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Foundation of Biochemistry

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  1. 1. Foundation of Biochemistry Course: B.Sc. Biochemistry Sub: introduction to biochemistry Unit 1
  2. 2. • Fifteen to twenty billion years ago, the universe arose as a cataclysmic eruption of hot, energy-rich subatomic particles. Within seconds, the simplest elements (hydrogen and helium) were formed. As the universe expanded and cooled, material condensed under the influence of gravity to form stars. Some stars became enormous and then exploded as supernovae, releasing the energy needed to fuse simpler atomic nuclei into the more complex elements. Thus were produced, over billions of years, the Earth itself and the chemical elements found on the Earth today. About four billion years ago, life arose—simple microorganisms with the ability to extract energy from organic compounds or from sunlight, which they used to make a vast array of more complex biomolecules from the simple elements and compounds on the Earth’s surface.
  3. 3. Biochemistry asks how the remarkable properties of living organisms arise from the thousands of different lifeless biomolecules. When these molecules are isolated and examined individually, they conform to all the physical and chemical laws that describe the behavior of inanimate matter—as do all the processes occurring in living organisms. The study of biochemistry shows how the collections of inanimate molecules that constitute living organisms interact to maintain and perpetuate life animated solely by the physical and chemical laws that govern the nonliving universe.
  4. 4. • A high degree of chemical complexity and microscopic organization. Thousands of different molecules make up a cell’s intricate internal structures. Each has its characteristic sequence of subunits, its unique three- dimensional structure, and its highly specific selection of binding partners in the cell. • Systems for extracting, transforming, and using energy from the environment, enabling organisms to build and maintain their intricate structures and to do mechanical, chemical, osmotic, and electrical work. Inanimate matter tends, rather, to decay toward a more disordered state, to come to equilibrium with its surroundings.
  5. 5. • A capacity for precise self-replication and self-assembly. A single bacterial cell placed in a sterile nutrient medium can give rise to a billion identical “daughter” cells in 24 hours. • Each cell contains thousands of different molecules, some extremely complex; yet each bacterium is a faithful copy of the original, its construction directed entirely from information contained within the genetic material of the original cell. • Mechanisms for sensing and responding to alterations in their surroundings, constantly adjusting to these changes by adapting their internal chemistry. • Defined functions for each of their components and regulated interactions among them.
  6. 6. A history of evolutionary change • Organisms change their inherited life strategies to survive in new circumstances. The result of eons of evolution is an enormous diversity of life forms, superficially very different but fundamentally related through their shared ancestry. • Despite these common properties, and the fundamental unity of life they reveal, very few generalizations about living organisms are absolutely correct for every organism under every condition; there is enormous diversity. • The range of habitats in which organisms live, from hot springs to Arctic tundra, from animal intestines to college dormitories, is matched by a correspondingly wide range of specific biochemical adaptations, achieved within a common chemical framework.
  7. 7. • Biochemistry describes in molecular terms the structures, mechanisms, and chemical processes shared by all organisms and provides organizing principles that underlie life in all its diverse forms, principles we refer to collectively as the molecular logic of life. Although biochemistry provides important insights and practical applications in medicine, agriculture, nutrition, and industry, its ultimate concern is with the wonder of life itself.
  8. 8. Cellular Foundations • The unity and diversity of organisms become apparent even at the cellular level. The smallest organisms consist of single cells and are microscopic. Larger, multicellular organisms contain many different types of cells, which vary in size, shape, and specialized function. Despite these obvious differences, all cells of the simplest and most complex organisms share certain fundamental properties, which can be seen at the biochemical level.
  9. 9. • Cells Are the Structural and Functional Units of All Living Organisms Cells of all kinds share certain structural features. The plasma membrane defines the periphery of the cell, separating its contents from the surroundings. It is composed of lipid and protein molecules that form a thin, tough, pliable, hydrophobic barrier around the cell. The membrane is a barrier to the free passage of inorganic ions and most other charged or polar compounds. • Transport proteins in the plasma membrane allow the passage of certain ions and molecules; receptor proteins transmit signals into the cell; and membrane enzymes participate in some reaction pathways. Because the individual lipids and proteins of the plasma membrane are not covalently linked, the entire structure is remarkably flexible, allowing changes in the shape and size of the cell. As a cell grows, newly made lipid and protein molecules are inserted into its plasma membrane; cell division produces two cells, each with its own membrane. This growth and cell division (fission) occurs without loss of membrane integrity.
