‹#› 1 Energy in chemical reactions Bioenergetics Kinetics Redox reactions © Department of Biochemistry, Jiří Dostál, 2010 ‹#› 2 Basic terms: system and energy •A portion of universe separated from surroundings •insulated – no communication with surroundings •closed – exchange of energy •open – exchange of energy, matter, and information • •Energy is the capacity of a system to do work •system = object, reaction mixture, cell, organism •Unit - joule (J), dimension J = kg m2/s2 • ‹#› 3 • the sum of all forms of energy in the system • the energy of all particles - atoms, molecules, ions • the total energy U cannot be determined • only the change (DU ) during some event in the system Internal energy (U) DU = U2 – U1 = Ufinal – Uinitial DU > 0 energy of system increases DU < 0 energy decreases DU = DQ + DW heat work ‹#› 4 Work Heat •they both describe energy in process (the change of energy) •heat (thermal energy) is associated with random motion of atoms and molecules, it is the transfer of energy between two bodies that are at different temperatures •work – useful energy, can be transformed to other forms of energy, and also completely into heat •the opposite is not possible Þ therefore heat is considered to be less utilizable form of energy (energy waste) yes no ‹#› 5 Exchange of heat in the human body Exogeneous heat intake: hot drinks, sun exposure ... Endogeneous heat production: meal, physical exercises Heat output: • radiation • conduction + convection (facilitated by wind) • sweat evaporation •human body cannot utilize heat, we need chemical energy of nutrients •heat just contributes to keeping body temperature •heat intake has limitations ‹#› 6 Heat output in various situations Output way Conditions (temperature, humidity) 20 °C / dry 30 °C / dry 36 °C / dry 36 °C / humid Radiation 61 % 46 % 0 % 0 % Conduction 26 % 27 % 0 % 0 % Sweat evap. 13 % 27 % 100 % 0 % Thermodynamically impossible heat transfer problem, body cannot remove heat, hyperthermia may be fatal. ‹#› 7 Hyperthermia (heat stroke) •more heat is accumulated in body than released, body temperature increases, •intensive exercize/work in hot humid weather, dehydration aggravates the condition •symptoms vary from headache, tachycardia to coma •treatment: careful and slow cooling the body Hypothermia •excessive heat loss, decrease of body temperature below 35 ºC •stay in cold wheather/water, wet clothing in wind etc. •aggravating factors: older people, hypothyreosis, alcohol •treatment: different rewarming methods ‹#› 8 Enthalpy corresponds to reaction heat at constant external pressure DH = H2 - H1 = DQp It expresses the difference in the bond energy of reaction products and reactants. DH < 0 exothermic reaction, the enthalpy of the reaction products is lower (the bonds are more stable) than that of the reactants DH > 0 endothermic reaction, the enthalpy of the products is higher than that of reactants. ‹#› 9 Enthalpy changes in specific processes •Heat of formation • elements ® 1 mol of compound + heat • •Heat of combustion • 1 mol of subst. + excess O2 ® combustion products + heat • •Heat of neutralization acid + hydroxide ® salt + water + heat • •Heat of solution substance(s) + water ® solution ± heat • •Heat of dilution concentrated acid + water ® diluted acid + heat Always pour acid into water, not the reverse way !!! ‹#› 10 Instant Cold/Hot Packs •Based on heat of solution •Consist of a pouch of water and a dry chemical •Striking the pack leads to the pouch breaking, dissolving the chemical and enthalpy change: (up to 80 °C or down to 0 °C, 20 min) •First-aid device at sports events etc. •CaCl2(s) ® Ca2+(aq) + 2 Cl-(aq) ΔHsoln < 0 