nanometers (nm) 1 nm = 10 angstrons (A) The flat, two-dimensional projections in which the structures of organic molecules are necessarily shown on the printed page are quite insufficient to describe the true, three-dimensional configuration of biomolecules. The configuration of a molecule in space can be represnted by the use of atomic models, of which there are two general types. Crystallographic models show the covalent skeleton and the bond angles and lengths, but they do not indicate the actual space occupied by the molecule. Space-filling models, on the other hand, show little detail with regard to bond angles and distances in the backbone, but they do show the van der Waals contour or surface of the molecule. While both types of model are useful in studying the structure of biomolecules, it isthe space-filling model that represents the molecule as it is "seen" by the cell or by one of its specific components, such as an enzyme. Actually, an enzyme "sees" much more than the three-dimenstional shape of its substrate. It see the location and sign of its electrical changes and the precise distance between them. It se~s the positions of uncharged polar groups, such as hydroxyl, carbonyl, and amide groups, which can potentially enter into hydrogen-bond formation. It sees the sizes and shapes of the nonpolar hydrocarbon areas on the surface of the biomolecule, which may provide important contact areas in hydrophobic interactions. Three-dimensional shape and surface topography are especially important in the case of macromolecules, particularly the proteins. A protein has only one specific thre~imensional conformation under normal intracellular conditions, called the native conformation, which is indispensable for its biological activity. Only the native conformation of an enzyme molecule has catalyticactivity. The three-dimensional confonnation of a large biomolecule cannot be extrapolated from a two-dimenstional structure on the printed page, nor can it be arrived at easily ith ordinary space-filling models. A very complex physical method, namely, x-ray diffraction analysis, is required to establish the precise conformation of a biological macromolecule. To date, the native conformations of only a few proteins are known exactly, but the charting of biomolecular structure by x-ray and computer methods has become a major objective of molecular biology. There is another way in which the sizes and shapes of biomolecules are of crucial importane. We have seen that in living cells there is a hierarchy of molecular organization, simple biomolecules are the building blocks of macromolecules, macromolecules, in turn, are the building blocks of supramolecular complexes, and supramolecular complexes are the building blocks of organelles. The

21

Origin 0/ Biomolecules

dimensions of biomolecules mut therefore be related to the dimensions of cells and their components (Table 2.7), just as the sizes and shapes of the component building blocks determine the sizes and shapes of cell macromolecules, such as proteins, so may the sizes and shapes of the component macromolecules determine the sizes and shapes of supramolecular complexes and thus, ultimately, of the cell organelles.

Table 2.7. Approximate dimensions and weight of some biomolecules and cell components Long dimension (A)

Alanine Glucose .. Phospholipid Myoglobin (a small protein) Hemoglobin (a medium-size protein) MyOsin (a large rod-shaped protein) Glutamate dehydrogenase130 (a very large globular protein) Ribosome of E.coli Bacteriophage X174 of E.coli Tobacco mosaic virus (a rod) Mitochondrion (liver cell) E.coli cell Chloroplast (spinach leaf) Liver cell

5 7 35 36 68 1,600 1,000,000 180 250 3,000 15,000 20,000 80,000 200,000

Weight (daltons)

89 180 750 16,900 65,000 470,000

2,800,000 6,200,000 40,000,000 1 x 10-12 gram 2 x 10-12 gram 1.3 x 10-10 gram 2 x 10-9 gram

Biomolecules in Relation to Cell Structure Throughout this book we shall relate the structures and dynamic functions of each type of biomolecule to the structures and functions of cell components, such as mitochondria, contractile systems, ribosomes, endoplasmic reticulum, chloroplasts, membranes, and cell walls. Before we begin the detailed study of biomolecules, therefore, it maybe useful to review the major structural elements of cells, with special reference of their dimensions and molecular composition. The first cell is the bacterium E.coli, which is the best-known member, biochemically and genetically speaking, of the great class of procaryotic cells. The second is the hepatocyte, or liver parenchymal cell, perhaps the most thoroughly studied eucaryotic cell. Both the E.coli cell and the hepatocyte are heterotrophic cells, which require carbon from the environment in a complex, reduced form. Biochemistry today is increasingly concerned with the structure of cell and their organelles. Some of the most illuminating recent progress has come from correlated biochemical and electronmicroscopic studies of cellular processes. Indeed, the dividing lins between biochemistry and cell biology, and between biochemistry and molecular biologysics, are becoming difficult to identify, since these areas form a logical continuum. The application of exact physical methods to the analysis of cell components is yielding significant data on the precision with which biom9lecules are constructed and ' with which they interact in their cellular functions.

3 ATOMS AND MOLECULES Living cells and organisms are made up of many kinds of molecules. These molecules, upon isolation, conform to the same laws of physics and chemistry that apply to lifeless matter. The living organism exhibits properties in addition to those described by the laws of physics and chemistry. Thus these biomolecules, the molecules characteristically found in living material, exhibit remarkable properties when they are arranged in complex and highly organizd living matter. This highly organized state is maintained by a constant supply of energy which must be provided for it. Each component of a living organism appears to have a specific function: This is true for large organs, such as the liver and brain, down to the individual molecules, such as nucleic acids, proteins, acetogenins, carbohydrates, and water. These molecules interact with each other in what has been called the molecular logic of the living state. One purpose of Part I is to explore some aspects of this logic, that is, to study the structures and functions of biomolecules and how they interact in forming the physical and chemical basis for living processes.

Physical and Chemical Properties of Atoms and Molecules If living material is composed of intrinsically lifeless molecules, why is living matter so different from lifeless matter? Put another way, why does the living organism qppear to be mqre than the sum of its inanimate parts? To answer these questions are the goals of those biologists who study biomolecules. Some understanding of the physics and chemistry of atoms and molecules is desirable before we explore the biomolecules themselves. The beautiful order and logic of atomic and molecular structure is reflected in the structures of the fantastically cO\'Tlplex .biomolecules of living organisms. Elements are not Changed in Chemical Reactions Elements are substances that cannot be broken down into simpler substances by chemical reactions. The matter of the universe is composed of 92 naturally occurring elements, ranging from hydrogen, the lightest, to uranium, the heaviest. Just four elements-oxygen, carbon, hydrogen, and nitrogen-are responsible for over 96% of the mass of most organisms. Others, such as calcium, phosphorus, potassium, and magnesium, are also consistently present but in smaller quantities. Some elements, such as iodine and copper, are known as trace elements because they are present only in minute amounts. Chemical elements cycle between organisms and the nonliving environment. Scientists have assigned each element a chemical symbol: usually the first letter or first and second letters of the English or Latin name of the element. For example, 0 is the symbol for oxygen, C for carbon, H for hydrogen, N for nitrogen, and Na for sodium (Latin natrium). Table 3.1. lists the elements that make up two representive organisms, a human and a typical non-woody plant, and briefly explains why each is important.

22

-Atoms and Molecules

23

Table 3.1. Elements that make up some representive organisms Approximate % of Total Approximate Element % of Total Mass of and Nonwoody Mass of Chemical Human Body Plant Importance or Functions Symbol Oxygen (0) Required for cellular respiration; present in most 65 78 organic compounds; component of water 18 Forms backbone of organic molecules; can form Carbon (C) 11 four bonds with other atoms Hydrogen (H) Present in most organic compounds; component 10 9 of water; hydrogen ion (H+) is involved in energy transformations * Nitrogen (N) Component of proteins and nucleic adds; 3 component of chlorophyll in plants Calcium (Ca) 15 Structural component of bones and teeth; calcium * ion (Ca 2+) is important in muscle contraction, conduction of nerve impulses, and blood clotting; associated with plant cell wall , Phosphorus (P) Component of nucleic acids and of phospholipids 1 * in membranes; important in energy transfer reactions; structural component of bone * Potassium (K) Potassium ion (K+) is principal positive ion (cation) * in interstitial (tissue) fluid of animals; important in nerve function; affects muscle contraction; controls opening of stomata in plants Sulfur (S) * Component of most proteins * Sodium ion (Na+) is principal positive ion (cation) Sodium (Na) * * in interstitial (tissue) fluid of animals; important in fluid balance; essential for conduction of nerve impulses; not essential in most plants • • Magnesium (Mg) Needed in blood and other tissues of animals; activates many enzymes; component of chlorophyll in plants • • Chloride ion (Cl-) is principal negative ion (anion) Chlorine (CO in interstitial (tissue) fluid of animals; important in water balance; essential for photosynthesis Iron (Fe) * Component of hemoglobin in animals; component * of cytochromes; activates certain enzymes • The asterisk indicates that these elements represent less than 10/0 of the total mass. Other elements found in very small (trace) amounts in animals, plants, or both include iodine (0, manganese (Mn), copper (Cu), Zinc (Zn), cobalt (Co), fluorine (F), molybdenum (Mo), selenium (Se), boron .(8), . silicon (Si), and a few others.

Atoms are the Basic Particles of Elements An atom is the smallest portion of an element that retains its chemical properties. Atoms are much smaller than the tiniest particle visible under a light microscope. By special scanning electron microscopy, with magnification as high as x 5 million, researchers have been able to photograph the positions of some large atoms in molecules.

Biomolecules

24 ATOMS CONTAINS PROTONS, NEUTRONS, AND ELECTRONS

Physicists have discovered a number of subatomic particles, but for our purposes we need consider only three: protons, neutrons, and electrons. Each proton has one unit of a positive electrical charge; a neutron is an uncharged particle with about the same mass as a proton. The mass of a proton or neutron is exceedingly small, much too small to be conveniently expressed in terms of grams or even micrograms. Such masses are expressed in terms Neutron of the atomic mass unit (amu) , also called the dalton in Nucleus honor of John Dalton, who formulated an atomic theory in Proton the early 1800s. An amu is equal to the approximate mass of a proton or neutron. Protons and neutrons make up almost all of the mass of an atom and are concentrated in the atomic nucleus. Each electron has one unit of negative electrical charge and an extremely small· mass (only about 1/1800 of the mass of a proton). The number of electrons in an electrically neutral atom equals the number of protons. the electrons do not have fixed locations, but are moving rapidly in the space qutside the atomic nucleus. The characteristics of protons, electrons, and neutrons are summarized below:

Particle

Charge

Mass

Location

Proton Neutron Electron

positive neutral negative

1 amu 1 amu negligible

nucleus nucleus outside nucleus

An Atom is Uniquely Identified by its Number of Protons Each kind of element has a fixed number of protons in the atomic nucleus. This number, called the atomic number, is written as a subscript to the left of the chemical symbol. Thus 1H indicates (_) that the hydrogen nucleus contains one proton, and 80 that the oxygen nucleus contains eight protons. It is the atomic number, the number of protons in its • e nucleus, that determines an atom's identity. • Protons Plus Neutron Determine Atomic Mass The atomic mass of an element is a number Hydrogen that indicates how massive an atom of that element is Carbon compared with an atom of another element. The value is determined by adding the number of protons to the number of neutrons and expressing the resulfin atomic mass units or daltons. The mass of the electrons is ignored because it is so small. The atomic mass is indicated by a superscript to the left of the chemical symbol. The common form of the oxygen atom, with eight protons and eight neutrons in its nucleus, has an atomic number of 8 and a mass number of 16 Fig. 3.1. Atomic structure of three elements haVing atomic mass units. It is indicated by the symbol 8160. differing numbers of electrons (e), protons Scientists commonly refer to the atomic weight (P), and neutrons (N). of an element. This is calculated as the ratio of its atomic mass compared with the atomic mass of the common form of carbon (12 amu). When this ratio is calculated, the atomic mass units cancel out. For this reason, atomic weight and atomic mass are numerically equal, but atomic weight is dimensionless. Therefore, while the atomic mass of the common form of oxygen is properly expressed as 16 atomic mass units, the atomic weight is simply 16.

@

226.025

(223)

91

Pa

90

Th

Cm (247)

Am (243)

Pu (244)

96

Np

Th

65

200.59

Pb

66

Dy

Cf (251)

Bk (247)

98 97

208.980

164.930

(252)

Es .

(257)

Fm

100

167.26

68

Er

67

Ho 99

Xe

54

..,.80

I. Ul

2 o

.c

II)

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OE Q)Q)

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t:

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~

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'"t:

-"0 CQ)

ti

:g a.

00 UlItS

~fi§

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Wc

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LL.l ~

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:.::® @ ® ® = Q)

from the nucleus, may have -liS ~ ~ different shapes. Regardless of ~ d: its sh~pe, a given orbital may • ..~ E contam no more than two i5 'ai £i ~ electrons. ~ -liS ::; (J) ...J Almost all of the volume '----v---" ''-_ _---,y,,-_ _--'J \~----=::::::=:::::;y;:::..--~ of an atom is empty space, Qi Qi Qi .c Ul because the electrons are quite -liS -liS for from the nucleus relative t o : ' : : ...J the size of the nucleus. If the nucleus of an atom were the size of an apple, the orbit of the nearest electron would be more than 1600 meters away. Consequently, the nuclei of two atoms never come

g,

•

£

28

Biomolecules

close enough in nature to interact with each other. It is for this reason that an atom's electrons, not its protons or neutrons, determine its chemical behaviour. This also explains why are the isotopes of an element, all of which have the same arrangement of electrons, because the same way chemically. Energy within the Atom Because electrons are attracted to the positively charged nucleus, it takes work to keep them in orbit, just as it takes work to hold an apple in your hand against the pull of gravity. The apple is said to posess potential energy; the ability to do work, because of its position; if you were to release it, the apple would fall and its ATOMIC STRUCTURE energy would be reduced. Conversely, if you were to move the apple to the top of a building, you would increase its potential energy. Similarly, electrons have potential energy of position. To oppose the attraction of the nucleus and move the electron to a more distant orbital requires an input of energy and results in an electron with greater potential energy. This is what happens during photosynthesis, when the energy in light excites electrons to higher energy levels. Moving an electron closer to the nucleus has the Carbon Atom (C) opposite effect; energy is released, usually as heat, and the electron ends up with less potential energy. During some chemical reactions, electrons are transferred from one atom to another. In such reactions, the loss of an electron is called oxidation, and the gain of an electrons is called reduction. It is important to realize that when an electron is transferred in this way, it keeps its energy of position. In organisms, chemical energy is stored in high-energy electrons that are transferred from one atom to another in reactions involving oxidation and reduction. Since the amount of energy possessed by an electron is related to its distance from the nucleus, electrons that are the same distance from the nucleus have the same energy, even if they are in different orbitals. Such electrons are said to occupy CH Methane ( 4) the same energy level. In a schematic diagram of an atom, the nucleus is represented as a small circle and he electron energy Fig. 3.5. Molecular structure of Carbon levels are drawn as concentric rings, with the energy level atom and Methane. increasing with distance from the nucleus. Be careful not to confuse energy levels, which are drawn as rings to indicate an electron's energy, with orbitals, which have a variety of three-dimensional shapes and indicate an electron's most likely location. Electrons orbit a nucleus in paths called orbitals. No orbital can contain more than two electrons, but many orbitals may be the same distance from the nucleus and, thus contain electrons of the same energy. Chemical Bonds: Linking Atoms Together A chemical bond is an attractive force that links two atoms to form a molecule. There are different kinds of chemical bonds (Table 3.2), but all strong chemical bonds results from an atom's tendency to attain stability by filling its outermost electron orbitals. Atoms can gain stability in the outermost orbitals by sharing electrons or by losing or gaining one or more electrons. In this section, we will first discuss covalent bonds, the strong·bonds that result from sharing of electrons. Then we'll examine hydrogen bonds, which are weaker than covalent bonds but enormously important to biology. Finally, we'll consider ionic bonding, which results when ions form as a consequence of the complete loss or gain of electrons by atoms.