  10. 10. • All cells have, for at least some part of their life, either a nucleus or a nucleoid, in which the genome the complete set of genes, composed of DNA—is stored and replicated. The nucleoid, in bacteria, is not separated from the cytoplasm by a membrane; the nucleus, in higher organisms, consists of nuclear material enclosed within a double membrane, the nuclear envelope. • Cells with nuclear envelopes are called eukaryotes (Greek eu, “true,” and karyon, “nucleus”); those without nuclear envelopes—bacterial cells—are prokaryotes (Greek pro, “before”).
  11. 11. • There Are Three Distinct Domains of Life • All living organisms fall into one of three large groups (kingdoms, or domains) that define three branches of evolution from a common progenitor. Two large groups of prokaryotes can be distinguished on biochemical grounds: archaebacteria (Greek arche-, “origin”) and eubacteria (again, from Greek eu, “true”). • Eubacteria inhabit soils, surface waters, and the tissues of other living or decaying organisms. Most of the well studied bacteria, including Escherichia coli, are eubacteria. The archaebacteria, more recently discovered, are less well characterized biochemically; most inhabit extreme environments—salt lakes, hot springs, highly acidic bogs, and the ocean depths.
  12. 12. Three Distinct Domains of Life
  13. 13. • Cells Build Supramolecular Structures • Macromolecules and their monomeric subunits differ greatly in size. A molecule of alanine is less than 0.5 nm long. Hemoglobin, the oxygen-carrying protein of erythrocytes (red blood cells), consists of nearly 600 amino acid subunits in four long chains, folded into globular shapes and associated in a structure 5.5 nm in diameter. In turn, proteins are much smaller than ribosomes (about 20 nm in diameter), which are in turn much smaller than organelles such as mitochondria, typically 1,000 nm in diameter. It is a long jump from simple biomolecules to cellular structures that can be seen with the light microscope.
  14. 14. Structural hierarchy in the molecular organization of cells
  15. 15. Biomolecules Are Compounds of Carbon with a Variety of Functional Groups • The chemistry of living organisms is organized around carbon, which accounts for more than half the dry weight of cells. Carbon can form single bonds with hydrogen atoms, and both single and double bonds with oxygen and nitrogen atoms. Of greatest significance in biology is the ability of carbon atoms to form very stable carbon–carbon single bonds. Each carbon atom can form single bonds with up to four other carbon atoms. Two carbon atoms also can share two (or three) electron pairs, thus forming double (or triple) bonds.
  16. 16.  Because of its bonding versatility, carbon can produce a broad array of carbon–carbon skeletons with a variety of functional groups; these groups give biomolecules their biological and chemical personalities.  A nearly universal set of several hundred small molecules is found in living cells; the interconversions of these molecules in the central metabolic pathways have been conserved in evolution.  Proteins and nucleic acids are linear polymers of simple monomeric subunits; their sequences contain the information that gives each molecule its three-dimensional structure and its biological functions.  Molecular configuration can be changed only by breaking covalent bonds. For a carbon atom with four different substituents (a chiral carbon), the substituent groups can be arranged in two different ways, generating stereoisomers with distinct properties. Only one stereoisomer is biologically active. Molecular conformation is the position of atoms in space that can be changed by rotation about single bonds, without breaking covalent bonds.  Interactions between biological molecules are almost invariably stereospecific: they require a complementary match between the interacting molecules.
  17. 17. WATER
  18. 18. O H H 1 molecule of water is made up of 2 hydrogen atoms bonded with 1 oxygen atom. The molecular formula is H2O STRUCTURE OF WATER
  19. 19. O STRUCTURE OF WATER The bond that forms water is a covalent bond
  20. 20. • Water consists of an oxygen atom bound to two hydrogen atoms by two single covalent bonds. – Oxygen has unpaired & paired electrons which gives it a slightly negative charge while Hydrogen has no unpaired electrons and shares all others with Oxygen – Leaves molecule with positively and negative charged ends Water is a Polar Molecule -has oppositely charged ends
  21. 21. Properties of Water Polar Molecule Cohesion And Adhesion High Specific Heat Density – Greatest At 4oc Universal Solvent Of Life Capillary Action Surface Tension Buoyancy
  22. 22. Polarity of Water • In a water molecule two hydrogen atoms form single polar covalent bonds with an oxygen atom. Gives water more structure than other liquids – Because oxygen is more electronegative, the region around oxygen has a partial negative charge. – The region near the two hydrogen atoms has a partial positive charge. • A water molecule is a polar molecule with opposite ends of the molecule with opposite charges.