hot pack •NH4NO3(s) ® NH4+(aq) + NO3-(aq) ΔHsoln > 0 cold pack ‹#› 11 Entropy is a thermodynamic property, a measure of disorder. It is defined as the amount of energy (heat) in the system that cannot be transformed to work: S = Qrev / T . A change in entropy is ΔS = S2 – S1 . in closed and open systems the entropy can either increase or decrease. Entropy ‹#› 12 Entropy increases Entropy decreases •Increasing temperature •Melting a solid •Evaporating a liquid •Dissolving a solid in a liquid •Mixing two substances •Increasing the number of particles during a reaction (decomposition) •Decreasing temperature •Freezing a liquid •Condensing a gas •Precipitating a solid from sol. •Separating two substances •Decreasing the number of particles during a reaction (synthesis) ‹#› 13 Hydrophobic interactions in aqueous environment hydrophobic int. two non-polar molecules in aqueous solution two non-polar molecules adhered to each other Hydrophobic interactions increase the total entropy of a system Water molecules are organized (cluster) – they have low entropy Six water molecules are free – they have high entropy ‹#› 14 Spontaneous processes are irreversible spontaneous energy They can happen without any continuing outside influence. ordered state disordered state ‹#› 15 Principal question of thermodynamics: Will a reaction proceed? YES / NO Neither enthalpy nor entropy can answer ‹#› 16 In closed and open systems, the driving force of chemical reactions or physical changes is the free energy change ΔG The Gibbs (free) energy G of a system is the energy that is available to do useful work as the result of chemical or physical change at constant temperature and pressure. The free energy change ΔG (= G2 – G1) is defined as DG = DH – T DS useful work heat lost or absorbed due to entropy change enthalpy change (reaction heat) ‹#› 17 DG is the criterion of process feasibility DG < 0 exergonic reaction prone to proceed spontaneously DG > 0 endergonic reaction that cannot proceed spontaneously under given circumstances (the reverse reaction is spontaneous) DG = 0 the system is at the equilibrium state DH the heat of reaction DH < 0 exothermic reaction DH > 0 endothermic reaction DS the entropy change DS > 0 the final state is very probable DS < 0 a very low probability of the final state Overview of thermodynamic functions ‹#› 18 What is the driving force of spontaneous processes? 1.to get lowest energy 2.to get maximal disorder Examples of spontaneous processes: • waterfall flows from up to down • heat is tranferred from warmer object to cooler one • sugar crystals spontaneously dissolve in water • gas expands into evacuated space • elemental iron slowly gets rusty • Na + H2O ® ½ H2 + NaOH • DG comprizes both aspects ‹#› 19 ΔS ΔH ΔG = ΔH – TΔS positive (decompositions) negative (exothermic) always negative; the reaction is spontaneous, practically irreversible positive (endothermic) as far as TΔS > ΔH (at higher temperatures), the reaction is spontaneous negative (syntheses) negative (exothermic) as far as TΔS < ΔH (at lower temperatures), the reaction becomes more favourable positive (endothermic) positive at all temperatures; the forward reaction cannot be spontaneous (the reverse rection is always spontaneous) spontaneous reactions are: •the entropy increases and heat is released • reaction is endothermic but accompanied with entropy increase • reaction is highly exothermic so that it overpoises entropy decrease ‹#› 20 Standard state of substances Substance Standard state definition Solid (s) Liquid (l) Gas (g) Solution Pure solid at given temperature* Pure liquid at given temperature* Pure gas, p = 1 atm = 101.3 