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Lecture 1 1 Chemistry of life Introduc5on to biochemistry 2 4 Types of Biomolecules 1-‐ Amino Acids Polypep5de or protein Lectures 2, 3 and 4 3 4 Types of Biomolecules 2-‐ Nucleo5des Lecture 5 Nucleic Acids 4 4 Types of Biomolecules 3-‐ Carbohydrates (monosaccarides) Disaccharide Polysaccharide (CH2O)n n>3 Lecture 5 5 4 Types of Biomolecules 4-‐ Lipids Hydrocarbon based Lecture 6 6 Lectures 7-‐11 Cellular Energy 7 Diabetes – Lecture 12 8 Week 1 – what are we going to talk about today? Stereochemistry Thermodynamics Water and Hydrogen bonding Acid-‐Base chemistry 9 Stereochemistry 10 6 simple elements combine to form molecules essen5al to an organism’s existence and func5on Carbon Hydrogen Nitrogen Oxygen Phosphorus Sulfur ~ 97% of composi5on of living organisms 11 Each of the macromolecules are organized around which element? 1-‐ Hydrogen 2-‐ Oxygen 3-‐ Carbon 4-‐ Nitrogen 12 Why are cellular organisms carbon based? Bonding versa5lity 13 Versa5lity of Carbon Linear – Branched -‐ Cyclic 14 Only single bonds can rotate freely – Conforma5on Ethane Freedom of Rota5on around single bonds allows for different conforma5ons Different conforma5ons are freely interconvertable 15 No Rota5on around double bonds Configura5on Ethylene Restricted rota5on of the double bond produces isomers with 16 different configura5ons Configura5on Configura>on (stereochemistry): FIXED spa5al arrangement of atoms (same bonds, different configura5on) Configura5on is conferred by 1-‐ Double bonds Geometric Isomers/cis-‐trans isomers 2-‐ Chiral centers Enan5omers Diasteromers 17 Geometric isomers or cis-‐trans isomers Differ in arrangement of groups with respect to a non-‐rota5ng double bond Geometric isomers cannot be interconverted without breaking double bonds **requires energy 18 Nonsuperimposable Chiral Enantomer Superimposable Achiral Iden5cal Structures -‐An atom agached to 4 different groups is a chiral center (asymmetric C) -‐A molecule is chiral if it cannot be superimposed on its mirror image -‐All chiral molecules contain at least one chiral center 19 Stereochemistry due to chiral centers 4 different subs5tuents bonded to a carbon may be arranged in 2 different ways in space (2 configura5ons) Enan5omers: 2 stereoisomers that are mirror images Diastereomers: 2 stereoisomers that are not mirror images Enan5omers (mirror image) Diastereomers (non-‐mirror image) Iden5cal chemical proper5es – very different biological proper5es 20 How many chiral carbons are there in D-‐Ribulose? n chiral carbons gives 2n stereoisomers 21 22 23 Stereoisomers of Modern Drugs 24 Common func5onal groups found in biochemistry CH3 Methyl Carbonyl Phosphoester Hydroxyl Suljydryl Carboxyl Phosphoryl Amino Ester Phosphoanhydride 25 Energy and Thermodynamics Study of heat and power 26 A living organism is an open system Exchanges both mager and energy with its surroundings Organisms derive energy from the surroundings Take up chemical fuels (glucose) or Total energy Absorb energy from sunlight remains constant Organisms convert energy to produce work Entropy or disorder increases Return some energy to the surroundings as heat Release end product molecules that are less well organized than the star5ng fuel. 27 2 Laws of Thermodynamics First Law : Conserva5on of energy The energy in the universe remains constant Energy cannot be created or destroyed but can change forms Example: A river flowing over a dam: energy is harnessed as electricity and used to produce heat or perform work. 28 2 Laws of Thermodynamics Second law: The entropy of the universe increases Entropy: a measure of a system’s disorder or randomness 29 2 Laws of Thermodynamics Implica5ons for biochemistry: Law 1-‐ naturally occurring (spontaneous) processes will always proceed towards the state with the least poten5al energy (releases energy) Law 2-‐ naturally occurring (spontaneous) processes must increase the disorder in the universe Thermodynamics tell us if a reac>on is possible and whether it will occur spontaneously 30 Synthesis of Macromolecules Spontaneous or not? Building organisms require energy!! 