29

Atoms and Molecules

Table 3.2. Chemical bonds. Types of Bond

Basis of Bonding

Energy

Bond Length

Covalent bond Ionic bond Hydrogen bond

Sharing of electron pairs Attraction of opposite charges Sharing of H atom

50-110 kcal/mol 3-7 kcal/mol 3-7 kcal/mol

0.28 nm (optimal) 0.26-0.31 nm

=1 kcal/mol

(between atoms that share H) 0.24-0.4 nm

van der Waals interaction

Interaction of' electron clouds

=0.1 nm

COVALENT BONDS CONSIST OF SHARED PAIR OF ELECTRONS

When two atoms attain stable electron numbers in their outer shells by sharing one or more pairs of electrons, a covalent bond forms. A hydrogen atom has one electron in its only shell, but two electrons wbuld be a more stable condition. Imagine two hydrogen atoms, initially far apart but coming closer and closer, until they b~gir, to interact. The negatively charged electron in each h~rogen atom is attracted by the positively charged proton in the nucleus of the other hydrogen atom. When the two atoms are close enough, the two electrons spend time between both I'luclei, and the two atoms are covalently bonded together, forming a molecule of hydrogen gas (H 2). The two hydrogen nuclei share the two electrons equally and completely. The two atoms do not come too close together, because their positively charged nuclei strongly repel each other. A certain distance between the coupled atoms gives the most stable arrangement. Pulling the atoms slightly farther apart would require an input of energy because of the "gluing" effect of the shared electrons. Pushing the atoms closer together would require energy because of the mutual repulsion of the protons. 50 the most stable arrangement of the covalently bonded hydrogens can also be described as an arrangement that has a minimum amount of energy and that is less reactive than are the individual atoms alone, each of which has an incompletely filled orbital in the K shell. A carbon atom has a total of six electrons; two electrons fill its inner shell and four are in its outer L shell. Because the L shell can hold up to eight electrons, this atom can share electrons with up to four other atoms. Thus it can form four covalent bonds. When an atom of carbon reacts wi!h.four hydrogen atoms, a substance called methane (CH4) forms, resulting from the overlap~~g of electron orbitals. Thanks to electron sharing, the outer shell of methane's carbon atom is filled with eight electrons, and the outer shell of each hydrogen 8tom is also filled. Thus four covalent bondseach bond consisting of a shared pair of electrons-hold methane together. Table 3.3 shows the covalent bonding capacities for some biologically Significant elements. Table 3.3. Typical bonding capabilities of some biologically important elements Elements

Hydrogen (H) Oxygen (0) Sulfur (5) Nitrogen (N) Carbon (C) Phosphorus (P)

Number of Covalent Bonds

1 2 2 3 4 5

Orientation of Bonds in Space Not only the number, but also the spatial orientation, of bonds is important. The four filled orbitals around the carbon nucleus of methane distribute themselves in space so that the bonded hydrogens are directed tothe corners of a regular tetrahedron with carbon in the center. Although the

30

Biomolecules

orientation of orbitals and shapes of molecLlles differ depending on the kinds of atoms and how they are linked together, it is essential to remamber that all molecules occupy space and have threedimensional shapes. The shapes of molecules contribute to their biological functions. Multiple Covalent Bonds A covalent bond is represented by a line between the chemical symbols for the atoms. Bonds in which a single pair of electrons is shared are called single bonds (for example, H-H, C-H). When four electrons (two pairs) are Propane (CH3CH2CH3) shared, the link is a double Fig. 3.6. Molecular structure of Propane. bond (C=q. In the gas ethylene (H 2C=CH2), two carbon atoms share two pairs of electrons. Triple bonds (six shared electrons) are rare, but there is one in nitrogen gas (N=N) , the chief component of the air we breathe. In the covalent bonds in these five examples, the electrons are shared more or less equally between the nuclei; consequently all regions of the bonds are identical. However, when electrons are shared unequally in a covalent bond, regions of partial electric charge exist. Unequal Sharing of Electrons So far we have discussed the covalent bonds that result from the equal sharing of electron:; between two nuclei. Now we want to consider the kind of covalent bond that results from unequal sharing of the electrons. Some atoms hold electrons to themselves more firmly than other atoms do. This characteristic is called electronegativity. Ethylene (CH2CH2) Highly electronegative atoms that form covalent bonds include Fig. 3.7. Molecular structure of Ethylene. oxygen and nitrogen. When these atoms are c,ovalently bonded to atoms with weaker electronegativity, such as carbon and hydrogen, the bonding pair of electrons is unequally shared between the two atoms, and the result is a polar covalent bond. For example, when oxygen is bonded to hydrogen, the bonding electrons spend much more time near the oxygen nucleus than near the hydrogen nucleus. Consequently, the oxygen end of the bond is slightly negative (symbolized &-- and spoken as "delta negative," meaning a partial unit charge), and the hydrogen end is slightly positive (0+). The bond is polar because these opposite charges are separated at the two ends of the bond. The partial charges that result from polar covalent bonds produce polar molecules or polar regions of large molecules. Polar bonds greatly influence the interactions between molecules that contain them, as we see in the interaction of water molecules in the liquid state. HYDROGEN BONDS

Both covalent and ionic bonds are regarded as rather powerful forces, but there are a number of other, weaker types of bonds as well. Perhaps the most important of these is called the hydrogen bond (a weak bond between the positive part of one molecule and the negative part of another molecule). Strangely enough, hydrogen bonds are important to life precisely because of their weakness. It takes very little energy to form them, and they are easily broken.

Atoms and Molecules

31

A hydrogen bond is fonned when two electronegative (negatively charged) atoms fulfill their shell requirements by sharing a single hydrogen atom. As an example water molecules may be held together by hydrogen bonds. The bonds NH are fonned because as we have AMMONIA ( 3) r-----------------~ seen, the electrons are more A B Nitrogen Relatively strongly attracted to the oxygen Negative (-I than to the hydrogens. This · means that the oxygen atoms, with their electrons drawn in close about them, become somewhat electronegative. This slight displacement leaves the (+1 hydrogen atoms (the protons Relatively from which the electrons were Positive Hydrogen drawn) with a slight positive charge. They therefore have a Fig. 3.B. Structure of an ammonia molecule. tendency to attract negatively charged particles, and the nearest such particle is likely to be the oxygen atom of the next water molecule. Thus, water molecules tend to join in a loose, but highly structured and constantly changing latticework. Hydrogen bonds last for only a very brief time-about 10 raised to the negative 11th power (1/ 100,000,000,000) second. But in their short existence, they bestow upon water unusual qualities of ..... __..... __ ~ __ .... being very fluid and, at the same time, relatively stable. The stability • • • • of water is quite easily demonstrated by considering the time it takes : : : l for it to boil. (Legend has it that if the pot is watched, the data may .-.--.-.--...... --..... be unreliable.) Why does water so strongly resist being changed to a • • • • gas? It is resistant because although the heat easily ruptures the bonds ~ ~ ~ between the molecules, they fonn new bonds with dazzling speed. • . •• The peculiarities of hydrogen bonding also explain why it is so difficult to freeze water. Ice is crystalline; that is, it is regular and Fig. 3.9. The hydrogen bonds here repeating in its molecular structure. But the constantly shifting are the small black dashes. molecules of water don't hold still long enough to encourage such a regular structure. The molecules are joined so weakly that they continually break and fonn new bonds. Water, then, is quite stable, a trait that has given it special importance in the processes of life. But, you may ask, just why is .this sort of stability so important to life? Actually, there are a number of reasons. For humans and other terrestrial air-breathers, for example, water's stabilitY, or resistance to change, enables us to breathe in a wide range of temperatures and in very dry environments by retarding the evaporation of moisture from our lungs. In addition, we are resistant to freezing because our bodies contain so much water. (The formation of ice crystals, of · course, can rupture delicate cell membranes.) Furthennore, the unequal charge distribution of water molecules, which pennits hydrogen bonding in the first place, also makes water a powerful solvent. Water can thus break down complex molecules, allowing their components to interact in new ways. In fact, all of the biochemical reactions in our bodies take place in water. Water's constantly changing structure also gives it a certain fluidity and movability that enab.1es it to pass through our bodies' tissues and to seep deep into the earth's crust to reach the roots of the largest trees. Furthennore, because of water's powerful tendency to join, it can fonn "columns" and move from those roots to the highest leaves as these "columns" are drawn along as units. It should be add that hydrogen bonds are important to a number of critical biological molecules. For example, they hold the two long chromosomal molecules together, as we will see, and they are important in maintaining the configuration of other large molecules.

Biomolecules

van der W~s Forus We have now looked at ionic bonds, covalent bonds and the weak hydrogen bonds. A fourth chemical bonding force is weaker yet: uan der Waals forces are weak and unspecific forces between non polar molecules. Such non·polar molecules have evenly distributed electrical charges (and so they lack positively and negatively charged areas). BecauSE! of complex molecular interactions when these non-polar molecules draw near each other, they generate slight positive and negative areas over each other's surfaces, causing them to draw together as closely as possible; as long as their outer electron shells do not overlap. The result is as if the non-polar molecules were ever-so-slightly sticky. It is because of van der Waal's forces that fat droplet mixed with water forms droplets. The fat droplets are non-polar and tend to cling together preventing any association with the polar water molecules. As we will see later, side groups of large molecules, such as proteins Polar co.;t~lent bond can be non-polar and thus van der Waals forces, acting with other bonding forces, can help hold these molecules in the complex and specific figurations. We will also see that van der Waal's forces are critical to H the structure and function of cell membranes. : Hydrogen bond

9-° :

forms between

Ions Form Bonds by Electrical Attraction • molecules. When one interacting atom--is much more electronegative than the other, a complete transfer of one or more electrons may take place. ~ For example, a sodium atom has only one electron in its outermost shell; this condition is unstable. A chlorine atom has seven electrons in Fig. 3.10. Hydrogen ' bonc:iing, its outer shell, another unstable condition. The reaction between sodium and chlorine makes both atoms more stable. When the two atoms meet, the highly electrqnegative chlQrine atom takes the single unstable electron from the SQdium atom. The result is two electrically charged particles, called ions. Ions are electrically charged particle& that form when atoms gain or lose one or more electrons. tru /----;~ ,.:'~ The sodium ion (Na+) has a +1 unit charge ~ ~ because it has one less electron than it has protons. Th~ outet'lllQst electron shell of the sodium Ion is . . Sodium atom (Na) Chlorine atrom (CI) full, with eight electrons, so the ion is stable. The (11 protons, 11 electrons) (17 protons, 17 electrons chloride ion (Cl-) has a ~ 1 unit charge because it has one mor~ electron than it has protons. This additional electron gives CI- an outer shell with a stable load of eight electrons. Negatively charged ions are caUed anions; positively charged ions are called cationl. Some elements form ions with multiple charges Sodium ion (Na+) by losing or gaining more than one electron to (11 protons. 10 electrons) achieve a stable electron configuration in their outer shell. Examples are Ca2+ (the calcium ion, created Fig. 3.11. Formation of sodium and chloride Ions. from a calcium atom that has lost two electrons), Mg2+ (magnesium ion), and AP+ (aluminum ion) . Two biologically important elements each yield more than one stable ion: Iron yields Fe2+ (ferrous jon) and Fe3+ (ferric ion), and copper yields Cu+ (cuprous ion) and Cu2+ (cupric ion). Groups of covalently bonded atoms that carry an electric charge are called complex ions; examples include NH/ (ammonium ion), 50/- (sulfate ion) and PO/- (Phosphate ion). Once they form, ions c;lre usually stable, and no more electrons are lost or gained. As stable entities, ions can enter into stable associations through ionic bonding. Thus stable solids such as sodium chloride (NaCl) and potassium phosphate (~P04) are formed. Although a very complex solid, bone has as one of its major components the simple ionic compound Ca3(PO4)2'

A...

t)

33

Atoms and Molecules

Ionic bonds: Electrical attractions Ionic bonds are the bonds fonned by electrical attractions between ions bearing opposite charges. In solids such as table salt (NaCl), the cations and anions are held together by ionic bonds. In solids, the ionic bonds are strong because the ions are close together. However, when ions are dispersed in water, the distance between them can be large; the strength of their attraction is thus greatly reduced. Under the conditions that exist in the cell, an ionic bond is less than one-tenth as strong as a covalent bond that shares electrons equally, so an ionic bond can be broken much more readily than a covalent bond. Not surprisingly, ions with one or more unit charges can interact with polar substances as well as with other ions. Such interaction results when table salt or any other ionic solid dissolves in water. The hydrogen bond that we described earlier is a weak type of ionic bond, because it is fonned by electrical attractions. However, it is weaker than most ionic bonds because the hydrogen bond is fonned by partial charges (0- and 0+) rather than by whole unit charges (+1 unit, -1 unit). Water molecules

,, ,, , , ."

_---.-::r

Chloride ion (CI)

/ ,,

:

Anion H attracts + pole of water

... .. ..

.. ... 1-.. . . . . . . . . . . . . .

I.._ _~

Cation (+) attracts - pole of water

Undissolved sodium chloride

Fig. 3.12. Water molecules surround ions.

Nonpolar Substances have no Attraction for Polar Substances We have been discussing the bonds that result from electrical attractions between positive and negative charges (ionic bonds and hydrogen bonds). Now let's return to a brief consideration of substances that have "pure" covalent bonds. These bonds form between atoms that have equal or nearly equal electronegativities-such as carbon and hydrogen-which share the bonding electrons equally. Such bonds are abundant in the compounds of hydrogen and carbon-the hydrocarbons. Molecules such as ethane (CH 3-CH3) and butane (CH 3-CH z-CH z-CH3) are small h~rocarbons, but in living systems, moiecules exist with h~rocarbon chains consisting of 16 or more carbon atoms.

Attraction between nonpolar molecules Nonpolar substances such as oils and fats show van der Waals attractions between molecules. These attractive forces operate only when nonpolar substances come very close to each other. The

34

Biomolecules

random variations in the electron distribution in one molecule create an opposite charge distribution in one molecule create an opposite charge distribution in the the adjacent molecule, and the result is . a brief, weak attraction. Although each such interaction is brief and weak at anyone site, the summation of many such interactions over the entire span of a nonpolar molecule can produce substantial attraction. Thus van der Waals interactions are important in holding together the long hydrocarbon chains that make up the inner portion of biological membranes. They also stabilize portions of the DNA double helix and the intricate folded structure of proteins.

Polar and nonpolar interactions When electrons are shared equally, the resultant covalent bonds are nonpolar and they do not interact with the charges of polar covalent bonds. Substances with only nonpolar bonds, such as butane and oils, will not interact with substances that have polar bonds or ionic bonds. This explains why oils will not dissolve in water. Oils are nonpolar hydrocarbons, while water is a highly polar substance. Substances with nonpolar covalent bonds are said to be hydrophobic (literally "waterfearing"), which refers to the fact that there is no attraction between nonpolar substances and water or other electrically charged substances. Hydrophobic substances are also called nonpolar substances. When hydrocarbons are dispersed in water, they slowly come together--dispersed molecules form droplets that form larger droplets. The forces that bring about this combining of molecules are sometimes called hydrophobic interactions, but this term is somewhat misleading. The interactions that combine nonpolar substances have less to do with forces between the nonpolar molecules than with the hydrogen bonding of the water that surrounds the molecules. Wh~n nonpolar substances such as hydrocarbons are introduced into water, they cause a disruption in the usual hydrogen bonding between water molecules. In the vicinity of the hydrocarbon, the water molecules form a hydrogen-bonded "cage" that surrounds the ·nonpolar hydrocarbons and pushes them together. These water cages can bring together dispersed nonPolar molecules into larger groups.

Significance of Polar Covalent and Hydrogen Bonds in Biological Systems Polar covalent bonds are the basis of hydrogen bonds which are perhaps the most important stabilizers of secondary and,tertiary structures of biological molecules. Polar bonds as such are responsible for polarity in some molecules, which are functionally important biomolecules. Some specific roles of these bonds in biological molecules are as follows: (i) Water as an unique solvent Water is a polar molecule because of the presence of polar covalent bonds in it. As such it can form hydrogen bonds with other molecules and can dissolve them. Other biologically important properties of water such as its fluidity, . formation of crystals, conductivity, physical constants etc. are also ·because of the potential to form polar and hydrogen bonds.

(il) Hydrogen bonds ·in nucleic acids The secondary structure of a double helical DNA molecule is stabilised due to the number of hydroen bonds. The two oppositely placed nucleotides of the two strands are linked through hydrogen bonds. Two hydrogen bonds are formed between adenine and thymine nucleotides and three between guanine and cytosine; . A=T G=C In secondary structure of RNA also hydrogen bonds are involved. The clover leaf model structure of transfer RNA is stabilised because of the hydrogen bonds in the double stranded regions of the RNA molecule. (ili) Hydrogen bonds in proteins Hydrogen bonds are also involved in the ·formation of secondary and sometimes in tertiary structure of proteins. In alpha helix and in beta sheet of the secondary structure of proteins, hydrogen bonds are the main stabilizing forces.