kPa, given temperature* c = 1 mol/l, given T, p = 101.3 kPa, pH = 0.00 Defined by convention, indicated by symbol ° Quite non-physiological conditions *often 298 K = 25 °C ‹#› 21 between the system that exists at the beginning of the process in its standard state (the reaction quotient Q = 1), i.e. all reactants, both reactants and products are of unit activity, in aqueous solution their concentration c = 1 mol l–1 (if H+ is a reactant, then also [H+] = 1, pH = 0), at specified temperature (usually 25 °C equal to 298 K), and atmospheric pressure 101.3 kPa, and the reaching the state with a minimal G value, that is the equilibrium state of the system in which the reactants and products has reached the concentrations corresponding with the equilibrium constant K. Standard Gibbs free energy change ΔG° for a reversible process represents the free energy change In biological systems, the standard state is defined by pH = 7.00; then the free energy changes are marked as ΔG°´ ‹#› 22 ΔG = ΔGº + RT ln [A]ai [B]bi [C]ci [D]di The ΔG of a reaction depends on the particular kind of reaction (expressed by the ΔGº term) and the initial concentrations of reactants and products (expressed by the second term equivalent to Q). The relationship between free energy and equilibrium for a reaction aA + bB D cC + dD If the equilibrium concentrations are put in as initial ones, the system is in its equilibrium state, DG = 0: 0 = DG° + RT ln K and DG° = – RT ln K ‹#› 23 The expression takes the same form as the equilibrium constant but is used for a initial state, not for a reaction at equilibrium. The initial nonequilibrium concentrations of the substances taking part in reaction a A + b B D c C + d D are used to calculate the reaction quotient Q: [A]ai [B]bi [C]ci [D]di Q = The equilibrium state is described by equilibrium constant K: K = [A]aeq [B]beq [C]ceq [D]deq The value of Q indicates what changes will occur in reaching equilibrium: When Q < K , the reaction has a chance to proceed in the forward direction. When Q > K , the reaction has a chance to proceed in the reverse direction. Equilibrium The general tendency of any spontaneous process is to reach an equilibrium state. ‹#› 24 Bioenergetics: Transformation of energy in the body Living organisms are open systems that have to receive permanently nutrients – compounds of high enthalpy (energy) and low entropy (due to their complex structure). Nutrients are transformed into waste metabolites of low enthalpy and high entropy (simplified structures - CO2, H2O, urea). The part of free energy gained by exergonic breakdown of nutrients drives endergonic reactions and processes (synthesis of complex molecules, performance of mechanical or osmotic work, etc.). The remaining part of acquired energy is released as heat into the surroundings. ‹#› 25 Nutrient Energy (kJ/g) Thermogenesis Lipids 38 4 % Saccharides 17 6 % Proteins 17 30 % Energy data of nutrients Thermogenesis is the heat production 3-5 h after meal. It is expressed in % of the nutrient ingested. Thermogenesis is the consequence of digestion, absorption, and metabolism of nutrient. ‹#› 26 Consider five sugar lumps Combusted in calorimeter sucrose ® CO2 + H2O + heat (100 %) Consumed in a cup of coffee sucrose ® CO2 + H2O + ATP (~70 %) + heat (~30 %) chemical energy content: 5 ´ 2.8 (g) ´ 17 (kJ/g) = 238 kJ sugar In both processes, oxygen excess is a necessary condition. ‹#› 27 Average fuel values of foods (kJ/100 g) Water Cucumbers Coca-Cola Lemons Apples Potatoes Bananas Eggs Beef Ham 0 50 180 100 200 280 400 500 600 700 Bread Roll Lentils Rice Pastry Pork Sugar Chocolate Butter Oils 900 1100 1200 1300 1400 1500 1700 