Bioenergetics 2_2024.pdf

BCCB2000 Foundations of Biochemistry Bioenergetics 2 Ricardo L. Mancera WARNING This material has been reproduced and communicated to you by or on behalf of Curtin University in accordance with section 113P of the Copyright Act 1968 (the Act) The material in this communication may be subject to copy...

BCCB2000 Foundations of Biochemistry Bioenergetics 2 Ricardo L. Mancera WARNING This material has been reproduced and communicated to you by or on behalf of Curtin University in accordance with section 113P of the Copyright Act 1968 (the Act) The material in this communication may be subject to copyright under the Act. Any further reproduction or communication of this material by you may be the subject of copyright protection under the Act. Do not remove this notice. 1 Learning Objectives  When you complete this lecture you will be able to  Describe coupled chemical reactions  Describe activation energy  Describe how ATP is the main biological ‘carrier’ of energy  Understand and describe the hydrolysis of ATP  Describe how electrons may transfer energy  Describe the main electron carriers  Calculate the free energy from reduction potentials  Describe the concept of ‘reduced bonds’  Calculate the biologically available Gibbs free energy in common small biomolecules  Use the biochemical knowledge learned to answer questions and solve problems Gibbs free energy (ΔG): recap  Negative free energy change (–ΔG) means the reaction is spontaneous   Positive free energy change (+ΔG) means the reaction is not spontaneous   Exergonic Endergonic Exergonic reactions can be used to ‘drive’ endergonic reactions if coupled 2 Coupled chemical reactions Couple the two reactions Exergonic reaction Endergonic reaction Exergonic reaction ‘driving’ the endergonic reaction Ferrier2014 Coupling exergonic and endergonic reactions  A coupled exergonic and endergonic reaction needs:  A common intermediate shared by the reactions that are coupled  A mechanism for energy transfer to occur Wigglesworth1997 3 Coupled chemical reactions  In a coupled reaction we can add the ΔG’º values for each reaction to get the ‘overall’ ΔG’º values for the combined reaction  For example, the enzymatic phosphorylation of glucose: (1) Glucose + Pi ΔG’º = +13.8 kJ / mole (2) ATP + H2O Glucose-6-phosphate + H2O ADP + Pi ΔG’º = ‒30.5 kJ / mole Sum of reactions Glucose + ATP  Glucose-6-phosphate + ADP Each reaction occurs as written from left to right so the sign of ΔG’º does not change:  ΔG’º = 13.8 + (‒30.5) = ‒16.7 kJ / mole  These ‘coupled reactions’ take place within the active site of hexokinase or glucokinase ATP: the universal energy carrier Nelson & Cox, 2000 4 ATP: the universal energy carrier ATP: the universal energy carrier  Free energy change for ATP hydrolysis is large and negative (ΔG’° = ‒30.5 kJ/mole), in part because:  The terminal phospho-anhydride bonds are ‘weaker’ and contribute to the instability of ATP compared with the products   Electrostatic repulsion among four negative charges is relieved by charge separation after hydrolysis The products are relatively more stable  Phosphate is resonance stabilised  ADP2- can ionize to ADP3-  The products (phosphate and ADP) are more soluble in water than ATP 5 ATP: the universal energy carrier Nelson & Cox, 2000 ATP: the universal energy carrier ΔG ΔG 6 ATP: the universal energy carrier  Standard free energy change for ATP hydrolysis is large and negative   The actual free energy of hydrolysis (ΔG) of ATP in cells is different from the standard free energy (ΔG’°) and depends upon the cellular concentration of ATP, ADP and Pi, and is about ‒50 to ‒65 kJ/mol   ΔG’º = ‒30.5 kJ/mol This reflects a large ratio of the cellular concentration of ATP with respect to that of ADP However, the hydrolysis of ATP has a large activation energy  About 200 to 400 kJ/mole Activation energy of ATP  Activation energies determine the kinetics of a reaction  Reaction kinetics determines how fast or how slow a reaction progresses  High activation energy gives a slow reaction and low activation energy gives a fast reaction  A favourable reaction (i.e. with ‒ΔG) can still be slow unless catalysed! 