Atoms and Molecules

35

(iv) Enzyme-substrate complex formation In some cases the intermediate complex formed during enzymic conversion of substrate to the product is stabilized among others through the hydrogen bonds. Eggs by the Dozen, Molecules by the Mole In modem biology, as in chemistry, the question "How much?" is as important as "What kind?" In this section, we will examine .briefly how chemists and biologists deal quantitatively with atoms and molecules. The molecular formula uses chemical symbols to identify different atoms, and subscript numbers to show how many atoms are present. For example, the molecular formula for methane is CH4 (each molecule contains one carbon atom and four hydrogen atoms), that for oxygen gas is 02' and that for sucrose (table sugar) is C12H220U' The hormone insulin is represented by the molecular formula C254H377N6S076S6! Although molecular formulas tell us what kinds of atoms and how many of each kind are present in the molecule, they tell us nothing about which atoms are linked to which. Structural formulas give us this information. Each compoood has a molecular weight (molecular mass); the sum of the atomic weights of the atoms in the molecule. The atomic weights of hydrogen, carbon', and oxygen are 1.008, 12.011, and 16.000, respectively. Thus the molecular weight of water (HP) is (2 x 1.008) + 16.000 = 18.016, or about 18. What is the molecular weight of sucrose (C12H22011)? Your calculations should tell you that the answer is approximately 342. If you remember the molecular weights of a few representative biological compounds, you will be able to picture the relative sizes of molecules that interact with one another. Experiments require quantitative information. Suppose we want to compare how sodium chloride (NaCl), potassium chloride (KCl) , and lithium chloride (LiCI) affect a biological process. At first you might think we could simply give, say, 2 grams (g) of NaCI to one set of subjects, 2 g of KCl to another, and 2 g of LiCI to the third. But because the molecular weights of NaCi, KCl, and LiCI are different, 2-gram samples of each of these substances contain different numbers of molecules. The comparison would thus nto be legitimate. Instead, we want to give equal numbers of molecules of each substance so that we can compare the activity of one molecule of one substance with that of one molecule of another. How can we measure out equal numbers of molecules? Molecular Weight and t~e Mole The weight of a molecule is simply the sum of the weights of all the atoms of which that molecule is composed. The unit of weight is, again, the dalton (one-twelfth the weight of an atom. .of 12C). Thus the molecular weight (MW) of water is 18 daltons. (We will ignore the tiny error introduced by the presence of traces of 170, 180, and 2H atoms among the predominant 1H and 160 atoms.) Why is it important to know the molecular weight of a compound? For one illustration, let us asswne that you wish to study the response of honebees to solutions of various kinds of sugars. One way to do this would be to make up several different solutions and see which one or ones the bees prefer to harvest. You might offer the bess the choice between, say, a 35% solution of sucrose (common table sugar) and a 35% solution of glucose (a natural component of honey). This would involve, in each case, dissolving 350 parts by weight (e.g., grams) of sugar in 650 parts (g) of \ALat~ thus producing 1000 g of each solution. But there is a problem with this approach. The:wtUmgness of the honeybee to respond to the presence of sugar dissolved in water is dependent on the number of sugar molecules in a given volume of the solution. And the sucrose molecule (MW=342) is almost twice as heavy as the glucose molecule (MW=180). So a 35% solution of glucose would contain almost twice as many molecules as a 35% solution of sucrose. To correct the problem, you should make your solution with the weights of sucrose to glucose in a ratio of 342: 180: Then you would have the same concentration of molecules in each-that is, drop for drop each solution would contain the same number of molecules.

36

Biomolecules

If you weigh out exactly 342 grams (g) of sucrose you will have weighed out 1 mole of it. That is, mole is the quantity of a substance whose weight in grams is numerically equal to the molewlar weight of that substance. Thus 1 mole of glucose weighs 180 g. Furthermore, if you dissolve 1 mole of a substance in enough water to make 1 liter (L) of solution, you have made a 1-molar (l-M) solution. (A specially graduated flask, is a useful device in which to do this. After the substance is put into the flask, enough water is added to bring the volume up to the etched line on the neck of the ·flask. The volume of this solution is then exactly 1 liter.) A 1-M solution of these sugar would probably be too strong for your experiment with the bees. It might be better to make up a liter of each solution containing 34.2 g and 18.0 g, respectively. Such solutions would be designated one-tenth molar (O.l-M) solutions. Drop for drop, these two solutions would still contain exactly the same number of molecules because they are of the same molarity. How many molecules are there in a mole? The number, as you might expect, is very large: It is approximately 6 x 1023 . (This number is called Avogadro's number, after the chemist ·who first attempted to determine it.) Note that this number applies to a mole of any substance, molecule or ion. Thus we can properly refer to a mole of hydrogen ion (1 g).

pH Most of the chemical reactions that occur in living matter are extremely sensitive to the degree of acidity in which they take place. Degree of acidity is measured on a scale of pH units. pH is defined as the negative logarithm of the concentration of hydrogen ions (H+) measured in moles per liter (M). Thus a O.l-M solution of hydrogen ions has a pH of 1. Changing 0.1 to scientific notation gives us 1 x 10- 1 . The negative of the exponent is 1, so the pH is 1. A solution containing 0.01 moles/liter (0.01 M) is 1 x 10-2M and thus has a pH of 2. It is 10 times less acidic than the first solution. Water molecules themselves break up into ions to a slight extent. A liter of pure water contains 10-7 moles of hydrogen ions (in the form of the hydronium ion H30+) and thus a pH of 7. But pure water is neutral-that is, neither acid nor basic. So any solution with pH of 7 is neutral. Basic (alkaline) substances are those, like ammonia, that combine with hydrogen ions, thus leaving even fewer than are found in pure water. Therefore, they form solutions with pH values greater than 7.

Chemical Changes We mentioned earlier that the relative proportions by weight of oxygen and hydrogen in water are 16:2. This is a reflection of the ratio of the numbers of atoms in the molecule: one atom of (At. Wt. = 16) to two atoms of hydrogen (At. Wt. = 1). If we pass an electric current through water containing enough ions to render it a good electrical conductor, we can decompose the water into constituents: hydrogen and oxygen. This is a chemical change. There has been a rearrangement of atoms. Or, to put it another way, certain bonds are broken and new bonds are formed. This chemical change can be expressed by an equation. On the left of the equation we write the molecular formulas of all the ingredients or reactants. In this case there is only one, H 20. We then draw an arrow and write in the formulas of all the products. Both oxygen and hydrogen when present as elements exist as diatomic molecules. Thus we indicate them by the formulas H2 and 02' Now our equation appears H 20 --? H2 + 02' But something is wrong. Matter cannot be created or destroyed. Every atom that participates in a chemical reaction must be accounted for in the products and vice versa. So we must use two molecules of water to get one molecule of 02' And the decomposition of two molecules will yield two molecules of H2. So our final, balanced equation reads 2Hp --? 2H2 + 02 Note that the prefixes in our equation tell us the relative number of atoms and molecules involved. If we are interested in the proportions by weight, we must multiply by the molecular weights. By doing so, we see that from 36 weight units of water we get 4 parts of hydrogen and 32 parts of oxygen by weight. If our unit of weight is the gram, we have decomposed 2 moles of water and

°

37

Atoms and Molecules

produced 2 moles of hydrogen and 1 mole of oxygen. A balanced equation can thus tell us not only the nature of the reactants and products in a chemical reactio but also the proportions in which the reactants are consumed and the products produced. And note too, that every gram of matter that is used as a reactant reappears' in the products. BoND ENERGY

To carry out the decomposition of water, an electric current must be passed through the water. Electricity is one form of energy. Energy is the capacity to do work. The decomposition of water requires the expenditure of energy. The reason for this is that it takes energy to break chemical bonds. However, energy is released when chemical bonds form. In fact, for any particular bond, say the covalent bond between hydrogen and oxygen, the amount of energy it takes to break that bond is exactly the same as the amount liberated when the bond is formed. This value is called the bond energy.

There are many forms of energy, for example, electrical, mechanical, and chemical. But all forms of energy are ultimately converted into heat. Therefore, it is convenient, at least for biologists, to measure energy in terms of units of heat. The unit we shall use most frequently is the kilocaloric (kcal). A kilocalorie is the amount of heat needed to warm 1 liter of water 1 Celsius degree (DC). The kilocalorie is also the unit used to described the energy content of foods. It is the "Calorie" so familiar to diet-conscious people. Returning to our example, we find that it takes 118 kcal of energy to decompose 2 moles of water into its elements. Actually, it takes more than 118 kcal to decompose a mole of water into its atoms, but as these atoms immediately bond together to form molecules of hydrogen and oxygen, some of that energy is given back. The bond energy of the H-G bond is 110 kcal. The bond energy of H-H bonds is 103 kcal, and that of 0=0 bonds is 116 kcal. The decomposition of two molecules of water requires the breaking of four H-G bonds and thus the input of 440 kcal. The formation of 2 moles of hydrogen yields 206 kcal (2 x 103), and the formation of a mole of oxygen yields 116 kcal. Substracting the sum of the energy released (206 + 116 = 322 kcal) from the energy consumed (440 kcal) gives us the net energy input of 118 kcal that we measured previously. Where has the energy gone? In a real sense, it isnow chemical energy stored in the bonds of the hydrogen and oxygen molecules. The energy stored in this reaction is also called free energy because, as we shall see shortly, it is still available to do work. It is useful to have a symbol for free energy, and we shall use the letter G (in honour of Josiah Willard Gibbs, who developed the concept of free energy). What is free energy? It is energy that can be harnessed to do work. The water stored behind a dam has free energy. When allowed to fall through turbines, it can generate electricity (another form of free energy). But for biologists, the most interesting form of free energy is the energy stored in chemical bonds. It, too, can be harnessed to do work. When we lift a weight, we are using the free energy stored in the bonds of food molecules to run a machine, our muscles, which do the work of lifting the weight. However, the coversion of free energy to work is not 100% efficient. As we contract our muscles, much of the free energy of our food is given off as heat. This energy is no longer free; there is no way we can harness the warmth of our muscles to accomplish biologically useful work. We shall indicate a change in free energy by the symbol LlG. In the decomposition of water, the energy needed is not lost but is stored in the products. By convention we indicate the storage of free energy with a plus sign. Therefore, the reaction is expressed as 2Hp ~ 2H z + 0z' LlG = + 118 kcal. You may have seen a chemistry professor ignite a mixture of hydrogen and oxygen. If not, simply accept my word that the result is a dramatic explosion. The equation for this chemical reaction is the reverse of the one we have just been studying and is expressed as

Biomolecules

38

2H2 + O 2 ~ 2Hp. And, as the explosion suggests, this time a release of energy occurs. In fact, the free energy change is once again 118 kcal. This is because it took only 322 kcal to break the H-H and 0 = 0 bonds, and 440 kcal were liberated by the 4 moles of H-O bonds that were formed. We express the fact that energy has come out of the reacting system by putting a minus sign before L\G. Therefore, the reaction is expressed as 2H2 + O2 ~ 2HP, L\G = - 118 kcal. It turns out that mitochondria exploit a similar reaction to secure the free energy they need. Mitochondria synthesize water using the hydrogen atoms removed from organic molecules, like glucose, and the oxygen atoms they take in as they respire. The process is called cellular respiration. The overall equation for the respiration of glucose is C6H 1P6 + 60 2 ~ 6C0 2 + 6Hp; L\G = - 686 kcal. The same equation describes the burning of glucose, and the same amount of free energy is released. But the energy of burning is released as heat, which is of little value to cells. The achievement of mitochondria is their ability to release the energy of glucose in small, descrete steps so that some of the energy can be trapped in ATP. OXIDATION-REDUCTION REACTIONS

Most of the chemical reactions that we have been studying involve shifts in the location of electrons. In the reaction of chlorine with sodium, for example, each sodium atom loses an electron to a chlorine atom. In such a case, the atom that loses the electron is said to have been oxidized. The atom that gains the electron is said to have been reduced. The two processes are inextricably linked. Whenever any substance gives up electrons, it is oxidized. The substance to which it gives its electrons is called the oxidizing agent. Note, however, that any substance acting as an oxidizing agent is itself reduced. Because oxidations always go hand in hand with reductions, we refer to these reactions as redox reactions. You may wonder why the reaction between sodium and chlorine is called an oxidation when oxygen has no part in it. The answer is simply a matter of history. Oxygen has a great affinity for electrons (it is the second most electronegative element) and is an excellent (as well as common) oxidizing agent. Because of its effectiveness in this role, it has supplied the name for all reactions of this type. (Actually, fluorine is even more electronegative than oxygen, and perhaps the process should rightly be called fluoridation. But that term has come to be used for quite a different process: the adding of trace amounts of the fluoride ion, F-, to municipal water supplies.) Some biological oxidations do indeed involve the addition of oxygen to the molecule being oxidized. When ethanol (the common beverage alcohol) is exposed to air and the action of vinegar bacteria, it is converted into acetic acid and water (vinegar). The equation for this reaction is C2HsOH + O 2 ~ CH3 COOH + Hp. The same equation, using electron-dot formulas. Note the atom of oxygen that has been added to the ethanol to form acetic acid. Because the difference in electronegativity between carbon and oxygen atoms is not as great as, say, that between oxygen and magnesium, the added oxygen atom does not remove electrons sufficiently far to form ions. It nevertheless does attract two electrons away from the carbon atom and much closer to itself. The most common kind of biological oxidation is accomplished by the removal of hydrogen atoms from a substance. The oxidation of ethanol also exhibits this kind of oxidation. Two hydrogen atoms each with its single electrons, are removed by the second oxygen atom. These unite to form a molecule of water. Although all oxidations go hand in hand with reductions, we often emphasize the oxidation aspect of the reactions when electrons are moving from a weakly electronegative atom to a strongly electronegative one. The reactions of sodium and chlorine, hydrogen and oxygen, ethanol and oxygen are all examples. In all of these cases, free energy is liberated. This is because weaker

Atoms and Molecules

39

bonds have been broken and strong bonds formed. We exploit oxidations characterized by a large -. ilG to heat our homes, to run our automobil~s, and-of special interest to us-to run all the metabolic activities of life. It is also possible to move electrons back from strongly electronegative atoms to weakly electronegative ones. In this case ilG is positive-that is, energy must be put into the reaction. We usually emphasize the reduction aspect of such redox reactions. To take an earlier example, decomposing 2 moles of water requires the input of 118 kcal of energy. We have oxidized oxygen atoms (they lose electrons) and reduced the hydrogen atoms. We have, in fact, broken strong bonds (H-O) and formed weaker bonds (H-H, 0-0). Thus weak bonds can serve as a means of storing energy. This energy is chemical, potential energy. As soon as we permit the electrons that are involved to return to more strongly electronegative atoms, the energy is liberated. The life continues to exist on this earth because it is able to exploit redox reactions. It uses them to store the sun's energy and then to liberate the resulting chemical energy as needed to run the activities characteristic of living things. In photosynthesis the energy of the sun is used to move electrons (associated with protons, thus hydrogen atoms) from strongly electronegative atoms (oxygen) to weakly electronegative ones (carbon). In this process strong bonds (H-O, C-O) are broken and weaker bonds are formed (C-H and especially 0-0). The reverse of this process is cellular respiration. In this case the electrons return to their original partners, and the energy that has been stored is liberated and used to run the organism. Thus the energy differences among chemical bonds make it possible to operate a kind of battery: storing the sun's energy and liberating it as needed to maintain the living process.