2200 3000 3800 ‹#› 28 chemical energy of nutrients heat NADH+H+ FADH2 proton gradient across IMM heat heat ATP 1 2 3 1 .. metabolic dehydrogenations with NAD+ and FAD 2 .. respiratory chain (oxidation of reduced cofactors + reduction of O2 to H2O) 3 .. oxidative phosphorylation, IMM inner mitochondrial membrane 4 .. transformation of chemical energy of ATP into work + some heat ¢.. high energy systems Energy transformations in the human body are accompanied with continuous production of heat work 4 ‹#› 29 ATP (adenosine triphosphate) is high energy compound N-glycosidic bond ester anhydride at physiological pH (7.4) adenine ribose triphosphate ‹#› 30 O O H O H 2 O O C H O P O P O O O ~ P O O O ~ + H+ H N N N N N H 2 ATP + P O O O H O 2 O O O H O H N N N N N H 2 2 ADP P O O O C H O P O O O ~ Hydrolysis of ATP decreases the number of negative charges ATP4- + H2O ® ADP3- + HPO42- + H+ ‹#› 31 Hydrolysis of ATP is exergonic reaction ATP + H2O ® ADP + Pi + energy DG°´ = -31 kJ/mol ADP + Pi + energy ® ATP + H2O DG°´ = 31 kJ/mol ATP formation is endergonic reaction and requires energy investment Compare ? ? ‹#› 32 Two ways of ATP formation in body 1. Oxidative phosphorylation in the presence of O2 (~ 95 % ATP) ADP + Pi + energy of H+gradient ® ATP 2. Substrate-level phosphorylation (~ 5 % ATP) ADP + macroergic phosphate-P ® ATP + second product Compare: Phosphorylation substrate-OH + ATP ® substrate-O-P + ADP (e.g. phosphorylation of glucose, proteins, etc., catalyze kinases) higher energy content than ATP ‹#› 33 Distinguish Process ATP is Oxidative phosphorylation produced Substrate-level phosphorylation produced Phosphorylation of a substrate consumed ! ‹#› 34 ATP is used for many processes • chemical work (endergonic biosyntheses) • mechanical work • electroosmotic work (ion transport across membranes) Body fluid Na+ K+ Ca2+ Cl- ECF ICF 140 10 4 160 2.5 0.0001 110 3 Average concentrations of ions (mmol/l) ‹#› 35 Less usual hydrolysis of ATP means that two ATP are consumed ATP + H2O ® AMP + PPi PPi + H2O ® 2 Pi (diphosphatase, pyrophosphatase) AMP + ATP ® 2 ADP (adenylate kinase) -------------------------------------------------- summary: 2 ATP + 2 H2O ® 2 ADP + 2 Pi ‹#› 36 High-energy phosphates and ΔG°´ of their hydrolysis (kJ/mol) Compound Phosphate derivative ΔG°´ Phosphoenolpyruvate Carbamoylphosphate 1,3-Bisphosphoglycerate Creatine phosphate ATP enolester mixed anhydride (acylphosphate) mixed anhydride (acylphosphate) phosphoamide double anhydride -62 -52 -50 -43 -31 ‹#› 37 Phosphoenolpyruvate pyruvate enolpyruvate phosphoenolpyruvate en + ol = hydroxyl attached to a double bond C=C ‹#› 38 1,3-Bisphosphoglycerate glycerol glyceric acid glycerate 1,3-bisphosphoglycerate anhydride ester ‹#› 39 Carbamoyl phosphate carbonic acid H2CO3 carbamoyl (acyl of carbamic ac.) carbamoyl phosphate carbamic acid phosphate ‹#› 40 Creatine phosphate glycine guanidinacetic ac. creatine creatine phosphate from arginine from methionine ‹#› 41 Other high energy nucleoside triphosphates have specialized functions •UTP – activation of glucose and other sugars •UTP + glucose-1-P ® UDP-glucose + PPi •(synthesis of glycogen) •GTP – gluconeogenesis, proteosynthesis •CTP – activation of choline, ethanolamine (synthesis of phospholipids) • ‹#› 42 UDP-glucose is activated glucose N-glycosidic bond O-glycosidic bond (ester type) CH2OH O uracil ribose glucose diphosphate ester anhydride ‹#› 43 Distinguish Compound Commentary Glucose metabolic fuel for most tissues, energy content = 17 kJ/g = 1700 kJ/100 g Glucose-6-P metabolically active form of glucose, enters various enzymatic reactions UDP-glucose activated form of glucose, energy-rich