31 Thermodynamics and Gibbs free energy Gibbs free energy (G): usable energy content of a biological system – Joules per mol (J.mol-‐1) Gibbs free energy Entropy (S) Enthalpy (H) -‐heat content of the system J·∙mol-‐1 Haner – Hbefore = ∆H < 0 Release heat to surroundings or absorb heat from the surrounding? -‐ measure of disorder J·∙K-‐1·∙mol-‐1 Saner – Sbefore = ∆S > 0 Increase or decrease entropy? Favorable or not favorable? 32 Free energy is a measure of the enthalpy and entropy G = H-‐T S ? Temperature in Kelvin For a process to occur, the overall change in free energy must be nega5ve Free energy of the products must be less than the free energy of the reactants G = Gproducts – Greactants < 0 33 Does a process occur spontaneously or not? G > 0 = not spontaneous: Endergonic -‐ requires input of energy G < 0 = spontaneous: Exergonic -‐ release of free energy G = 0 = equilibrium: No net change in the energy of the system ∆G does NOT indicate how fast the reac5on will proceed (rate) -‐ only whether it will occur! 34 Effects of H and S on spontaneity ∆H G = H-‐T S ∆G = ∆H-‐ T∆S ∆S -‐ + release of heat; increase in entropy -‐ -‐ spontaneous at all temps -‐ -‐ release of heat; decrease in entropy – + -‐ spontaneity depends on temp + increase in heat; increase in entropy; -‐spontaneity depends on temp + -‐ increase in heat; decrease in entropy; -‐unfavorable at all temps. 35 Let’s consider Gibbs Free Energy in terms of chemical reac>ons The entropy of a substance increases with its volume. example: gas molecules Entropy is therefore a func5on of concentra5on If entropy changes with concentra5on, so must free energy ** The free energy change of a chemical reac5on depends on the concentra5ons of reactants and products 36 Change in free energy is related to the concentra5ons of the reac5ng substances A + B C + D When a reac5on is at equilibrium Equilibrium constant Keq = [C]eq[D]eq [A]eq[B]eq Equilibrium – rate of product forma5on equals the rate at which product is converted to reactant Equilibrium -‐ No net change in [ ] of reactants Equilibrium DOES NOT mean that the [ ] of reactants and products are = 37 A + B C + D Keq = [C]eq[D]eq [A]eq[B]eq What does a large (>>1) Keq mean? A-‐ The concentra5on of products and reactants are equal at equilibrium B-‐ The forma5on of products is favored at equilibrium C-‐ The forma5on of reactants is favored at equilibrium 38 Chemical equilibria and free energy A reac5on’s free energy change depends on 2 things: 1) A constant term dependent only on the reac5on itself 2) A variable term dependent on the concentra5on of reactants and products Constant Term ΔG°’ -‐ Standard Free energy change of a reac5on under standard condi5ons. Concentra5ons = 1.0 M Temp = 25°C Atmospheric pressure = 1.0 atm What does ΔG°’ tell us? Which direc5on and how far a reac5on must go to reach equilibrium when ini5al concentra5ons of each component is 1.0 M 39 Standard free energy is related to the equilibrium constant ∆G°’ = -‐ RT ln Keq R = gas constant (8.315 J/mol·∙K) T = temperature (Kelvin) 40 But -‐ biochemical rxns don’t occur at standard state. We need to factor in the variable components C + D A + B ∆G = ∆G°’ + RT ln [C]c [D]d [A]a[B]b R = gas constant (8.315 J/mol·∙K) T = temperature (Kelvin) ∆G is a func5on of the actual concentra5ons of reactants and the temperature ∆G is a measure of the distance from equilibrium 41 Many of life’s chemical reac5ons are unfavorable; How do we drive these reac5ons to occur? 