  23. 23. 24 slightly positive charge slightly negative charge hydrogen bond between (+) and (-) areas of different water molecules Water molecules form Hydrogen bonds
  24. 24. – Water has a variety of unusual properties because of attractions between these polar molecules. – The slightly negative regions of one molecule are attracted to the slightly positive regions of nearby molecules, forming a hydrogen bond. – Each water molecule can form hydrogen bonds with up to four neighbors.
  25. 25. “Universal” Solvent • A liquid that is a completely homogeneous mixture of two or more substances is called a solution. – A sugar cube in a glass of water will eventually dissolve to form a uniform mixture of sugar and water. • The dissolving agent is the solvent and the substance that is dissolved is the solute. – In our example, water is the solvent and sugar the solute. • In an aqueous solution, water is the solvent. • Water is not really a universal solvent, but it is very versatile because of the polarity of water molecules.
  26. 26. • Water is an effective solvent as it can form hydrogen bonds. – Water clings to polar molecules causing them to be soluble in water. • Hydrophilic - attracted to water – Water tends to exclude non polar molecules. • Hydrophobic - repelled by water
  27. 27. Acids • Acids dissociate in water to increase the concentration of H+. – Have many H+ ions – Sour taste – HCl is hydrochloric acid or stomach acid
  28. 28. Bases • Bases combine with H+ ions when dissolved in water, thus decreasing H+ concentration. – Have many OH- (hydroxide) ions – Bitter taste – NaOH = sodium hydroxide or baking soda
  29. 29. Acids and Bases • An acid is a substance that increases the hydrogen ion concentration in a solution. • Any substance that reduces the hydrogen ion concentration in a solution is a base. – Some bases reduce H+ directly by accepting hydrogen ions. • Strong acids and bases complete dissociate in water. • Weak acids and bases dissociate only partially and reversibly. 1
  31. 31. • The structure of any molecule is a unique and specific aspect of its identity. • Molecular structure reaches its pinnacle in the intricate complexity of biological macromolecules, particularly the proteins. Although proteins are linear sequences of covalently linked amino acids, the course of the protein chain can turn, fold, and coil in the three dimensions of space to establish a specific, highly ordered architecture that is an identifying characteristic of the given protein molecule
  32. 32. Weak Forces Maintain Biological Structure and Determine Biomolecular Interactions • Covalent bonds hold atoms together so that molecules are formed. In contrast, weak chemical forces or noncovalent bonds, (hydrogen bonds, van der Waals forces, ionic interactions, and hydrophobic interactions) are intramolecular or intermolecular attractions between atoms. None of these forces, which typically range from 4 to 30 kJ/mol, are strong enough to bind free atoms together. The average kinetic energy of molecules at 25°C is 2.5 kJ/mol, so the energy of weak forces is only several times greater than the dissociating tendency due to thermal motion of molecules. Thus, these weak forces create interactions that are constantly forming and breaking at physiological temperature, unless by cumulative number they impart stability to the structures generated by their collective action.
  33. 33. •Van der Waals • Van der Waals forces are the result of induced electrical interactions between closely approaching atoms or molecules as their negatively-charged electron clouds fluctuate instantaneously in time. These fluctuations allow attractions to occur between the positively charged nuclei and the electrons of nearby atoms. • Hydrogen Bonds • Hydrogen bonds form between a hydrogen atom covalently bonded to an electronegative atom (such as oxygen or nitrogen) and a second electronegative atom that serves as the hydrogen bond acceptor.
  34. 34. • Ionic Interactions • Ionic interactions are the result of attractive forces between oppositely charged polar functions, such as negative carboxyl groups and positive amino groups. These electrostatic forces average about 20 kJ/mol in aqueous solutions. Typically, the electrical charge is radially distributed, and so these interactions may lack the directionality of hydrogen bonds or the precise fit of van der Waals interactions.
  36. 36. Atom – the smallest unit of matter “indivisible” Helium atom
  37. 37. electron shells a) Atomic number = number of Electrons b) Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. c) Electron shells determine how an atom behaves when it encounters other atoms
  38. 38. Electrons are placed in shells according to rules: 1) The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.