derivative for endergonic synthesis ‹#› 44 •Synthesis of glycogen requires energy: • 5-Glc + UDP-Glc ® 5-Glc-O-Glc + UDP • • •Cleavage of glycogen does not require energy: • (Glc)n + Pi ® (Glc)n-1 + Glc-1-P low-energy substrate (glucose unit) high-energy reagent (activated glucose) high-energy substrate (glycogen macromolecule) low-energy reagent (phosphate) Compare ‹#› 45 Consider and distinguish: low energy content = stable = non-reactive (CO2, H2O, NH3, Pi) high energy content = unstable = reactive thioesters (CoA-S-acyl), mixed anhydrides (acylphosphates), nucleoside triphosphates (ATP, GTP, UTP) etc. Ä ‹#› 46 Low-energy compounds and their activated forms Low-energy compound High-energy activated form Commentary CO2 carboxybiotin to make carboxybiotin, one ATP mol. is needed; the cofactor of carboxylation reactions, e.g.: pyruvate + carboxybiotin ® oxalocetate + biotin NH3 carbamoyl phosphate two ATP molecules are needed for NH3 activation; substrate for urea and pyrimidine synthesis Pi ATP there are two ways of ATP production in the body R-COOH R-CO-S-CoA (acyl-CoA) unusual ATP decomposition occurs in activation: R-COOH + ATP + CoA-SH ® R-CO-S-CoA + AMP + PPi ‹#› 47 Kinetics: Basic terms •reaction: S ® P (S = substrate, P = product) •Definition of reaction velocity (rate): • this definiton is for average rate, instantaneous rate: d[S]/dt ‹#› 48 Initial velocity vo •The highest value of velocity •It is not influenced by the decrease of substrate nor the reverse change of product •Determined from kinetic curves at the time t = 0 ‹#› 49 Factors influencing reaction rate •Concentration of substrate [S] •temperature •The presence of catalyst or inhibitor • •In enzyme reactions •Concentration of enzyme [E] •pH ‹#› 50 Kinetic equation for reaction S ® P v = k [S] = k [S]1 Þ reaction of 1. order k = rate constant ‹#› 51 Concentration of substrate drops down during reaction Þ kinetic curve for substrate Reaction rate is determined from kinetic curve equilibrium ‹#› 52 Reaction of 0. order •Reaction rate does not depend on the substrate concentration •v = k [S]0 = k . 1 = k = constant •At great excess of substrate S in enzyme rections only in laboratory conditions ‹#› 53 Oxidation Reduction Loss of electrons Dehydrogenation (loss of 2H)a,b Oxygenation (gain of O)d,e Gain of electrons Hydrogenation (gain of 2H)c Deoxygenation (loss of O) a H = electron + proton = e- + H+ b also desaturation c also saturation d electronegative oxygen atom shifts bond electrons to itself = relative loss of electrons e specific case is hydroxylation Redox reactions are based on electron transfer ! ‹#› 54 Cu2+ + Fe ® Cu + Fe2+ reduction of cupric ion to copper Zn + Cu2+ ® Zn2+ + Cu oxidation of zinc ribose ® deoxyribose deoxygenation C(s) + O2 ® CO2 oxidation (combustion) of carbon Different types of redox reactions – examples – Loss and gain of electrons – Oxygenation and deoxygenation – Dehydrogenation and hydrogenation CH3CH2-OH CH3CH=O dehydrogenation of ethanol to acetaldehyde – 2H + 2H CH3–C–COOH O CH3–CH–COOH OH hydrogenation (reduction) of pyruvate to lactate ‹#› 55 Oxidation Ared - n e- ® Aox Reduction Box + n e- ® Bred Ared + Box D Aox + Bred Oxidation and reduction occur always simultaneously. Aox /Ared Box/ Bred Redox pairs Redox pair = half-cell = partial reaction Redox reaction = combination of two redox pairs ‹#› 56 A redox pair are two species, which differ each from other in the oxidation number of one or more atoms of the same element (mostly also in the number of electrons). One component of a redox pair is more oxidized and can give the second one by reduction (in a "half-reaction" of a particular