4 WATER Because it is ubiquitous, water is often regarded as a bland, inert liquid, a mere space filler in living organisms. Actually, water is a highly reactive substance with unusual properties, and it is very different, both chemically and physically, from most other common liquids. The first cells to arise in the primordial sea had to learn to cope with the singular properties of water, and ultimately, as ws shall see, living organisms evolved means of exploiting these properties. We now recognize that water and its ionization products, hydrogen and hydroxyl ions, are important factors in determining the structure and biological properties of proteins, nucleic acids, lipids, membranes, and many other cell components. Physical Properties and Structure of Water When compared with other common liqUids, water has a high melting point, boiling point, heat of vaporization, specific heat, heat of fusion, and surface tension. These properties indicate that the forces of attraction between ·molecules in liqUid water, and thus its internal coheSion, are relatively high. For example, in Table 4.1 we see that the heat of vaporization of water is considerably higher than that of any of the other common liquids listed. The heat of varporization is a direct measure of the amount of energy required to overcome the attractive forces between adjacent molecules in a liquid so that individual molecules can escape from each other and enter the gaseous state. Table 4.1. Heat of vaporization of some common liquids at their boiling point (1.0 atm) Water Methanol Ethanol n-Propanol Acetone Benzene Chloroform

LlHvap cal gm- 1 540 263 204 164 125 94 59

The strong intermolecular forces in liquid water are caused by the electrical polarity of the water molecule, which in turn is the consequence of the specific arrangement of electrons in its oxygen and hydrogen atoms. The oxygen atom shares a pair of electrons with each of the two hydrogen atoms, through overlap of the Is orbitals of the hydrogen atoms with two hybridized Sp3 orbitals of the oxygen atom. Each of these electron pairs has about one-third ionic and two-thirds covalent character. From spectroscopic and x-ray analyses, the precise bond angles and lengths have been determined. The average H-G-H bond angle in water is 104.5°, which represents a slight deviation from a 40

Water

41

perectly tetrahedral arrangement of the four possible Sp3 orbitals of the oxygen atom, which would have an angle of 109.5°. One explanation for this deviation is that the unpaired electrons of oxygen tend to repel the paired electrons. The average H-O interatomic distance of 0.965A. This arrangment of electrons in the water molecule gives it electrical asymmetry. The more electronegative oxygen atom tends to attract the single electrons of the hydrogen atoms, leaving the hydrogen nuclei bare. As a result, each of the two hydrogen atoms has a local partial positive charge (designated a+). The oxyge atom, in tum, has a local partial negative charge (designated a -) located in the zone of the unshared orbitals. Thus, although the water molecule has no net charge, it is an electrical dipole. The degree of separation of positive and negative charges in dipolar molecules is given by the dipole moment, a measure of the tendency of a molecule to orient itself in an electrical field. From the dipole moment of water, it o has been calculated that each hydrogen atom has a partial 10 positive charge of about +0.33 x 10- electrostatic unit (esu) and the oxygen atom a partial negative charge of about -0.66 x 10-10 esu. Hydrogen Bonding (a) (b) The dipolar nature of the water molecule is largely FIg. 4.1. (a) Structure of water molecule. (b) responsible for the attractive forces between water Tetrahydral hydrogen bonding of water molecules, since a strong electrostatic attraction occurs molecules in ice. • between the partial negative charge on the oxygen atom of one water molecule and the partial positive charge on the hydrogen atom of an adjacent water molecule. This type of electrostatiC interaction, which has a small covalent component, is called hydrogen bonding. Because of the nearly tetrahedral arrangement of the electrons about the oxygen atom, each water molecule tends to hydrogen-bond with four neighbouring water molecules. An important property of hydrogen bonds is that they are much weaker than covalent bonds. The H bonds in liquid water are estimated to have a bond energy of only about 4.5 kcal mole-I, compared with 110 kcal mole- l for the H-O electron-pair bonds in waer. (Note: Bond energy is the energy required to break a bond.) Another important property of hydrogen bonds is that they have a high degree of directionality, which is conferred by the characeristic arrangement of the bonding orbitals of the hydrogen and oxygen atoms. Hydrogen bonds also possess a specific bond length, which differs from one type of H bond to another, depending on the structural geometry and the electron distribution in the molecules involved. In ice, each water molecule is hydrogen-bonded; the length of the hydrogen bond is 1. 77 A. Hydrogen bonds therefore form and remain stable only tmder speCific geometrical conditions. Hydrogen bonding between water molecules occurs not only in liquid water but also in ice and in water vapour. In the most common crystalline form of ice, called ice I, each water molecule is hydrogen-bonded with exactly four nearest Fig. 4.2. The open structure of ice. neighbours and the average oxygen-oxygen distance is 2.76A.

42

Biomolecules

In liquid water, each water molecule at 0° is hydrogen-bonded at any given time with an average of about 3.4 other water molecules; the average oxygenoxygen distance is only slightly greater than in ice, about 2.90A at 15°C and 3.osA at 83°C. It has been estimated from the heat of fusion of ice that only about 15 percent of the hydrogen bonds in ice are broken when ice is melted to water at OoC. Thus liquid water can be regarded as slightly broken-down ice, with considerable short-range molecular order but no long-range order. : Strong attractions between water molecule still exist in water at 100°C, as is indicated by its high heat of vaporization. In fact, hydrogen bonding between water molecules is completely overcome only when water vapour is heated to nearly 600°C. Hydrogen bonds are not unique to water. They form between a highly electronegative atom such as oxygen, nitrogen, or flUOrine, and a hydrogen atom covalently bonded to another such electronegative atom. Hydrogen bonds may form between two molecules or between two parts of the same molecule. Fig. 4.3. Directionality Single hydrogen bonds between two solute molecules in aqueous systems are of the hydrogen very weak because the surrounding water molecules compete to form hydrogen bond. bonds with the solutes. However, when a number of hydrogen bonds exist between two structures, the energy required to separate them is much greMer than the sum of the bond energies of the individual hydrogen bonds. This phenomenon is called cooperative hydrogen bonding, and it.is characteristically seen in proteins and some nucleic acid molecules, which may contain dozens or even hundreds of cooperating hydrogen bonds. Such bonding yields structures that are surprisingly stable in water. H R H R'C/ R'

/j\

I

II

o

C H

H

"

o

I 0,

I

'v"N, /C'-/

H

(a)

(b)

Between a hydroxyl group and H 20

Between a carbonyl group and H 20

C

I

N

C

I

I

Thymine

17C" /C~ N 0

/"..,H/N, /"-.

II

R

/C~ /CH3

o

H . I C

I

,

0 (c)

Between two peptide chains

I H

H

I

H" /N~ /N-H C ~C Adenine

.11 N

I C

'C~ 'N

\

II

N-CH

/

(d) Between complementary base pairs in DNA

Fig. 4.4. Some hydrogen bonds of biological importance.

The Kinetics of Hydrogen Bonding We have seen that between ice and liqUid water there is only a smaIl difference in the amount of hydrogen bonding. This may appear surprising in view of the rigidity of ice and the fluidity of liquid water. The explanation lies in the rate at which hydrogen bonds are made and broken. Although at any given time most of the molecules in liquid water are hydrogen-bonded, the half-life of each hydrogen bond is only about 10-10 to 10- 11 seconds. The structure of liquid water is therefore only statistical, since it is the result of an average over both space and time. Consequently, liquid water is

43

Water

at once fluid and stable. The apt term "flickeming clusters" has been applied to the short-lived icelike groups of water molecules in liquid water. The precise structure of such clusters and the nutnber of hydrogen bonds per cluster are still unknown and are the subjects of much current research. The rate at which hydrogen bonds form and break'in aqueous systems greatly exceeds the rate of making and breaking of most covalent bonds. This fact endows hydrogen bonds with a great biological advantage over covalent bonds in certain types of biomolecular phenomena.

Water is the Principal Solvent in Organisms Because its molecules are polar, water is an excellent solvent, a liquid capable of dissolving many different kinds of substances, especially polar and ionic compounds. Previously in this chapter, we discussed how polar water molecules pull the ions of ionic compounds apart so that they dissociate. Because of its solvent properties and the tendency of the atoms in certain compounds to form ions when in solution, water plays an important role in facilitating chemical reactiol1s. Substances that interact readily with water are said to be hydrophilic ("water-loving"). Not all substances in organisms are hydrophilic, however. Many hvdrophobic ("water-hating") substances found in living things are especially important because of their ability to form structures that are not dissolved by water. Hydrogen-Bonding makes Water Cohesive and Adhesive Water molecules have a very strong tendency to stick to each other, that is, they are cohesive. This is due to the hydrogen bonds among the molecules. Water molecules also stick to many other kinds of substances, most notably those with charged groups of atoms (M (Il) or molecules on their surfaces. These adhesive forces explain how water makes things wet. -,-A combination of adhesive and cohesive forces accounts for the tendency, termed capillary action, of water to move in narrow tubes, even against the force of gravity. For example, water moves through the microscopic space between soil particles to the roots of plants by capillary action. Because of the cohesive nature of water molecules, any + i force exerted on part of a column of water will be transmitted to the colun as a whole. The major mechanism of water movement in plants depends on this fact. Water has a high degree of surface tension because of the cohesiveness of its molecules, which have a much greater attraction for each other than for molecules in the air. Thus, water molecules at the surface crowd together, producing a strong layer as they are pulled Fig. 4.5. The cohesive and adhesive properties of downward by the attraction of other water molecules beneath them. water account for Water Clings to Polar Molecules capillary action. (a) In the smaller tube, The polarity 0 f water 'causes it to be attracted to other polar adhesive forces attract molecules. When the other molecules are also water, the attraction is water molecules to referred to as cohesion. When the other molecules are of a different changed groups on substance, the attraction is called adhesion. It is because water is cohesive the surfaces of the that it is a liquid, and not a gas, at moderate temperatures. tube. (b) In the largediameter tube, a The cohesion of liqUid water is also responsible for its surface smaller percentage of tension. Small insects can walk on water because at the air-water interace the water molecules all of the hydrogen bonds in water face downward, causing the molecules line the glass. of the water surface to cling together. Water is adhesive to any substance with which it can form hydrogen bonds. That is why substances containing polar molecules get "wet" when they are immersed in water, while those (such as waxes) that are composed of nonpolar molecules do not. The attraction of water to substances like glass with surface electrical charges is responsible for capillary action: if a glass tube with a narrow diameter is lowered into a beaker of water, the water

I

I

I

44

Biomolecules

will rise in the tube above the level of the water in the beaker, because the adhesion of water to the glass surface, drawing it upward, is stronger than the force of gravity, draWing it down. The narrower the tube, the greater the electrostatic forces between the water and the glass, and the higher the water rises. Water Helps Maintain a Stable Temperature Raising the temperature of a substance involves adding heat energy to make its molecules move faster-that is, to increase the kinetic energy (energy of motion) of the molecules. The term heat refers to the total amount of kinetic energy in a sample of a substance; temperature refers to the average kinetic energy of the particles. Water has a high specific heat; that is, the amount of energy required to raise the temperature of water is quite large. A calorie is a unit of heat energy (equivalent to /4.184 joules) that equals the amount of heat required to raise the temperature of 1 gram of water 1 degree Celsius. The specific heat of water is therefore 1 calorie per gram of water per degree Celsius. Most other common substances have much lower specific heat values. The high specific heat of water results from the hydrogen-bonding of its molecules. Some of the hydrogen bonds holding the water molecules together must first be broken to permit the molecules to move more freely. Much of the energy added to the system is used up in breaking the hydrogen bonds, and only a portion of the heat energy is available to speed the movement of the water molecules (thereby increasing the temperature of the water). Conversely, when liquid water changes to ice, additional hydrogen bonds must be formed, liberating a great deal of heat into the environment. Because so much heat input is required to raise the temperature of water (and so much heat is lost when the temperature is lowered), oceans and other large bodies of water have relatively constant temperatures. Thus, many organisms living in the oceans are proVided with a relatively constant environmental temperature. The properties of water are crucial in stabilizing temperatures on the surface of the Earth. Although surface water is only a thin film relative to Earth's volume, the quantity is enormous compared to the exposed land mass. This relatively large mass of water resists both the warming effect of heat and the cooling effect of low temperatures. In addition, hydrogenbonding gives ice unique properties that have important environmental consequences. The high water content of organisms helps them maintain relatively constant internal temperatures. Such minimizing of temperature fluctuations is important because biological reactions can take place only within a relatively narrow temperature range. Because its molecules are held together by hydrogen bonds, water has a high heat of vaporization. To change 1 gram of liquid water into 1 gram of water vapour, 540 calories of heat ar~ required. The heat of vaporization of most other common substances is much less. As a sample of watE!r is heated, some molecules are moving much faster than others (that is, they have more heat energy). These faster-moving molecules are more likely to escape the liqUid phase and enter the vapour phage. When they do, they take their heat energy with them (thus lowering the temperature of the sample). For this reason the human body can dissipate excess heat as sweat evaporates from the skin, and a leaf can keep cool in the bright sunlight as water evaporates from its surface. HYDROPHOBIC INTERACTIONS

Water also disperses many compounds containing non-polar or hydrophobic groups in the form of micelles providing such compounds also contain strongly polar groups. This type of "solubilization" is made possible by hydrogen bonding, not between solute and solvent molecules but among the solvent molecules. Compounds containing both highly hydrophobic and highly polar groups are called amphipathic. The most common amphipathic biomolecules that tend to form micelles are fatty acids and polar lipids. A simple example is the sodium salt of the long-chain fatty acid oleic acid. Because it long hydrocarbon chain is intrinSically insoluble in water, there is very little tendency for sodium oleate (a soap) to dissolve in water in the form of a truly molecular solution. However, it readily disperses in

45

Water

water to form micelles in which the negatively charged carboxyl groups are exposed to the water phase and the non-polar, insoluble hydrocarbon chains are hidden within the micellar structure. Such micelles have a net negative charge, Polar and they remain suspended because of mutual repulsion. Micelles may contain hundreds or even thousands of molecules of head amphipathic substance. The characteristic internal location of the nonpolar groups in micelles in the result of the tendency of water molecules to assume the state of maximum randomness, or entropy. In this state, water has enormous internal cohesion because of hydrogen bonding. To insert a nonpolar structure such as hydrocarbon chain into water would require energy since the surrounding water molecules would be forced into a more regular arrangement, one with less entropy or randomness than in pure water. Put in another way, micelles form because water likes water more than it likes nonpolar structures. Within micelles, there are additional attractive forces between adjacent hydrocarbon structures through van aer Waals interactions. It must be emphasized Nonthat there is no true stoichiometric bonding between hydrocarbon groups in a polar micelle. For this reason, the term hydrophobic interaction is more properly tail used than hydrophobic bond to refer to the association of the hydrophobic portions of amphipathic molecules. Compared with hydrogen bonds, hydrophobic interactions have relatively little directionality, but they do tend to produce systems of high stability. As we shall see later, many other cell components are amphipathic and tend to form structures in which the nonpolar, hydrophobic parts are hidden from water-in particular, the proteins and the nucleic acids. Effect of Solutes on Water Structure The presence of an ionic solute such as NaCl causes a distinct change in the structure of liquid water since each Na+ and Cl- ion is surrounded by a shell Fig. 4.6. Formation of a soap of water dipoles. These hydrated ions have a geometry somewhat different from micelle in the clusters of hydrogen-bonded water molecules; they are more highly ordered water. and regular in structure. Table 4.2 gives the average interionic distance in aqueous NaCI solutions as a function of hte concentration of NaCI. We see that at 0.15 M NaCl, the approximate concentration of NaCl in blood plasma (and of K+ salts in the cytoplasm of cells). Na+ and Cl- ions are separated by only about 19A, on the average. Since each hydrated Na+ and Cl- ion is 5 to 7A in diameter and a tetrahedral cluster of five water molecules is about 5A in diameter, it is clear that a considerable change must occur in the three-dimensional structure and properties of liquid water when NaCI is dissolved in it in a concentration approximating that occurring in biological fluids. Salts therefore "break" the structure of water. Table 4.2. Average interionic distance in solutions of NaCI Concentration (M)

0.001 0.010 0.10 0.150 1.00

Distance (A)

94 44 20 19 9.4

The effect of a solute on the solvent is manifest in an other set of properties, namely, the colligative properties of solutions, which are dependent on the number of solute particles per unti volume of solvent. Solutes produce such characteristic effects in the solvent as depression of the