  39. 39. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to N would like to O would like to Gain 4 electrons Gain 3 electrons Gain 2 electrons
  40. 40. Why are electrons important? 1) Elements have different electron configurations  different electron configurations mean different levels of bonding
  41. 41. 3
  42. 42. Electron Dot Structures Symbols of atoms with dots to represent the valence-shell electrons 1 2 13 14 15 16 17 18 H He:            Li Be  B   C   N   O  : F  :Ne :                    Na Mg  Al  Si  P S :Cl  :Ar :        
  43. 43. Chemical bonds: an attempt to fill electron shells 1. Ionic bonds – 2. Covalent bonds – 3. Metallic bonds
  44. 44. IONIC BOND bond formed between two ions by the transfer of electrons
  45. 45. Formation of Ions from Metals  Ionic compounds result when metals react with nonmetals  Metals lose electrons to match the number of valence electrons of their nearest noble gas  Positive ions form when the number of electrons are less than the number of protons Group 1 metals  ion 1+ Group 2 metals  ion 2+ • Group 13 metals  ion 3+
  46. 46. Formation of Sodium Ion Sodium atom Sodium ion Na  – e  Na + 2-8-1 2-8 ( = Ne) 11 p+ 11 p+ 11 e- 10 e- 0 1+
  47. 47. Formation of Magnesium Ion Magnesium atom Magnesium ion  Mg  – 2e  Mg2+ 2-8-2 2-8 (=Ne) 12 p+ 12 p+ 12 e- 10 e- 0 2+
  48. 48. Some Typical Ions with Positive Charges (Cations) Group 1 Group 2 Group 13 H+ Mg2+ Al3+ Li+ Ca2+ Na+ Sr2+ K+ Ba2+
  49. 49. Ions from Nonmetal Ions In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals Nonmetal add electrons to achieve the octet arrangement Nonmetal ionic charge: 3-, 2-, or 1-
  50. 50. Fluoride Ion unpaired electron octet     1 - : F  + e : F :     2-7 2-8 (= Ne) 9 p+ 9 p+ 9 e- 10 e- 0 1 - ionic charge
  51. 51. Ionic Bond • Between atoms of metals and nonmetals with very different electronegativity • Bond formed by transfer of electrons • Produce charged ions all states. Conductors and have high melting point. • Examples; NaCl, CaCl2, K2O
  52. 52. Ionic Bonds: One Big Greedy Thief Dog!
  53. 53. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions. 4
  54. 54. 4
  55. 55. COVALENT BOND bond formed by the sharing of electrons
  56. 56. Covalent Bond • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing particles, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O, SiC
  57. 57. 5
  58. 58. Bonds in all the polyatomic ions and diatomics are all covalent bonds
  59. 59. when electrons are shared equally NONPOLAR COVALENT BONDS H2 or Cl2
  60. 60. when electrons are shared but shared unequally POLAR COVALENT BONDS H2O
  61. 61. Polar Covalent Bonds: Unevenly matched, but willing to share.
  62. 62. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen. 6
  63. 63. METALLIC BOND bond found in metals; holds metal atoms together very strongly
  64. 64. Metallic Bond • Formed between atoms of metallic elements • Electron cloud around atoms • Good conductors at all states, lustrous, very high melting points • Examples; Na, Fe, Al, Au, Co
  65. 65. Metallic Bonds: Mellow dogs with plenty of bones to go around.
  66. 66. Ionic Bond, A Sea of Electrons 7
  67. 67. Metals Form Alloys Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter.
  68. 68. Formula Weights • Formula weight is the sum of the atomic masses. • Example- CO2 • Mass, C + O + O 12.011 + 15.994 + 15.994 43.999
  69. 69. References/Sources • All images are from Lehninger Principles of biochemistry by Nelson and Cox except 1. iaSO4qwZ86u7YrskOad2mmojw=s85 2. Y8ayo7ilfE1gUg4wia77gXThu05L6zCZJohN_Qvb1nQ=s93 3. wePD6IGe1T7ayJ8keL78slq_JCPeEKlgcUA=s139 4. nKRoWVEcQ=s16 5. zVQyHNfbhJ1tMVtMTDTrP9HG5MKOBKZk_OC49ioWw=s116 6. mkRU14yNN0viCE7qqp5TJQi4yC2=s88 7. caEj81y24_fc6U1F2g_1uXaE8ZhDzc=s125 Books/ Web resources • Lehninger Principles of biochemistry by Nelson and Cox •