redox reaction). Do not confuse redox pairs and conjugate pairs A conjugate pair is a couple that consists of an acid and a base (that differ just in one proton H+ ). Conjugate pair: R-COOH / R-COO- carboxylic acid / carboxylate ion Redox pair: R-COOH / R-CH=O carboxylic acid / aldehyde ‹#› 57 Substance donates H+ = acid Substance donates H = reducing agent Substance accepts H+ = base Substance accepts H = oxidizing agent Consider the difference ! ‹#› 58 Oxidation number Is the charge an element has in a simple ion it forms or the it would hypothetically have, if the shared electron pairs in covalent bonds are assigned to the more electronegative element sharing the pair of electrons. Examples: Oxidation number of sulfur (x) in sulfuric acid: H2SO4 2 ´ (+I ) 4 ´ (–II ) 2 + (– 8) = – 6 – 6 + x = 0 + VI Different oxidation numbers of nitrogen: -IIINH3 N20 N2IO NIIO NIIIO2– NVO3– The algebraic sum of the oxidation numbers of elements in a molecular compound equals zero and, in a polyatomic ion, it must equal the charge on the ion. ‹#› 59 Rules to assign oxidation numbers – The oxidation number of any free element is zero, even when the atoms are combined with themselves (e.g. O2, P4, S8). – No regard is paid to covalent bonds between atoms of the same species. – An element may have more than one oxidation number, if it forms a variety of compounds. – The oxidation number of hydrogen in a compound or an ion is + I except in ionic hydrides (– I). – The oxidation number of oxygen in a compound or in an ion is –II except in peroxides (it takes on a – I). – Metals generally have only positive oxidation numbers in compounds. – The oxidation number of alkali metals equals always + I, of alkaline earth metals always + II. – Nonmetals have negative oxidation numbers when combined with metals, positive oxidation numbers when combined with more electronegative nonmetals. ‹#› 60 Oxidation numbers of carbon in organic compounds -IV C H 4 C H 3 C H 2 O H -III -I -III C H 3 C O H O III C H 3 C H O -III I C O O IV C H 3 C H 2 C H 3 -III -II -III ‹#› 61 The strength of oxidants and reductants (their tendency to gain or lose electrons) is expressed for particular redox pairs by standard electrode potentials E°. Standard electrode potential E° is the potential of redox pair (both oxidized and reduced form at c = 1 mol/l) established relatively to the potential of 0.00 V of the standard hydrogen electrode (H+/H2 pair under standard state conditions). ‹#› 62 Standard state of a redox pair: – both oxidized and reduced form of a redox pair at c = 1 mol/l – specified temperature, usually 25 °C – atmospheric pressure 101.3 kPa is important only when there is a gaseous component of the redox pair H2 H+ [H+] = 1 mol / l pH2 = 101.3 kPa E0(H+/H) = 0.000 V (25 °C) Standard hydrogen electrode – reference electrode electrode - inert metal Aox Ared [Aox] = [Ared] = 1 mol/l Redox pair to be measured in the standard state ‹#› 63 Half-reaction (Redox pair) E° (V) Na+ + e− " Na Zn2+ + 2 e− " Zn Fe2+ + 2 e– " Fe 2 H+ + 2 e− " H2 Cu2+ + 2 e− " Cu I2 + 2 e− " 2 I − Fe3+ + e− " Fe2+ O2 + 4 H+ + 4 e− " 2 H2O Cr2O7− + 14 H+ + 6 e− " 2 Cr3+ + 7 H2O MnO4− + 8 H+ + 5 e− " Mn2+ + 4 H2O H2O2 + 2 H+ + 2 e− " 2 H2O − 2.71 − 0.76 – 0.44 0.00 0.34 0.54 0.76 1.23 1.33 1.51 1.77 Examples of standard electrode potentials (25 °C) ‹#› 64 The guiding principle: Under standard conditions, any oxidant will react with any reductant with a lower, more negative E° Na+ / Na Zn2+ / Zn Fe2+ / Fe H+ / ½ H2 Cu2+ / Cu I2 / 2 I− Fe3+ / Fe2+ O2 / 2 H2O Cr2O72− / 2 Cr3+ MnO4− / Mn2+ H2O2 / H2O strong oxidants weak oxidants strong reductants weak reductants ‹#› 65 If the