Biomolecules

46

freezing point, elevation of the boiling point, and depression of the vapour pressure. They also endow a solution with the property of osmotic pressure. One gram molecular weight of an ideal nondissociating nonassociating solute dissolved in 1,000 grams of water at a pressure of 760 mm of mercury depressed the freezing point by 1.86°C and elevates the boiling point by 0.543°C. Such a solution also yields an osmotic pressure of 22.4 atmospheres in an appropriate apparatus. Since aqueous solutions usually deviate considerably from ideal behaviour, these relationships are quantitative only at infinite dilution, i.e. on extrapolation to zero concentration of solute. A second large class of substances readily dissolved by water includes nonionic but polar compounds such as sugars, simple alcohols, aldehydes, and ketones. Their solubility is due to the propensity of groups, such as the hydroxyl groups of water molecules to hydrogen~bond with polar functional t sugars and alcohols and the carbonyl oxygen atom of aldehydes and ketones. Solvent Properties of Water Water dissolves or disperses many substances because of its dipolar nature. It is a much better solvent than most common liquids. Many crystalline salts and other ionic compounds readily dissolve in water but are nearly insoluble in nonpolar liquids such as chloroform or benzene. Since the crystal lattice of salts, such as sodium chloride, is held together by very strong electrostatic attractions between alternating positive and negative ions, considerable energy is required to pull these ions away from each other. However, water dissolves crystalline sodium chloride because the strong electrostatic attraction between water dipoles and the Na+ and Cl- ions, greatly exceeds the tendency of Na+ and Cl- to attract each other. Ion solvation is also aided by the tendency of the solvent to oppose the electrostatic attraction between positive and negative ions. This is given by the dielectric constant D, which is defined by the relationship e 1e 2 F=-2Dr where F is the attractive force between two ions of opposite change, e1 and e2 are the charges on the ions, and r is the distance between. them. As you can see in Table 4.3. water has an extremely high dielectric constant and benzene a very low one. The attractive force between Na+ and Cl- ions at given distance, in water is only one-fortieth that in benzene, a factor which greatly favours hydration of the ions and dissolution of the crystal1attice. Table 4.3. Dielectric Constants of some liquids (20°C)

o Water Methanol Ethanol Acetone Benzene

Hexane

80 33 24 21.4 2.3 1.9

Water Ionizes The covalent bonds within a water molecule sometimes break spontaneously. In pure water at 25°C, only lout of every 550 million water molecules undergoes this process. When it happens, one of the protons (hydrogen atom nuclei) dissociates from the molecule. Because the dissociated proton lacks the negatively charged electron that it had been sharing in the covalent bond with oxygen, its own positive charge is not counterbalanced. Thus, a positively charged ion, H+, is produced, which is usually associ?lted With another water molecule to form a hydronium (H30+) ion. The rest of the dissociated! water molecule, which has retained the shared electron from the covalent bond, is

47

Water

negatively charged and forms a hydroxide ion (OH-). This process of spontaneous ion formation is called ionization: H+ + ~O hydrogen ion water hydroxide ion At 25°C, a liter of water contains 1/10,000,000 (or 107) mole of H+ ions. (A mole is • defined as the weight in grams ~. that corresponds to the summed atomic masses of all of the atoms in a molecule. In the case 100° C • ~ of H+, the atomic mass is 1, (1 '-'0 0 and a mole of H- ions would 0 weight 1 gram. One mole of any Q-o substance always contains 6.02 .0.,.-,6) x 10 23 molecules of the substance.) lberefore, the molar concentration of h~en ions (represented as IH+)) in pure water is 10-7 mole/liter. A more convenient way to express the hydrogen ion 50° C concentration of a solution is to use the pH scale in which pH is defined as the negative logarithm of the hydrogen ion concentration in the solution: pH = -log IH+) Since the logarithm of the hydrogen ion concentration is simply the exponent of the molar concentration of H+, the pH equals the exponent times -1 . Thus, pure water, w!th an [H+) · of 10-7 mole!liter, has a pH of 7. Recall that for every H+ ion formed when water dissociates, an OH- ion is also formed. In other words, the dissociation of water produces H+ and OH- in equal amounts. Therefore, a pH value of 7 indicates neutrality of the pH Fig. 4. 7. Water can exist as a gas (water vapour), a liquid, or a solid (ice). scale. Any substance that dissociates to form an excess of H+ ions when dissolved in water is called an acid. Acidic solutions have pH values below 7. The stronger an acid is, the more H+ ions it produces and the lower its pH. For example, hydrochloric acid (HC!) , which is abundant in your stomach, ionizes completely in water, so 10-1 mole per liter of HCl will dissociate to form 10-1 mole per liter of H+ ions, giving the solution a pH of 1. The pH of champange, which bubbles because of the carbonic acid dissolved in it, is about 2.

a

.1J c1

.'-.,.

.0

D

•

Biomolecules

48

A substance that combines with H+ ions when dissolved in water is called a base. By-cOfi1l)ining

with H+ ions, a base lowers the H+ concentration in the solution. Basic (or alkaline) solutions, therefore, have pH values above 7. Very strong bases, such as sodium hydroxide (NaOH), have pH values of 12 or more. ACIDS ~ PROTON DONOIlS; BASES ARE PROTON ACCEPTORS

Water molecules have a slight tendency to iQni~e-that is, to dissociate into hydrogen ions (H+) and hydroxide ions (OH-). In pure water, a very small number of water molecules ionize. This slight tendency of water to dissociate is reversible as hydrogen ions and hydroxide Ions reunite to form water:

Because each water molecule splits into one hydrogen ion and one hydroxide ion, the concentrations of hydrogen ions and hydroxide ions in pure water are exactly equal (0.0000001 or 10-7 moles per liter for each ion). Such a solution is said to be neutral, neither acidic nor basic (alkaline). An acid is a substance that dissociates in solution to yield hydrogen ions (H+) and an anion. Acid -7 H+ + anion An acid is a proton donor. (Recall that a hydrogen ion, or H+, is nothing more than a proton.) An acidic solution has a hydrogen ion concentration that is higher than its hydroxide ion concentration. Acidic solutions tum blue litmus paper red and have a sour taste. Hydrochloric acid (HCI) and sulfuric acid (H 2S04) are examples of inorganic acids. Lactic acid (CH3CHOHCOOH) from sour milk !U1d acetic acid (CH3COOH) from vinegar are two common organic acids. A base is defined as a proton acceptor. Most bases are substances that dissociate to yield a hydroxide ion (OH-) and a cation when dissolved in water. A hydroxide ion can act as a base by accepting a proton (H+) to form water. Sodium hydroxide (NaOH) is a common inor~anic base. . NaOH -7 Na+ + OHOH- + H+ -7 Hp Some .bases do not dissociate to yield hydroxide ions directly. For exampl~, ammonia (NHJ acts as bqse by accepting a proton from water, producing an ammonium ion (NH4+) and releasing a hydroxide ion. NH3 + Hp -7 NH/ + OHA basic solution is one in which the hydrogen ion concentration is IQwer than the hydre»dde ion concentration. Basic solutions tum red litmus paper blue and feel slippery to the touch. In later encounter a number of organic bases such as the purine and pyrimidine bases that are components of nucleic acids. pH is a Convenient Measure of Acidity The degree of a solution's acidity is generally expressed in tel111$ of pH, defined as the negatioe logarithm (base 10) 0/ the hydrogen ion concentration (expresseq in mole$ per liter). pH = - 10g10 IH+) The brackets refer to concentration; therefore the term IH+) means "the concentration of h~rogen ions," which is expressed in moles per liter bec~use we are interested in the number of hydrogen ions per liter. Because the range of possible pH values is very broad, a logarithmic scale (with a tenfold difference between successive units) is more convenient than a Unear scale. Hydrogen ion concentrations are nearly alwaYS less than l mole per liter. One gram of hydrogen ions dissolved in 1 liter of water (a 1-molar ·solution) may not sound very impressive, but such a solution woUld be extremely acidic. The logarithm of a number less than one is a negative number; thus the negative logarithm corresponds to a positive pH value.

49

Water Distilled Water Milk 6.6 Blood 7.4

I Egg white 8.0

Black coffee 5.0 Tomato 4.6 Vinegar 3.0 Stomach gastric juice 2.0

I I

Baking soda 9.0 Milk of magnesia 10.5 Household ammonia 11.0 13.0

Hydrochloric acid 0.8 -

Fig. 4.8. A pH meter is an electronic device used to measure the acidity

0/ a

solution.

Whole number pH values are easy to calculate (Table 4.4). For instance, consider our example of p~re watf2r, which has a hydrogen ion concentration of 0.0000001 (10-7) moles per liter. The logqrith11l is -7. The negative logarithm is 7; therefore the pH is 7 .

Table 4.4. The relation of pH to hydrogen ion concentration

Sl4bstance Gastric juice Pure water, neutral solution Household ammonia

{H+j

0.01, 100.0000001, 10-7 0.00000000001, 10- 11 2

log (H+j

pH

-2 -7 -11

2 7 11

If the hydrogen ion concentration of a solution is known, the hydroxide ion concentration can be easily calculated. The product of the hydrogen ion concentration and the hydroxide ion concentration is 1 x 10-14 .

[H+) [OH-) = 1 x 10- 14 In pure (freshly distilled) water, the hydrogen ion concentration is 10-7 ; therefore the hydroxide concentration is also 10-7 . Such a solution, in which the concentrations are equal, is said to be neutral. Acidic solutions (those with an excess of hydrogen ions) have pH values smaller than 7; basic solution (those with an excess of hydroxide ions) have pH values greater than 7. A more acidic solution has a lower pH value. The hydrogen ion concentration of a solution with pH 1 is ten times that of a solution wit~ pH 2. The contents of most animal and plant cells are neither strongly acidic nor strongly basic but instead are an essentially neutral mixture of qcidic and basic substances. Any substantial change in the pH of the cell is incompatible with life. The pH of living cells ordinarily ranges from around 7.2 to 7.4.

Buffers Minimize pH Change Many homeostatic mechanisms operate to maintain appropriate pH values. For example, the pH of human blood is about 7.4 and must be maintained within very narrow limits. Should the blood

50

Biomolecules

become too acidic (for example, as a result of respiratory disease), coma and death may result. Excessive alkalinity can result in over-excitability of the nervous system and even convulsions. Organisms contain many natural buffers. A buffer is a substance or combination of substances that resists changes in pH when an acid or base is added. A buffering system includes a weak acid or a weak base. A weak acid or weak base does not ionize completely. That is, at any given instant only a fraction of the molecules are ionized; most are undissociated. One of the most common buffering systems is important in human blood. Carbon dioxide, produced as a waste product of cellular metabolism, enters the blood, the main constituent of which is water. The carbon dioxide reacts with the water to form carbonic acid, a weak acid that dissociates to yield a hydrogen ion and a bicarbonate ion. The buffering system is described by the following expression: H+ + HC0 3Hp < CO 2 + < Water Bicarbonate Carbon ion dioxide acid Note that the expression is not an equation and does not need to be "balanced." As indicated by the arrows, all the reactions are reversible. Because carbonic acid is a weak acid, unciissociated molecules are always present, as are all the other components of the system. The expression describes the system when it is at equilibrium, when the rates of the forward and reverse reactions are equal and the ralative concentrations of the components are not changing. If a system is at equilibrium, it can be "shifted to the right" by adding reactants or removing products. Conversely, it can be "shifted to the left" by adding products or removing reactants. Hydrogen ions are the important products to consider in this system. The addition of excess hydrogen ions has the efect of temporarily shifting the system to the left, as they combine with the bicarbonate ions to form carbonic acid. Eventually a new equilibrium is established; at this pOint the hydrogen ion concentration is similar to the Original concentration. If hydroxide ions are added, they combine with the hydrogen ions to form water, effectively removing a product and thus shifting the system to the right. More carbonic acid then ionizes, replacing the hydrogen ions that were removed. Organisms contain many weak acids and weak bases, thus maintaining an essential reserve of buffering capacity and avoiding pH extremes.

J\n Add and a Base React to Form a Salt When an acid and a base are mixed together, the H+ of the acid unites with the OH- of the base to form a molecule of water. The remainder of the acid (an anion) combines with the remainder of . the base (a cation) to form a salt. For example, hydrochloric acid reacts with sodium hydroxide to form water and sodium chloride: HCI + NaOH ~ H20 + NaCI A salt is a compound in which the hydrogen ion of an acid is replaced by some other cation. Sodium chlOride, NaCl, is a compound in which the hydrogen ion of HCl has been replaced by the ~ljon

Ng+.

When a salt, an acid, or a base is dissolved in water, its dissociated ions can conduct an electrical current; these substances are called electrolytes. Sugars, alcohols, and many other substances do not form ions when dissolved in water, they do not conduct an electrical current and are referred to as nonelectrolytes. Cells and extracellular fluids (such as blood) of animals and plants contain a variety of dissolved salts that are the source of the many important mineral ions essential for fluid balance and acid-base balance. Nitrates and ammonium from the soil are the important nitrogen sources for plants. In animals, nerve and muscle function, blood clotting, bone formation, and many other aspects of body

51

Water

function depend on ions. Sodium potassium, calcium, and magnesium are the chief cations present; chloride, bicarbonate, phosphate, and sulfate are important anions (Table 4.5). Table 4.5. Some B!ologically important ions

Name

Formula

Charge

Sodium Potassium

Na+ K+ H+ Mg2+ Ca2+ Fe2+ or Fe3+

1+ 1+ 1+ 2+ 2+ 2+ [iron (II)] or 3+ [iron (III)] 1+ 1121-

H~rogen

Magnesium Calcium Iron Ammonium Chloride Iodide Carbonate Bicarbonate Phosphate Acetate Sulfate Hydroxide Nitrate Nitrite

NH4+ Cl-

ICO32HC03PO 34 CH3COO5°4 2OHN0 3N0 2-

v

3-

12111-

The body fluids of terr~trial animals differ considerably from sea water in their total salt content. However, they tend to resemble sea water in the kinds of salts present and in their relative abundances. The total concentration of salts in the body fluids of most invertebrate marine animals is equivalent to that in sea water, about 3.4%. Vertebrates, whether terrestrial, freshwater, or marine, have less than 1% salt in their body fluids. Although the concentration of salts in cells and body fluids of plants and animals is small, the concentrations and relative amounts of the respective cations and anions are kept remarkably constant. Any marked changes results in impaired cellular functions and may lead to death.

The Fitness of the Aqueous Environment for Living Organis.ms Living organisms have effectively adapted to their aqueous environment and have even evolved means of exploiting the unusual porperties of water. The high specific heat of water is useful to the cell because cel~ water acts as a "heat buffer," allowing the temperature of the cell to remain relatively constant as the temperature of the environment fluctuates. Furthermore, the high heat of evaporation of water is exploited as an effective means for vertebrates to lose heat by evaporation of sweat. The high degree of internal cohesion of liquid water, due to hydrogen bonding, is exploited by higher plants to transport dissolved nutrients from the roots up to the leaves during the process of transpiration. Even the fact that ice has a lower density than liquid water and therefore floats has important biological consequences in the ecology of aquatic organisms. But most fundamental to all living organisms is the fact that many important biological properties of cell macromolecules, particularly the proteins and nucleic acids, derive from their interactions with water molecules of the surrounding medium, as we shall see.

5 CHEMICAL BONDS All matter, living and non-living alike, is composed of chemical elements, substances that cannot be broken down into simpler substances by chemical reactions. The matter of the universe is composed of 92 naturally occurring elements, ranging from hydrogen,the lightest, to uranium, the heaviest. In addition to the naturally occurring elements, about 17 elements heavier than uranium have been made by bombarding elements with sub-atomic particles in devices known as particle accelerators. About 98% of an organism's mass is composed of only six elements-oxygen, carbon, hydrogen nitrogen, calcium and phosphorus. Approximately 14 other elements are conSistently present in living things, but in smaller quantities. Some of these, such as iodine and copper, are known as traCe elements because they are present in such minute amounts. Table 5 .1 lists the elements that make up a living organism and explains why each is important. Instead of writing out the name of each element, chemists use a system of abbreviations called chemical symbols-usually the first one or two letters of the English or Latin name of the element. For example, 0 is the symbol for oxygen,C for carbon, Cl for chlorine, N for nitrogen and Na for sodium (its Latin name is natrium). Chemical symbols for the elements found in living organisms are given in Table 5.1. THE ATOM

, Imagine a bit of gold being divided into smaller and smaller pieces. The smallest possible particle of gold that could be obtained would be an atom of gold. The atom is the smallest sub-division of an element that retains the characteristic chemical properties of that element. The sub-division of any kind of matter ultimately yields atoms. This is true no matter what physical state matter may assumesolid, liquid or gas. Atoms are almost unimaginably small-much smaller than the tiniest particle visible under a light microscope. By special scanning electron microscopy, with magnification as much as 5 million times, researchers have been able to photograph some of the larger atoms such as uranium. Atomic Structure An atom is composed of smaller components called subatomic particles. For our pupose, we need consider only three types-protons, neutrons and electrons: (a) protons, which carry a positive charge; (b) electrons, which carry a negative charge and (c) neutrons, which carry no charge. Protons and neutrons make up most of an atom's mass and form the dense atomic nucleus. The electrons,which are much smaller that either protons or neutrons, rotate around the nucleus. Each element has a different number of protons, with the same number of electrons. Because the number of positive charges on protons is equal to the number of negative charges on electrons, atoms are electrically neutral. 52

53

Chemical Bonds

Table 5.1. Elements that make up the human body

Name

Chemical Symbol

Approximate composition by mass (%)

Oxygen

°

Carbon

C

18

Hydrogen

H

10

Nitrogen Calcium

N Ca

3 1.5

Phosphorus

P

1

Potassium

K

0.4

Sulphur Sodium

S Na

0.3 0.2

Magnesium Chlorine

Mg CI

0.1 0.1

Iron

Fe

Trace amount

Iodine

65

Trace amount

Importance or function

Required for cellular respiration; present in most organic compounds; components of water. Backbone of organic molecules; can form four bonds with other atoms. Present in most organic compounds; component of water. Component of all proteins and nucleic acids. Structural component of bones and teeth; important in muscle contraction, conduction of nerve impulses and blood clotting. Component of nucleic acids; structural component of bone; important in energy transfer. Principal positive ion (cation) within cells; important in nerve function; affects muscle contraction. Component of most proteins. Principal positive ion in interstitial (tissue) fluid; important in fluid balance; essential for conduction of nerve impulses. Needed in blood and other body tissues. Principal negative ion (anion) of interstitial fluid; important in fluid balance; component of sodium chloride. Component of haemoglobin and myoglobin; component of certain enzymes. Component of thyroid hormones.