1 -‐ You can make an unfavorable reac5on favorable by adjus5ng the concentra5ons of reactants and products Dihydroxyacetone (DHAP) Glyceraldehyde 3-‐ phosphate (GAP) Phosphate -‐Reac5on is endergonic (∆G° ‘ = 7.5 kJ/mol) -‐DHAP will not spontaneously convert to GAP When [GAP] decreases DHAP will spontaneously convert to GAP [DHAP] [GAP] Le Chatelier’s Principle 42 Many of life’s chemical reac5ons are unfavorable; How do we drive these reac5ons to occur? 2 -‐ We couple endergonic reac5ons with exergonic reac5ons Unfavorable Glucose + Pi → Glucose 6-‐Phosphate Favorable ∆G = + 13.8 kJ·∙M-‐1 ATP → ADP + Pi ∆G = -‐ 30.5 kJ·∙M-‐1 Glucose + ATP → Glucose 6-‐P + ADP ∆G = -‐ 16.7 kJ·∙M-‐1 Reac5on is now favorable Cells can enable unfavorable reac5ons by coupling them with favorable reac5ons (e.g., ATP hydrolysis) 43 Real life example How do we make large ordered proteins from individual amino acids? Amino Acids → Proteins ∆G = posi5ve (endergonic) ATP → AMP + P -‐ P ∆G = nega5ve (exergonic) By coupling thermodynamically unfavorable processes with sources of high energy, cells are able to synthesize large polymers!! 44 Living organisms and steady state We con5nuously ingest high enthalpy, low entropy nutrients and convert them to low enthalpy, high-‐entropy waste products Living organisms only come to equilibrium when this process is disrupted Equilibrium = Death However, we do maintain a steady state or homeostasis -‐ Proper5es of the cell are maintained over 5me. The system reacts to changes in these proper5es to restore homeostasis 45 Chemistry of Water Study of an aqueous environment 46 Aqueous solu5ons and proper5es of water: 1-‐ Organisms are mostly water (human 70%) 2-‐ Influences the shape and func5on of biomolecules 3-‐ Medium of most biochemical reac5ons 4-‐ Important for transport of nutrients and waste 5-‐ Water may itself par5cipate in chemical reac5ons 47 Structure of H20 O forms covalent bonds with 2 H atoms H2O has 2 unshared pairs of electrons-‐ Tetrahedral shape O is more electronega5ve than the H’s (O has a stronger agrac5on for the electrons) Unequal electron sharing between the O and the H’s This creates a dipole or polarity where the O has a par5al -‐ charge and the H’s have a par5al + charge 48 Hydrogen Bonding H-‐bonds result from agrac5on between the O atom of 1 H2O molecule and the H atom of another water molecule H-‐ bonds result in cohesiveness and high surface tension of 49 water Structure of Ice -‐ H2O par5cipates in 4 hydrogen bonds Accepts H bonds Donates to H bonds Regular la•ce-‐like structure 50 Hydrogen bonding in H2O depends on the phase No H bonds ~ 3.4 H bonds 4 H bonds H20 molecules are in con5nuous mo5on – H-‐ bonds are randomly breaking and forming At room temperature why does ice melt spontaneously? 51 What would ∆S be? What would ∆G be? Hydrogen bonds are not unique to water N ·∙·∙·∙ H Common H bonds O ·∙·∙·∙ H C – H bonds do not form hydrogen bonds 52 Other types of bonds in biochemistry Noncovalent Covalent: Atoms share electrons (London Dispersion forces) Rela5ve strengths of different bonds Non covalent interac5ons are individually weak but strong in large numbers! 54 Polar – hydrophilic – soluble in water – charged or able to H bond Nonpolar – hydrophobic – insoluble in water Amphiphilic – contains both polar and nonpolar components 55 BP Oil Spill A. Polar B. Nonpolar C. Amphiphilic 56 A. Polar B. Nonpolar C. Amphiphilic 57 Hydrophobic effect What happens when a nonpolar substance is added to an aqueous solu5on? 58 Polar head Ordered H2O forms cages around hydrophobic por5on Hydrophobic effect Hydrophobic tails cluster: less ordered water Hydrophobic tails sequester from water: Ordered water is minimal Entropy Entropy increases as H20 becomes less ordered 59 Remember Entropy? Entropy: a measure of a system’s disorder or randomness 2nd Law of Thermodynamics The total entropy of the universe is always increasing 60 The hydrophobic effect is very important in protein folding The hydrophobic effect drives the nonpolar components to cluster away from the water and effec5vely increases the amount of disordered water 61 Acid Base Chemistry 62 Why talk about acids and bases? Many biological molecules have functional groups that undergo acid-base reactions; therefore, the properties of these molecules are affected by acidity of the solutions in which they are surrounded. pH affects the structure and activity of biological molecules 63 Water ionizes to yield a hydrogen ion (proton) and a hydroxide ion H20 Equilibrium constant for ioniza5on of water H⁺ + OH⁻ Keq = [H⁺][OH⁻] [H2O] Keq[H2O] = [H⁺][OH-‐] Kw = [H⁺][OH-‐] 10-‐14 = [H⁺][OH-‐] Keq[H2O] is redefined as Kw; Kw = ioniza5on constant of water; Kw = 10 -‐14 (determined by conduc5vity measurements) The product of [H+] and [OH-‐] in any solu5on must be equal to 10 -‐14 64 The product of [H+] and [OH-‐] in any solu5on must be equal to 10 -‐14 Changes in the [H+] are balanced by changes in the [OH-‐] Neutral solu5on= [H⁺] = [OH-‐] = 10 -‐7 (pure water) Acidic solu5on= [H⁺] > 10 -‐7 Basic solu5on = [H+] < 10 -‐7 65 pH is defined by the concentra5on of hydrogen ions pH = -‐log [H+] [H+] < 10 -‐7 [H⁺] = [OH-‐] = 10 -‐7 [H⁺] > 10 -‐7 66 Acids and Bases -‐ Act to alter the pH of a solu5on Acid – proton donors Base – proton acceptor Strong acids and bases – ionize completely Example: HCL + H20 H+ + Cl-‐ HCl donates a proton (H+) increasing the [H+] of the solu5on = decreases pH 67 Weak acid and bases -‐ubiquitous in biological systems Weak acids and bases do not dissociate completely in H2O Tendency of an acid to lose its proton is defined by Keq or Ka HA H+ + A-‐ Acid Conjugate base Keq = [H+][A-‐] = K a [HA] Acid dissocia1on constant pKa = -‐ log Ka The larger the acid’s Ka, the smaller it’s pKa, and the stronger the tendency to lose a proton-‐ The stronger the acid 68 Determining the pKa of an acid – Acid 5tra5on curve Ace5c acid solu5on 1-‐ Prior to 5tra5on all the acid is in protonated form (HA) 2-‐ Added base causes protons to dissociate from the acid un5l all the acid is in its conjugate base [A-‐]. OH-‐ equivalents added (frac5on of total NaOH required to convert all acid to deprotonated form) 3-‐ At the midpoint [HA] = [A-‐] pH = pKa B uffering 4-‐ occurs one pH unit above and below the pKa 69 pH Buffers Buffer -‐ Aqueous systems that resist changes in pH when small amounts of acids (H+) or bases (OH-‐) are added Buffering region – flat zone around the pKa of a solu5on Strong acid Why are buffers important in biological systems? Various aspects of biological systems work best at well-‐defined pHs Buffers keep pHs within well-‐defined ranges 70 Enzymes have op5mum pH’s – If enzymes don’t func5on -‐ reac5ons slow and the organism dies 71 Rela5ng pH, pKa, and buffer concentra5on Henderson – Hasselbalch equa5on Relates the pH of a solu5on to the pKa of an acid and the [HA] and [A-‐] pH = pKa + log [A-‐] [HA] Prac1cal way to predict the pH of a solu1on when given the [acid] and [conjugate base] Or to calculate the [acid] and [conjugate base] at a given pH. 72 Biological Buffering Bicarbonate system – maintains blood pH Bicarbonate Carbonic acid 73 H+ + HCO-‐ 3 H2CO3 H2O + CO2 What happens in metabolic acidosis? Blood pH decreases due to an accumula1on of H+ How does the body agempt to quickly adjust for the acidosis? Increased ven1la1on: “blowing off” CO2 How does losing CO2 decrease the amount of H+? ShiJs the equilibrium of the equa1on towards making CO2 Which decreases the amount of H+ H+ + HCO-‐ 3 H2CO3 H2O + CO2 Le Chatelier's principle 74 Next week Amino Acids Proteins 75