difference ΔE° between both redox pairs is greater than 0.400 V, the reaction is irreversible (i.e. proceeds to completion) even under various non-standard concentrations of the reactants. If the difference between both E° is less than 0.400 V, then the reaction will reach equilibrium, the position of which depends on the initial concentrations of components of both redox pairs; the direction of such a reaction may be reversed. Electrode potentials E under non-standard conditions for a redox pair a Aox + n e– ® b Ared Nernst equation RT nF ln E = E0 + [Aox]a [Ared]b [Aox] and [Ared ] relevant concentrations of reactants R = 8.314 l kPa K–1 mol–1 F = 96 500 C mol–1 n = number of moles of electrons transferred E, E0 el. potentials in volts 0.059 n log E = E0 + [Aox]a [Ared]b (in volts; t = 25 °C) ‹#› 66 Redox reactions in biological systems Most biological redox reactions are catalyzed by enzymes. Oxidative metabolism of nutrients rich in hydrogen releases energy required to carry out any work of the body. Some synthetic pathways, e.g. synthesis of fatty acids or cholesterol also include several redox reactions, but those are predominantly reductions (reductive syntheses). Redox reactions serve to also other specialized purposes, e.g. hydroxylations of numerous compounds foreign to the cells (xenobiotics) and dehydrogenation of alcohols. Nutrients ® CO2 + NADH, FADH2 + heat ‹#› 67 Reduced cofactors NADH+H+ or FADH2 are reoxidized in the respiratory chain within the inner mitochondrial membrane. The oxidation of nutrients proceeds through several dehydrogenation steps Dehydrogenations are catalyzed by the enzymes dehydrogenases. The two atoms of hydrogen are taken off from substrates and are accepted by the oxidized forms of cofactors NAD+ or FAD. + substrate to be oxidize oxidized cofactor + 2nd redox pair 1st redox pair dehydrogenase oxidized substrate reduced cofactor ‹#› 68 Vitamin nicotinamide (niacin) is part of NAD+ nicotinic acid (pyridine-3-carboxylic) nicotinamide ‹#› 69 NAD+ (nicotinamide adenine dinucleotide) ribose diphosphate ribose N-glycosidic linkage N-glycosidic linkage addition of hydride H- anion ‹#› 70 NAD+ is cofactor of dehydrogenases •NAD+ is oxidant – takes off 2H from substrate •one H adds as hydride ion (H-) into para-position of pyridinium cation of NAD+ •the second H is realesed as proton (H+) and binds to enzyme molecule 2 H = H– + H+ ‹#› 71 Redox pair of cofactor oxidized form NAD+ reduced form NADH aromatic ring aromaticity disturbed positive charge on nitrogen electroneutral species high-energy compound ‹#› 72 Riboflavin (vitamin B2) comprizes colourless ribitol and yellow dimethylisoalloxazine colourless ‹#› 73 FAD (flavin adenine dinucleotide) vazba 2H adenine ribosa difosfát ribitol dimethylisoalloxazin ribose diphosphate accepting two H atoms ‹#› 74 Redox pair of cofactor oxidized form FAD reduced form FADH2 aromatic system aromaticity disturbed electroneutral species electroneutral species high-energy compound ‹#› 75 Redox pairs in the respiratory chain E°´ (V) NAD+ + 2 H+ + 2 e− " NADH + H+ FAD + 2 H+ + 2 e− " FADH2 FMN + 2 H+ + 2 e− " FMNH2 2 cytochrome b (Fe3+) + 2 e− " 2 cytochrome b (Fe2+) ubiquinone + 2 H+ + 2 e− " ubiquinol 2 cytochrome c (Fe3+) + 2 e− " 2 cytochrome c (Fe2+) 2 cytochrome a3 (Fe3+) + 2 e− " 2 cytochrome a3 (Fe2+) ½ O2 + 2 H+ + 2 e− " H2O − 0.320 a) a) 0.030 0.100 0.235 0.385 0.816 a) Flavoproteins exhibit variable values of E°´ (0.003 – 0.091 V) which depend on the protein part of the enzyme. Electrode potentials in biological systems are related to pH value 7.00 and temperature 30 °C; the symbols are E°' and E' instead of E° and E, respectively