Other elements found in very small amounts in the body include manganese (Mn), copper (Cu), Zinc (Zn) , cobalt (Co), fluorine (F), molybdenum (Mo), selenium (Se) and a few others. They are called trace elements.

Each kind of element has a fixed number of protons in the atomic nucleus. This number, called the atomic number, is written as a subscript to the left of the chemical symbol. Thus lH and 80 indicate that the hydrogen nucleus contains one proton and the oxygen nucleus has eight protons. It is the atomic number, the number of protons in the nucleus,that determines the chemical identity of the atom. The total number of protons plus neutrons in the nucleus is termed mass number and is indicated by a superscript to the left of the chemical symbol. The common form of oxygen atom, with eight protons and eight neutrons in its nucleus, has an atomic number of 8, and a mass number of 16. It is indicated by the symbol 16p. The number of neutrons in an atom varies, although commonly the number of neutrons is equal to the number of protons. Variations in the number of neutrons do not change the chemical characteristics of the element but do produce various isotopes of the element. Each isotope has a characteristic number of neutrons. Many isotopes are unstable. As they revert to a stable condition, radiation is given off. Electrons are arranged in a number of orbitals or energy-level 'shells,' that differ in their average distance from the nucleus. The dis-tance of an orbital from the nucleus predicts the energy state of the electrons: the greater the distance, the higher the energy state.

Biomolecules

54

The number of electrons in each orbital is predictable: two in the orbital closest to the nucleus, eight in the next orbital, eighteen in the third, thirty-two in the fourth and so on. However, eight electrons in the outermost orbital beyond the first is particularly stable. If the outer orbital has more or less than eight electrons, the atom will have a certain degree of instability. The instability causes these atoms to react with other atoms by forming chemical bonds.

•

Proton

Q

Neutron

A hydrogen atom has: 1 proton 1 electron o neutrons

e

Electron

A Carbon 2ltom has: 6 protons 6 electrons 6 neutrons

CHEMICAL BoNDS

The chemical properties of an element are determined primarily by the number and arrangement of electrons in the outermost energy level (electron shell) . In a few elements, called the noble gases, the outermost shell is A nitrogen atom has: filled. These elements are A oxgen atom has: 7 protons 8 protons chemically inert, meaning 7 electrons 8 electrons that they will not readily com7 neutrons 8 neutrons bine with other elements. Two such elements are helium with two electrons (a complete inner shell) and neon with ten electrons in the outermost. Energy level of an atom is referred to as valence electrons. When the outer shell of an atom contains fewer than eight electrons, the atom tends to lose, ~ain, or share electrons A sulphur atom has: A phosphorus atom has: to achieve an outer shell ·of 16 protons 15 .protons eight (zero or two in the case 16 electrons 15 electrons 16 neutrons of the lightest elements). 15 neutrons The elements in a given Fig. 5.1. Simplified diagrams of atoms of the elements hydrogen, carbon, oxygen, compound are always present nitrogen, phosphorus and sulfur. in a certain proportion by mass.This reflects the fact that atoms are attached to each other by chemical bonds in a precise way to form the compound. A chemical bond is the attractive force that holds two atoms together.~ Each bond represents a certain amount of potential chemical energy.

55

Chemical Bonds

The atoms of each element form a specific number of bonds with the atoms of other elements-a number dictated by the number of valence electrons. Chemical bonds may be weak or strong. Strong bonds hold the atoms of a molecule together. The amount of energy which is usually released in the formation of strong bond like covalent bond ranges from 50 to 200 kcal/mol. On the other hand the amount of energy released in the formation of weak bonds is only 1 to 5 kcal/mol. Thus the strong bonds are responsible for maining the structure of large molecules. However, it is not always true. The stability of three-dimensional structure of large molecules often stems from weaker bonds. Many of such weak bonds can act together to form very stable structure. This type of joint working of weak bonds help in maintaining stable association with other molecules. For example: when two protein molecules bound together by four weak bonds and if two of the four bonds break, the protein molecule will still bind together by the other two bonds and thus facilitating the reformation of the broken ones. Thus such bonding pattern plays a crucial role in maintaining biological structures. The two principal types of chemical bonds are covalent bonds and ionic bonds. IONIC BONDS

One extreme of polarity is the ionic bond, in which electrons are pulled completely from one atom to the other. The number of protons in the nucleus remains unchanged so the loss or gain of electrons produces an atom with a net positive or negative charge. Such electrically charged atoms are termed ions. battery Ions (from the Greek ion, 'going') are atoms or molecules that positive negative bear electrical charges as a result of having gained or lost one or electrode electrode (anode) more electrons. An atom that loses one or more electrons becomes (anode) positively charged. Such an atom moves towards the negative electrode or cathode in an electric field (the space between positively and negatively charged electrodes through which an electric current passes) and is known as a cation. An atom that gains electrons be~omes negatively charged, moves toward the anode in an electrical field and is known as an anion. Being oppositely charged, anions and cations are also attracted to and tend to bind to each other. The binding of anions with cation constitutes an ionic bond. Because they are at a lower level of potential energy, the atoms in ionic Fig. 5.2. Mouement of anions and cations to the cathods and compounds such as sodium chloride are in a more stable anode, respectiuely, in an arrangement than unbonded sodium and chlorine atoms. electric field. Atoms with five, six or seven electrons in their outer shell tend to gain electrons from other atoms and become negatively charged anions (e.g. CI-, chloride ion). Charged particles, both anions and cations, play many important roles in biological roles in biological systems, such as the transmission of nerve impulses and the contraction of muscles. An ionic bond is the force of electrical attraction between two oppositely charged ions. When held together by ionic bonds, oppositely charged ions form an ionic compound. Sodium chloride is a good example of an ionic compound. A sodium atom, with atomic number 11, has two electrons in its inner shell, eight in the Clsecond, and one in the third shell by obtaining seven electrons from other atoms because it would then have a vast excess of negative charge. Instead, it gives up the single electron in its third shell to some electrons acceptor, leaving the second shell as C __ l-.. Clthe complete outer shell. A chlorine atom, atomic number 17, , has two electrons in its inner shell, eight in the second , and seven in the third shell. The chlorine atom achieves a complete outer shell not by losing the seven electrons in its third shell (for SodIum chlorIde crystal.

8

8 8

Biomolecules

56

it would then have a vast positive charge) but by accepting an electron from an electron donor, such as sodium to complete its third shell. When an electron donor such as sodium meets an electron acceptor such as chlOrine, the electron may be transferred completely from the donor to the acceptor. The sodium ion now has 11 protons in its nucleus and ten electrons around the nucleus, giving it a net charge of 1+. The chlorine ion has 17 protons its nucleus and 18 electrons around the nucleus, so it has a net charge of 1. These ions attract each other as a result of their opposite charges. They are held together by this electrical attraction to form sodium chloride (NaCn, common table salt. This transfer of one or more electrons from one atom to another and the binding together of two ions of opposite charge (the ionic bond) results in the formation of an ionic compound . Ionic bonds occur between electron donors and electron acceptors. Whether an ionic compound is in solid from or is dissolved in water, its constituent particles (ions) do not share electrons. Because of this, the term molecule does not adequately explain the properties of ionic compounds such as NaCI. Chemists simply refer to them as compounds.

Sodium,the donor gives its one electron in orbital 3 to complete the third orbital of chlorine,which lacks one.

SOdium has 11 electrons: 2 in the first orbital 8 in the second orbital 1 in the third orbital

Chlorine has 17 electrons: 2 in the first orbital 8 in the second orbital 7 . in the third orbital

Ionic bond

A sodium ion has one more proton than electrons and therefore has positive charge.

Chlorine has one more electron than protons, producing a negaitive charge.

Fig. 5.3. The formation of ionic bonds.

57

Chemical Bonds COVALENT BONDS

In covalent bonds, electrons are not donated from one atom to another, but are shared between atoms. Thus, a covalent bond consists of a pair of electrons that are shared between two atoms and that occupy two stable orbitals, one of each atom. If a single pair of electrons is shared between a pair of atoms, it is called single bond. Sometimes, two atoms share two or three pairs of electrons to form double or tripl'e bonds. Covalent bonds are more stable than the units that result from ionic bonds. If a molecule contains atoms of two or more elements, the substance is a molecule of a compound. If the joined atoms are of the same element, the substance is the molecular form of that element : molecular oxygen, molecular nitrogen and so on. How Covalent Bonds are Formed We can gain some understanding of covalent bonding by visualizing the orbitals of two or more atoms and imagining what happens when the atoms closely approach each other. Consider what happens, for example, when two hydrogen atoms unite with an oxygen atom in forming the familiar compound water (HP). Oxygen, with atomic number 8, has eight electrons each in its Is, 25 and 2p orbitals and one in each of the other 2p orbitals. Because each orbital can accommodate two electrons, oxygen has two vacancies in its outermost energy level. The hydrogen atom's only electron is in a Is orbital, at the first energy level-thus hydrogen has one vacancy. When two hydrogen atoms are dose to an oxygen atom, the two half-filled 2p orbitals of the oxygen atom and the spherical s orbital of each hydrogen atom overlap. The outer orbitals of all three atoms can now be considered filled since they are now sharing electrons. In the newly formed molecule, two pairs of electrons are attracted by and actually reside for -a substantial fraction of the time with all three nuclei. This sharing of electrons and their attraction to all the nuclei constitute the two covalent bonds that are formed. The nuclei cannot approach more closely because of their

•

2 hydrogen..atoms + 1 oxygen atom

Water molecule

Fig. 5 .4. Formation of coualent bonds.

58

Biomolecules

mutual repulsion of each other. The electrons and nuclei maintain those distances from each other at which all attractive and repulsive forces are at equilibrium, thereby forming a stable bond. Polar Covalent Bonds In a water molecule, the oxygen nucleus attracts the shared electrons much more strongly than do the small hydrogen. The result is an unequal distribution of charge within the molecule. The end with the oxygen atom is somewhat negative because it has more electrons moving about it at any given moment than there are protons in its nucleus. Conversely, the end of the molecule with the hydrogen atoms is slightly positive because of the general absence of electrons from that region. These small differences in charge are denoted by the small Greek letter delta: 8+ and 8-. The pairs of oppositely charged en.ds in such molecules are called electric dipoles. In electric field , the negative end of the molecule will be oriented towards the positive pole and the positive end of the molecule towards the negative pole. Because there is a difference between the two ends or poles of these molecules, they are known as polar molecules. The bonds responsible for producing polarity are termed polar covalent bonds. The polarity of the water molecLlles is crucial to the biochemical reactions that underlie many biological processes. In molecules such as H 2 , 02' and F2' which are found of pairs of identical atoms, the shared pair of electrons is visualized as spending as much time with one nucleus as with the other. The covalent bonds in these molecules are non-polar.

Fil:ld off

.Hydrophobic Interactions Another kind of bonding occurs among molecules that have little or no affinity for water. These molecules tend to group together, when in an aqueous environment, forming droplets. For Field on example, When oil is added to water, the two do not mix. This Fig. 5.5. Orienta tion of polar molekind of interaction is seen in biological membranes, which are cules in an electric field . composed largely of lipids. The oil molecules remain together not so much because they are r.lttracted to each other but because they are forced together by surrounding water molecules. Weaker Interactions In addition to ionic and covalent bonds, there are also several types of weaker interactions among atoms and molecules. These weaker interactions are largely responsible for the three dimensional shapes of many large biological molecules such as proteins and for some of their functional properties. HYDROGEN BONDS

Water is a very special compound with many unique features important to living systems. Many of these special features can be attributed to the hydrogen bonds that exist between water molecules. A single water molecule is polarized, which means that although the oxygen and hydrogen atoms share electrons, .the oxygen atom attracts the electrons towards itself and away from the hydrogen atoms. Therefore, the area of the water molecule occupied by the oxygen atom is slightly more negatively charged than the area occupied by the hydrogen atoms. Water molecu1es form lattices when the negatively charged oxygen area of one water molecule is linked to the more positively charged hydrogen area of another water. molecule. In effect, a hydrogen ·atom forms a bridge between two water molecules. This bridge is a hydrogen bond. Because water molecules are linked by hydrogen bonds, water is a very cohesive substance and has a high surface tension. A large amount

59

Chemical Bonds

of energy is necessary to heat water or to change water from a solid to a liquid or from a liquid to a gas. The polarized nature of water molecules makes it a nearly universal solvent, and the heat + capacity of water allows life to exist over a wide temperature range. Hydrogen bonds enhance the average molecular weight and are also responsible for the reduction of free molecules. At the time of Few electrons formation and breaking of hydrogen bonds exothermic and endothermic reaction occur respectively. In molecules like alcohol, water, carboxylic acids, primary and secondary amines hydrogen bonds are formed between donor and acceptor groups. However, hydrogen uands fail Nagilive charge to develop between the molecules of saturated hydrocarbons. There are two types of hydrogen bondings : inter-molecular and intramolecular. The inter-molecular hydrogen bonding tends to aggregate molecules under the Poslilive charge influence of shape of the molecule resulting the Fig. 5.6. Water is a polarized molecule because the oxygen

-

atom attracts the electrons, giving a portion of

greater or increased molecular weight of a the water molecule a negative charge. The portion structure thus developed. Intra-molecular hydrogen of the molecule formed by the hydrogen Qtoms bonding occur in the molecules like alcohol, water, consequently is more positively charged. carboxylic acids, primary and secondary amines. The intra-molecular hydrogen bonding develop only when the proton donor and proton acceptor sites on the same molecule appear in a favourable spatial configuration. Hydrogen Bonding in Biological System Water is the universal biological solvent. The amount of water in living cells varies from 60% to 70%. However, in blood plasma the amount of water is as high as 90%. Water is the most classical example for hydrogen bonding. In all living systems the biomolecules which are nearer to water molecules have a greater impact on the structure and conformation of biopolymers. These biopolymers undergo inter molecular hydrogen bonding among themselves and with adjoining water molecules and almost all polar molecules are inter-linked by a network of hydrogen bonded water molecules in a living system. The most predominent and extensive hydrogen bonding within a cell is responsible for proper organization of non polar lipid molecules. Hydrogen Bonding in Proteins Proteins are the macromolecules present in all biological systems. Every protein molecule is a polypeptide molecule showing both intra-molecular and inter-molecular hydrogen bonding. Inter-molecular hydrogen bonds ocCur between the polypeptides whereas intra-molecular hydrogen bondings occur in helicular conformation of peptides. In the secondary structure of protein hydrogen bonds develop between carboxylic (COOH) group of one amino acid and amino group (NH 2) related to the same peptide chain. Formation of hydrogen bond in the side chains of amino acids is probably a guiding force in transforming a-helix into tertiary and quaternary structure of protein. The hydrogen bond in polar side chains and non polar side chains decide the extent of affinity of water towards protein (water solubility). Hydrogen Bonding in Carbohydrates Carbohydrates are among the most widely distributed compounds in both plant and animal kingdoms. Plants and animals contain large quantities of carbohydrates as reserve food material. In

60

Biomolecules .

Fig. 5. 7. Water is a cohesive substance because hydrogen bonds form between water molecule. The negative charged part of one molecule is linked to the more positively charged portion of another molecule by hydrogen bonds.

disaccharides and polysaccharides hydrogen bonds occur abundantly with bond distance varying trom' 2.65 to 3.03A. There are two intra-molecular hydrogen bonds in sucrose. In polysaccharides the aggregation of cellulose fibrils into microfibrils is due to extensive hydrogen bonding. Multiple hydrogen bonds bring about stability and organization of primary cell wall in plant cells.

Hydrogen Bonding in Deoxyribonucleic Acids DNA is composed of two chains or back bones running in the anti parallel fashion. Both the chains coil around each other forming a helicular structure. This helicular nature of DNA is done to extensive hydrogen bonding . Hydrogen bonds are formed between purine and pyrimidine bases. To correct adenine and thymine together two hydrogen bonds are developed. Similarly between cytosine

61

Chemical Bonds

Polypeptide' chain

Fig. 5.B.

~ ·structure formed between two stretches of a polypeptide chain in antiparallel orientation. Hydrogen bonds between the two sections of the polypeptide stabilize their association

H

I

H-C-OH

I

C=O

I

H

I

C=O

I

H-C~H

I

OH-C-H

OH-C-H

OH-C-H

OH-C-H

I I

I

I

H-C~H

H-C~H

H-C~H

H-C~H

I I

I

I

H

H

D-fructose

D-galactose

and quanine three h~rogen bonds are formed . These multiple weak h~rogen bonds provide stability to DNA molecule. During replication and transcription several hydrogen bonds break down release both the chains in the related part of DNA move apart and the remaining hydrogen bonds will facilitate the reformation of broken ones. Thus such molecules can maintain stable associations if enough weak bonds work together. CO-ORDINATE BONDS

Co-ordinate bonds are basically the covalent bonds in which the two shared electrons come from the same atom. In this pairing of two atoms, one atom has one spare pair of electrons and is called the donor atom, and the other atom lacks a pair of electrons and is called acceptor atom. After pairing or formation of covalent bonds, this pair of electrons becomes the property of both the atoms and held in common. The co-ordinate bond is represented by an arrow pointing from the donor atom to the acceptor atom as shown in the diagram given below. PEPTIDE BONDS

H T

~.~

~1 \N -I

-N H

\

.

.

I

~

C-N

~O

'J ,.H T . :. 1 = H H

_-. .,-. .

T

~ ". H ~\.~--

§

Q,

'cl'c T

T

c/ ', .,-N

r

Proteins are the building stones and are the most essential Fig. 5.9. Hydrogen bonding between macromolecules in living beings. Proteins are chemically made of two nucleotides (G-C). amino acids. Amino acids are amphoteric molecules and each one has an amino (NH z) group at its one end and a carboxylic (COOH) group at its other end. In all living cells, these amino acids

62

Biomolecules

come and react together to form protein. For this, the various amino acids are linked together. This linkage is developed between the NH2 group of one amino acid and COOH group of other amino acid. In fact, this is the amide linkage which is called peptide bond. A chain with one peptide bond is the monopeptide, with two peptide bonds, dipeptide and with more than two peptide bonds is called polypeptide chain, which is the protein. DISUlPHIDE BONDS

In some cases, a protein is made up of two or more than two polypeptide chains. These polypeptide chains are held together by crosslinks formed by disulphide bonds. Such bonds are developed only when two or more polypeptide chains come together and develop disulphide bonds simply by removing one atom of hydrogen from the sulph~ryl group of two cysteine residues of the different polypeptide chains. Disulphide bonds are also covalent bonds. These are the strongest bonds formed between different polypeptide chains. Sometimes, the disulphide bonds are formed even between cysteine residues at different loci in the same peptide chain. I I I I o 0 0 0 I I I I O-CH 2. SH + H.S. CH2-D-.07O--CH2-S-S--CH2-D I I I I o 0 0 0 I I I I Chain I Chain II Disulphide bond. GLYCOSIDIC BoNDS

In living organisms, carbohydrates occur CH,OH mainly in three forms-monosaccharides, oligosaccharides and polysaccharides. Oligosaccharides are composed of various units of monosaccharide~ '----O--~ which are held together by glycosidic bonds. The glycosidic Ilnlcag6 most common oligosaccharides encountered in nature are disaccharide like Maltose. Maltose has H OH H OH two molecules of glucose which are linked together D-Glucose D-GIucose by glycosidic bonds. These bonds are formed between the hydroxyl group or the C-l carbon atom of one molecule and the hydroxyl group on the C-4 carbon atom of another molecule of glucose as shown in the diagram. INTERMOLECUlAR FORCES

Ultimately, the behaviour and properties of a molecule are determined by its structure. Structure determines the mechanism by which molecules interact with other molecules and it is these interactions upon which all purification methods depend. Let us summarize the basic intra-and intermolecular forces which are so important in determining the properties and behaviour of a molecule in all the various purification systems encountered in biochemistry. Ionic (Coulombic) Forces These are the strong electostatic forces between oppositely charged ions. These are for example the major forces involved in the structure of inorganic salts such as sodium chloride, and they account for the high melting point of such compounds.

Dipole-dipole Forces Most molecules are not ionic but because their bonding electrons are distributed unevenly over the molecule, they have permanent dipoles. Different atoms have different electronegativities, i.e. they

63

Chemical Bonds

have different electron-attracting power, and so the molecule is partly polarised: eg HC1. 0Cl

~ ~0 ~ ~~ ~

>

Dipole moment

The' dipole moment, p, is the product of one of the charges q, in the dipole and the. distance of separation, r of the lie li+ a· charge$. (p = qr) . The dipole moment of Fig. 5.10. Dipole-dipole interactions. a molecule as a whole is a vector sum of the group moments. Dipole moments are usually defined in terms of the Debye, 0 (0 = 3.336 X 10-30 C m). The positive end of a molecule with a permanent dipole will be attracted to and align itself at, the negative end of a similar molecule; dipole-dipole interaction takes place. A molecule will not have a permanent dipole, however, if it is completely symmetrical, even if it consists of atoms with different electronegativities, e.g. CC1 4 .

Van'der Waals Forces and London - Dispersion Interaction Van der Waals is a weak attractive force that is produced when two atoms approach one another clo~ely. In an atom any fluctuation in the distribut!on of its electrons gives rise to transient dipoles. Van der Waals force is operative for small distances and it is being considered as short range force. The proposal on existence of Van der Waals force was made by Dutch physicist Van der Waals. The fQrce would be originated when two dipoles come · together within very short distance. There are different types of forces recognized based on the nature of dipoles. One is Keeson force formed due to the interaction of two permanent dipoles i.e., positive end of one dipole is in close proximity to the negative end of the other dipole, (a) Intermolecular hydrogen bonding between water molecules which elicites attraction between the molecule; another type is the Debye force. These are the dipole-induced dipole attraction formed due to the polarization of a non-polar molecule in the electric field of a dipolar molecule leading to an attraction between the tWo molecules. Van der Waals interactions occur among all types of polar and non polar molecules, particularly, they are responsible for the attachment among the molecules of nonpolar liquids and solids. The attraction declines rapidly as the distance between these two atoms (b) intramoleular hydrogen bonding in salicylic acid increases. Thus, force is effective only when atoms are close to one another. However, if atoms get too close together, they get repelled by the negative charges in their outer electron shell. The energy of the Van der Waals interaction attains about 1 kcallmol. OH This value is slightly higher than the Fig. 5.11 . Examples 0/ hydrogen bonding. kinetic energy of molecules at 25°C. In

,

64

Biomo/ecules

aqueous environment its energy value has 1 to 2 kcal/mol. The interaction between antibody with antigen molecule are the Van der Waals contact. Another Van der Waals interaction is the enzyme and their specific bound substrates. Van der Waals forces also called London dispersion forces also help hold aggregates of smaller molecules together. For Instance, the phospholipids that are major components of cell membranes are held together mainly by these forces. London-dispersion interaction is caused by the attraction of electrons of one molecule by the nucleus of another molecule. This type of interaction occurs both in non-polar and polar molecules. Generally London-dispersion interaction increases with molecular mass because the number of electrons that can be attracted increases with molecular mass. Because of this H 2 , F2 and Cl 2 are gasses at room temperature while Br2 is a liquid and 12 a solid under these same conditions. The dipole-dipole interaction, the dipole-induced dipole interaction and the London-dispersion interaction are often included under the general heading of Van der Waals interactions. The important facts to remember are that the size of the force is determined by the relative polarizability of the electnms of the constituent atoms. This In tum depends on the size of the atoms and the degree to which the electrons arB held by the nucleus. Generally, in a given family of ~Iemerits , the larger atoms are more polarizable as their outer electrons are more shielded from and are therefore less tightly held by, the nucleus.

6 CELL AS A . UNIT OF LIFE .The cell is the smallest but complete expression of the fundamental structure and function of all organisms and that is why it is also called a unit of biological activity. It is delimited by a semipermeable membrane and capable of self reproduction in a suitable non-living medium. The body of allUving organisms (bacteria, blue green algae, plants and antmals) except viruses has cellular organization and may contain one or many cells. The organisms with only one cell in their body are called unicellular organisms (e.g. bacteria, blue green algae, some algae, protozoans etc). The organisms having many cells in their body are called multicellular organisms (e.g. most plants and animals). Any cellular organism may contain only one type of cells-prokaryotic cells or eukaryotic cells. It was Robert Hooke who first of all in 1665 observed under a microscope honeycomb like structures in cork, and he applied the term 'cell' (L. cella; small room) for the same. Previously it was believed that this is only one component of the cell which is cell wall. But in 1831, one more peculiar structure was observed by Robert Brown. He gave the name 'nucleus' . Later on in 1835, the word 'sarcode' was proposed for the jelly like material present inside the cell by Dujardin. In 1840, Purkinje replaced the word sarcode by protoplasm which is now used universally. VIRUS~

Virus particles outside their host cells consist of a core of nucleic acid, either DNA or RNA, surrounded by a coat of protein. All degrees of complexity exist in viruses, from simple particles consisting of core with a single nucleic acid molecule and a coat of protein molecules of a singfe type of more complex particles with coats made up of more than 50 different kinds of proteins. A few viruses infecting animal cells such as the influenza and herpes visuses, are surrounded by an outer membrane derived from the plasma membrane of their hosts. The nucleic acid molecules of viral particles may be either linear or circular. Some of the linear molecules consist of a single nucleotide chain, while others are in double helical form. The viruses infecting plant cells usually contain linear RNA molecules, viruses infecting animal or bacterial cells may contain either RNA or DNA molecules, in either linear or circular form. The protein coat enclosing the nucleic acid core, depending on the virus is either rod-spherical or lollipop shaped. The coat of the much studied tobaccomosaic virus, for example, is a rod shaped structure about 15 by 300 nm long, built up from more than 2000 identical protein units. The RNA molecule of this virus winds into a helix extending through the axis of the rod. The viruses infecting bacteria called bacteriophages, are among the most complex viral particles known. Best studied of these are the bacteriophages attacking the bacterium Escherichia coli. The "T-even" E. colt bacteriophages (T2 , T4 , and T6) have a polyhedral head enclOSing a DNA core and a tail containing several different proteins. The tail is complex in structure and consists of a collar at the 65

66

Biomolecules

a

point of attachment to the head, cylindrical b~eplate. The base plate carries six long hair like extensions, the tail fibres . The life-cycle of the T-even bacterio90 phages illustrates the general pattem by which virus particles infect their host cells. Free bacteriophage particles randomly collide with bacterial cells. If T-even cell virus particle head protein subunits ~~~[~~~ colloides with E. coli cell, the tail fibres 'recognize' and bind specific sites on the bacterial cell wall. The head and tail sheath then contract and inject the DNA core into coIlar-----' the bacterial cell, while the proteins of the ~"\!(--sheath virus remain outside. Viruses are best classified as non-living matter when they are outside their host cells. In this form, they carry out none of the tail fiber's activities of life and are inert except for the capacity to attach specifically to their host cells. They can be crystallized in this nonliving form and stored indefinitely without change or damage. The viral nucleic acid molecule carries only the information required to direct the host cell machinery to make more viral particles and is active in this function only when inside the host cell. Thus a virus particle probably represents nothing more or less than a fragment of a nucleoid or chromosome derived from a once living Fig. 6.1. Viruses. cell that is reduced to a set of coded directions for making additional particles of the same kind.

sheath and a

ONAc~e-----~~~lii

PROKARYOTIC

CEUS

Prokaryotes are divided into two major groups, the eubacterla and archaebacterla . Both groups share many properties that clearly separate them from eukaryotes. Most species of bacteria are eubacteria. Eubacteria can be divided into two major groups : the non-photosynthetic and photosynthetic bacteria. The bacterium E. coli is used as an example of non-photosynthetic eubacteria.

Non Photosynthetic Eubacteria E. coli is a bacterium that commonly inhabits the intestinal tract of humans and other animals. It is a cylindrical cell about 2 Ilm long and 1 Ilm in diameter with a volume of about 1.6 Ilm3. On its surface are a number of filamentous appendages called flagella, usually six, by which it rapidly propels itself. One cm3 (about one gm) of packed E. coli contains 50 x 109 cells. An individual cell grows by increasing its length while maintaining a constant diameter. The cell divides into two daughter cells by the forming of a partition through the middle of the cylinder. The genetic constitution of E. coli allow the organisms to grow and divide in a medium containing only a few kinds of inorganic ions, and a source of organic carbon, for example, the sugar glucose. Thus DNA of this bacterial cell contains genes for all of the enzymes needed for the synthesis of all the amino acids, nucleosides, fatty acids and other components needed to make macro molecules,

Cell as a Unit of Life

67

using only a simple organic molecule, such as glucose and inorganic salts as a starting material. In glucose containing medium, E. coli doubles in size and divides every 40 minutes at the optimum temperature of 37°C. If amino acids, nucleosides and other nutritionally useful organic molecules are added to the minimal culture medium, the rate of cell growth and division increases. The increase occurs ::-4,~&/iiie~l-lloluble RNA because provision of useful nutrient relieves the cell of the need to synthesize these components. By addition of rich variety of nutrients to the medium and E. coli, cell approaches the upper limit in its reproductive rate , doubling all of its contents and dividing every 20 minutes. The Fig. 6.~. A Prokaryote cell . . upper limit of growth rate is probably set by the maximum rate at which E. coli can synthesize macro molecules with a generation time of only 20 minutes and a single cell could give rise to more than 3 x 1011 cells in just 12 hours. Prokaryotic cells usually grow and reproduce much more rapidly than eukaryotes. For instance, for a mammalian cell, seven hours is about the shortest generation time or cell cycle time-the time a cell needs to go from one cell division to the next . Some protozoa which are unicellular eukaryotes, have generation times as short as two [lours nutritionally rich medium, and one kind of yeast cell can divide every 75 minutes; but none of the eukaryotes approach the rapid proliferation rates common among prokaryotes. Prokaryotes are structurally simple cells. Most types have a rigid cell wall made of polysaccharides, peptides and lipids laid down outside the cell. A few types, for example, a small bacteria known as mycoplasma, lacks an extracellular wall. The rigid wall of rod-shaped bacteria maintains the cylindrical shape of the cell. vacuale Other species of bacteria have walls that produce a spherical or spiral form . The cell wall provides mechanical protection, particularly against osmotic pressure. Immediately inside the cell wall of E. coli is the plasma membrane which completely encloses the cell. The plasma membrane consists of a double layer of lipid molecules with many associated protein Fig. 6.3. Diagram 0/ a Mycoplasma .

68

Biomolecules

molecules. In cqntrast to the cell wall, which is porous and therefore cytoplasm capsule penetrable by molecules and ions, the plasma membrane severely restricts the diffusion of molecules and ions in and out of the cell. Thus the cell wall membrane serves the critical role of retaining desired substances inside the plasma cell, although it also limits diffusion into the cell of environmental substances membrane necessary to sustain cell metabolism. Certain specialized proteins bound to the lipid bilayer of the plasma membrane greatly enhance the inward passage of inorganic ions, sugars, amino acids, nucleosides and other dissolved RNA materials which are useful to the cell. Other proteins bound' in the plasma membrane of a bacterial cell catalyze the process by which the energy ribosomes contained inorganic molecules are converted into a chemically . unstable form. The intracellular contents of prokaryotes such as E. coli are present flagella in two major structural parts, the nucleotide and the cytoplasm. The nucleotide consists of a single DNA molecule (the chromosome) condensed into a irregularly shaped fibrous network, which occupies a few percent of Fig. 6.4. Bacterium. total cell volume. It is thought that the nucleotide is attached at one point to the plasma membrane. This attachment of the chromosome to the membrane may help both the control of chromosome replication and in the separation of daughter chromosomes during the cell division. The cytoplasm of E. coli contains approximately 25,000 tiny p~cles called ribosomes, floating in a solution called the cytosol. Each ribosome is a machine for synthesizing proteins. The cytosol, which contains a large variety of ions, small organic molecules and enzymes, is the site where the cell carries out most of its metabolic activities.

Nucleoid (DNA)

PhotOlynthetic lameflle with phycobiiolOmeI

(DNA)

A

Fig. 6.5. Prokaryotic cells. A-A cell of blue-green algae, B-A Bacterial cell.

Much has been learned about the molecular biology of the cell from the study of prokaryotes, in particular from the study of E. coli. In part, bacteria were chosen as research materials because they are functionally and structurally far less complex than any of the eukaryotic cells. In addition, the fast growth rate and low number ·of nutritional requirements of bacteria such as E. coli constitute a great practical advantage for research because large number of cells can be obtained in a few hours with a simple and inexpensive culture medium. E. coli is not the smallest type of cell known. Some bacteria, the mycoplasmas, have volume as small of 0.02 11m3, compared to a minimum volume for E. coli of 1.6 11m3. Mycoplasmas lack cell walls, and their chromosomes can be as small as one fifth of the chromosome in E. coli. These are the smallest chromosomes known among bacteria. The mycoplasmas were identified about 1900 as the cause of respiratory diseases in animals and gained attention during World War II as the causative agents of pneumonia among U.S. Army recruits. Mycoplasmas are sometimes referred to as PPLO, which stands to pleuropneumonia like organism. Due to their small size and small amount

69

Cell as a Unit of Life

of DNA, the mycoplasmas are no doubt genetically and functionally iess complex than E., coli. However, they require a nutritionally complicated medium for growth and grow slowly and hence are less convenient to use in research. Nevertheless, the study of mycoplasmas has intensified during recent years, and these simplest of known cells may well provide unique insight into principles ot cell organization and operation. Mycoplasmas are sometimes referred to as the minimum cell because they approach the minimum genetic and molecular complexity necessary to sustain the life and reprodudion of a cell.

Photosynthetic Bacteria The photosynthetic bacteria probably arose from non--photosynthetic bacteria very early in the course of evolution, perhaps as early as 3.1 billion years ago. Most photosynthetic bacteria are obligate photoautotrophs. Photoautotroph means requiring only light, water, inorganic ions and CO 2; obligate means that for growth light is necessary because these bacteria cannot use organic compounds like sugars as an alternative source of energy. Photosynthetic bacteria are widely distributed in fresh and salt water and in soil. The enormous mass of photosynthetic bacteria growing in the oceans generates much of the oxygen in the Earth's atmosphere. Whenever seen under an electron microscope, a photosynthetic bacterium is enclosed by a rigid wall and immediately inside the wall, by a plasma membrane. As in other bacteria, the cytoplasm is rich in ribosomes, and a nucleoid is present. In contrast to other kinds of bacteria, however, photosynthetic bacteria often have extensive internal membranes that contain light-absorbing pigments and the machinery for photosynthesis. Photosynthesis is the process by which the energy of light is captured and used to synthesize sugar, starting with carbon dioxide and water. EUKARYOTIC CELLS

Eukaryotic cells may be unicellular organisms, such as protozoans and unicellular algae, or they may be cells that make up the tissues and organs of multicellular organisms. Though eukaryotic cells have different shape, size and physiology but all the cells are typically composed of plasma membrane, cytoplasm and its organelles, viz, mitochondria, ribosomes, endoplasmic reticulum~ Golgi complex etc., and a true nucleus. Shape

The shape of the cells varies. Cells are found in varieties of shapes. Some cells, such as amoebae and leucocytes, change shape. Some other cells possess a typical shape, which is more or less specific and fixed viz. spermatozooids, erythrocytes, epithelial cells, nerve cells and muscle cells. The shape of the cells varies mainly due to the functional adaptation, surface tension, viscosity of the protoplasm, mechanical action excited by the adjoining cells and rigidity of the cell membrane. Thus, the cells become cuboidal, polygonal, columnar, flat or plate like structures. Nerve cells, which serve for transmitting electrical impulses over long distances, possess long extensions. These extensions may, sometimes, be several centimeters in length. Muscle cells are also elongated so that force of action may be exerted properly in one direction. Epithelial cells are flattened . The supporting cells of the plant possess thick walls. When masses of cells are packed together, they take the shape of polyhedral solids, having many faces. Thus shape allows a close pack of the cells, although regular polyhedral of four, six and twelve sides can be packed even without interstics. 14-sided polyhedron satisfies the condition of minimal surface. After observing the cells takirlg into consideration of three dimensions and reconstruction, it is revealed that many cells of animals and plants are on an average very close to fOl,lrteen faces . Numbers Protozoa are single celled, while other animals are multicellular. Rotifers consist of a few hundred cells. Hydra , too, is made up of a few thousand of cells. In higher animals, generally, there are millions and trillions of cells. For example, man has about 26 trillions of cells in the body. In the human brain, the grey matter alone is made up of 92 billion cells. The size of the animal depends on

Biomolecules

70

the number of cells and not on the size of cells. Swift moving and active anin'lals, like insects arid birds, possess fewer cells per unit volume in comparison to the sluggish and lethargic animals. Size

secretory

granull/l~

Goigi

rough endoplasmir: reticulum

The size of the different cells ranges rlbosom"-!::::::~;;J.n within broad limits. Some plant and animal ph.gocvtosls cells are visible to the naked eyes, such as eggs of certain birds have a diameter of several centimeters. But great majority of cells are visible under the microscope. The smooth smallest living cells are endoplasmic reticulum found among bacteria and viruses where the microboCfI•• peroxl.om. mltoehondr~ size ranges from .1 to . Fig. 6.6. A generalized eukarvotic cell. 1~ (1~ = 0 .001 mm) . . .Organisms like pleuropneumonia, so called "elementary bodies" have been observed with diameters of 100~m. These appear to be a resting form of the bacteria which may grow into bodies of 250 ~m in diameter during the' active metabolism of the organism. It would seem that 20o-250~m in diameter is the lower limit for the size of inactive, living cells. This lower limit may be set by minimum number and size of the component necessary for independent cellular existance. Usually the vast majority of cells lies in the range of 5 to 20~ in diameter. The length of unicellular diatoms is about 100~. And the size of Amoeba proteus is 1mm (1000~) in length. The size of human R-B-C is 7-8~m diameter. In human beings, the volume of nerve cells varies from 200~3 and 15000~3 . The long nerve is also an example of the extremes in cell sizes, the range of which is quite remarkable in the living world . .

A

TYPICAL CEll.

Under the ordinary microscope, only few cell organelles like mitochondria, Golgi bodies, chloroplasts and nucleus are visible . However, under the electron microscope, several other cytoplasmic organelles such as ribosomes, Iysosomes, endoplasmic reticulum, nuclear membrane, plasma membrane etc. are

seen. Some important cellular components are being described here under the following heads :

Plasma Membrane The outermost boundary of the animal cell is called "plasma membrane". It is invisible under a light microscope but under the electron microscope, it appears to be composed of two dense layers which are called the outer dense layer and inner dense layer. Both these layers are about 20A.

71

Cell as a Unit of Life

in thickness, and are separated from each other by a less dense area of about 35A, in thickness. By this way the total thickness of the plasma membrane is about 75A. Before the advance of the electron microscope, people used to think of plasma membrane as been stretched tightly over the cell. But now we know that single plasma membrane can have as many as 3000 microvilli. Another peculiarity of the structure of the membrane is that, it remains connected with the endoplasmiC reticulum.

Fig. 6.7. Structures typically seen in thin·sectioned plant cells.

The membrane itself may be perforated with tiny holes through which certain materials may cross. Also, as we will see shortly, materials actually move across the membranous material itself. There may also be other kinds of molecules associated with the membrane. For example, various sorts of carbohydrates may be found attached to the outer side of the membrane, specific carbohydrates that determine the cell type. The underside of the membrane may be attached to a sort of internal support (skeleton) for the cell, called the microtrabecular lattice. In a word, plasma membranes are more than fluid filled sacs. They are living responding structures. Thus, the more membrane area an organism possesses, the greater its control over its internal environment.

Biomolecules

72

The microtrabecular lattice It has been recently discovered that the cell organelles may not float freely in an amorphous cytoplasm. Instead, they are held in place by a complex bridge .work called the microtrabecular lattice, a web-like system of microtubules and microfilaments that forms an internal cellular framework upon which many organelles are suspended. It appears as a maze like network of hollow fibres, extending throughout the cell, connecting and suspending the organelles in a kind of three-dimensional web. Researches are already hard at work unlocking the secrets of this gland network (others are not conceived-it exists). There is also important evidence that the lattice holds even enzymes in place. It has been suggested that precise spatial arrangement of enzymes would increase their efficiency by encouraging a specific sequence of interactions. For example, enzyme B might be held near enzyme A, so that it might more easily interact with the product of A, enzyme C would be near B and so forth. Such structural organizations presumably would be an improvement over random enzyme movement through the cell. Cell Walls Plant cells have cell walls, animal cells do not. Cell walls are non-living, rather inflexible, highly permeable, and strengthened by mats composed of cellulose fibers and other compounds in a tough and complex matrix. T~is is why trees are not limp.

·microtubule Fig. 6.8. Ultrastructure of a typical animal cell as seen In the electron microscope.

Cell as a Unit of Life

73

The cell walls of plants are commercially important in a number of ways. For example, we count cell walls to hold up our own walls as we frame our houses with wood. Also, it is for the cellulose in cell walls that we have revealed. Vast areas of our forests in response to wheedling commercials 'are designed to increase our demand for "disposable" paper commodities. The cellulose of plant cell walls is also valuable as a major component of celluloid, rayon, cotton and hemp (Hemp was once provided by legally cultivating a plant called Cannabis saliua, later known as "killer weed" or reefs). Another component of cell walls, lignin, was long considered a totally useless byproduct of paper manufacturing. However, researchers worked hard to find ways to alter so that it could be sold as the outer border of the nuclear ,membrane. Furthermore, there is evidence that it is confluent with the plasma membrane. The hypothesis is partly based on notions of how the nucleus may have arisen. This arrangement would result in an open channel from the nucleus to the outside of the cell. The cell might thus be able to easily transport products manufactured in the nucleus directly to the outside. But more important, with such a connection, the nucleus might be able to react quickly to changes in the cell's environment.

Golgi Bodies In 1898, Camillo Golgi, an Italian cytologist, was experimenting with some cell-staining procedures and discovered that when he used certain stains such as silver nitrate, "peculiar bodies" appeared in the cells. These structures had never been noticed before, but when other workers looked for them using the same stains, they turned up in a variety of cells, However because these could not be seen in living cells, there was a great argument over whether they really were cell structures, or were just artifiCial or debris produced by the staining process itself. The electron microscope resolved the debate. Indeed, these strange bodies did exist, and appropriately enough they were named Golgi bodies. It was found that they had a characteristic and identifiable structure no matter what kind of cell they were found in. In every case, they appeared as a group of tiny flattened vesicles, lying roughly parallel to each other, somewhat like pancakes. Even after their existence was confirmed, an argument continued over their function. What did they do for a living? We now know that they serve as a sort of packaging center of the cell. They have been linked to manufacturing, warehOusing and shipping centers (as well as swiss finishing schools). Their role is indeed complex. Products formed by the endoplasmic reticulum are stored (and in some cases modified) in the Golgi complex. The complex also manufactures many polysaccharides including ones that will be secreted by the cell. Enzymes and other proteins, as well as certain carbohydrates, are collected in these bodies and packaged into sacs or vesicles. In this way, they are kept apart from the rest of the cell. In some cases, the packages break away from the Golgi complex and move to the plasma membrane where the enclosed molecules are excreted from the cell. Lysosomes Lysosomes are somewhat spherical cytoplasmic organelles and are in general, d~1:inctly unimpressive bodies. It is believed that lysosomes are packets of digestive enzymes that are syntheSized by the cell and packaged by the Golgi bodies. The packaging is important because if these enzymes were floating free in the cell's cytoplasm, the cell itself would be digested. Christian de Duva, who discovered the lysosomes, called them "suicide bags", and the dramatic description is not entirely univarrented, since they can actually destroy the cell that bears them. So why would cells have ever developed such a risk to themselves? In some cases, the destruction of cells is beneficial to the organism. For example, the cells could be old and not functioning well, or they might be in a part of the body that was undergoing reduction as a part of a normal developmental process, such as in the webbing between the fingers of a developing embryo. Lysosomes may also help dispose off unwanted mitochondria, red blood cells, or bacteria (fragments of all these have been found within the organelles). Interestingly, malfunctioning lysosomes have been associated with a number of human diseases, including cancer. n.upturing lysosomes have also been accused of contributing to the ageing process,

74

Biomolecules

Mitochondria outer membrane perimitochondrial space These may be filamentous matrix cristae or granular structures. The size changes, depending on the physiological conditions of the cells. Mitochondria are usually .51l to III in diameter and range in length from 21l to 71l. The number in cells, such as liver cells mitochondria may be over a thousand. Electron microscope has revealed the basic structure of the mitochondria. Each mitochondrion is enclosed in a double membrane. The outer membrane forms an uninterrupted boundary but the inner Fig. 6.9. Three dimerlllollai structure of mitochondrion. membrane is continuously extended into folds, which project into the inner space of the mitochondrion. The inner folds are known as cristae. The inner space of the mitochondrion is filled with a fluid, which is rich in enzymes. Other enzymes are found in the membranes. The mitochondrial membranes are similar to those of the cell membranes. They are also made up of double layers of phospholipid molecules sandwiched between layers of protein molecules. Similar to cell membranes, mitochondrial membranes can also expand and contract. Mitochondria render oxidation of food substances. They change the potential energy of different food materials into a form of energy, which can be used by the cell for its various activities. It is, therefore, regarded that mitochondria may gather in most regions of the cells. Due to their functions, mitochondria are also called power house of the cell. Endoplasmic Reticulum The term 'microsome' was introduced to the modern cytology by Claude (1943) to represent one of the submicroscopic cellular components isolated by centrifugation. These are now known to be broken parts of the endoplasmic reticulum after the studies of Kollman (1953). Porter (1945) was the first man to describe their electron microscopic strudure Smooth Endoplasmic reticulum in cultured cells. He described them as lace like reticulum. There are membrane bounded sacs in the form of double membrane or cisternae. The term cisternae was introduced by Sjostrand (1953) to describe long and elongated rod-like parts of the endoplasmic reticulum measuring 50-200A. A. Wiess (1953), described the vesicles of the endoplasmic reticulum Rough which usually has a diameter Endoplasmic reticulum of 25-50oA while Bradfield Fig. 6.10. Two types of endoplasmic reticulum WIth lIucleus. (1913) came across another

Cell as a Unit of Life

75

type of endoplasmic reticulum and termed them as tubules, they measure from SO to 100A. in diameter. On the basis of presence or absence of ribosomes or RNP particles, the endoplasmic reticulum (E.R) is distinguished in two varieties: the rough walled endoplasmic reticulum and smooth walled endoplasmic reticulum. They are also known as granular or a.-~ type of endoplasmic reticulum.

Rough-walled Endoplasmic Reticulum Rough-walled endoplasmic reticulum (R.E.R) is that variety which bears the ribosomes at the external surface of the cisternae. The distribution of the ribosomes can be circular, spiral or rosette type. The particles may be induced to leave the surface of the cisternae or the capacity to associate again with the particles. Usually the rough-walled endoplasmic reticulum are richly distributed in those cells which are engaged in the synthesis of the proteins.

Smooth-walled Endoplasmic Reticulum The other large division of the system owes its identity in parts to the absence of the particles and therefore commonly referred to c