7 ATP: the universal energy carrier  ATP is the ‘currency’ of energy in the cell  Continuously formed and consumed  Resting human consumes 40 kg/day  Strenuous exercise in human uses 0.5 kg/min  Different concentrations in different cells:  rat hepatocyte = 3.38 mM  human erythrocyte = 2.25 mM  E. coli cell = 7.90 mM Transfer of energy 8 Oxidation and reduction reactions  Commonly called redox reactions  Consist of a molecule, atom, or ion that is reduced and a molecule, atom or ion that is oxidised  Redox reactions always occur together  The reduced molecule (or atom or ion) gains electrons  The oxidised molecule (or atom or ion) loses electrons Transfer of energy by transfer of electrons  Electrons are transferred from one or more redox reactions in the cell to specific electron carriers  The electron carriers then transfer the electrons to other redox reactions and electron carriers in the inner mitochondrial membrane  The flow of electrons through these other carriers and redox reactions in the inner mitochondrial membrane eventually contribute to the production of ATP  ATP is then used as the energy carrier in the cell 9 Transfer of energy by transfer of electrons  Electrons are transferred in two related ways: 1) 2) Through a series of redox reactions in the cell  Some reactions are sequential  Some reactions occur in different locations of the cell By specific carriers  These ‘carry’ electrons from a redox reaction in one location to a redox reaction in another location  NAD+ / NADP+  FAD  Coenzyme A  Coenzyme Q  Others Oxidation and reduction reactions   Oxidising ‘agents’  Molecules (or atoms or ions) in a redox reaction that have the propensity to accept electrons  Gain electrons in a redox reaction and so are reduced Reducing ‘agents’  Molecules (or atoms or ions) in a redox reaction that have the propensity to donate electrons  Lose electrons in a redox reaction and so are oxidised 10 Oxidation and reduction reactions  In addition to a role in generating energy, redox reactions in the cell also have other roles  Provide electrons for the synthesis of molecules  Provide a mechanism of reaction for some enzymes Important oxidation and reduction reactions Note that all of these redox reactions are catalysed by dehydrogenases Dehydrogenases transfer hydrides (H+ atoms with two electrons) 11 Electron carriers  NAD+ and NADP+    Water soluble electron carriers: transfer two electrons as a hydride ion FAD (flavin adenine dinucleotide) and FMN (flavin mononucleotide)  Strongly (sometimes covalently) bound to proteins  Can transfer one or two electrons Other electron carriers  Ubiquinone (coenzyme Q): hydrophobic benzoquinone  Cytochrome proteins containing iron bound in a haem prosthetic group  Iron-sulphur proteins: iron not in haem but as a complex of sulphur and iron atoms joined to the protein through cysteine residues Electron carrier: NAD+/NADH Reduced substrate + NAD+ 2H+ + 2e‒ Oxidised substrate + NADH + H+ Nicotinamide functional group 2H+ +2e‒ Note that what is actually transfered to NADH is a hydride ion H‒: H‒ = H+ + 2e‒ The other H+ is in solution Position of phosphate in NADP 12 Electron carrier: FAD/FADH2 Nelson & Cox, 2008 Electron carrier: ubiquinone (coenzyme Q) Nelson & Cox, 2008 13 Electron carrier: cytochromes Fe3+ + e‒ = Fe2+ Nelson & Cox, 2008 Electron carrier: iron-sulphur proteins Nelson & Cox, 2008 14 Reduction potentials  Reduction potentials are a way of tracking the number of electrons stored or transferred  Reduction potentials determine the free energy of a redox reaction  Reduction potentials (Eº) are measured in volts (V)  Potential is measured with reference to a ‘standard’ reduction:  H+ + e‒ = ½ H2 Eº = 0 V  Standard = 1 M concentrations of oxidant and reductant at pH = 0 (or pH = 7 for biological reactions (E’º)) Reduction potentials   Reduction potentials are calculated by the Nernst Equation:  E = E’º ‒ RT/nF·ln[electron acceptor]/[electron donor]  E = E’º ‒ RT/nF·lnQ where  E = reduction potential (V)  E’º = standard reduction potential at pH 7 (V)  R = gas constant (8.315 J K‒1 mol‒1)  T = temperature (K)  n = number of electrons exchanged in reaction  F = Faraday constant (9.6485 x 104 C mol‒1)  ln = natural logarithm (base e)  Q = reaction quotient = [oxidised product of A]c [reduced product of B]d [reduced reactant A]a [oxidised reactant B]b 15 Reduction potential values Larger values indicate tendency to accept electrons and be a better oxidising agent Electrons are transferred from lower E’º to higher E’º Lower values indicate tendency to donate electrons and thus be better reducing agent. Nelson & Cox, 2008 Reduction potentials and free energy   Reduction potentials can be used to calculate the free energy of a redox reaction:  ΔG’º = ‒nFΔE’º  ΔG = ‒nFΔE where  ΔG’º = change in standard free energy (usually expressed as kJ mol‒1)  ΔE’º = change in standard reduction potential at pH 7 (V)  ΔG = change in free energy (kJ mol‒1) at non-standard conditions  ΔE = change in reduction potential at non-standard conditions (V)  n = number of electrons exchanged in reaction  F = Faraday constant (96,485 C mol‒1 or 96,485 J V‒1 mol‒1) 16 Reduction potentials and free energy Calculation of the standard free energy of the pyruvate to lactate reaction The overall reaction lactate + NAD+ Pyruvate + NADH + H+ consists of two half-reactions Pyruvate + 2H+ + 2eNAD+ + H+ + 2e- lactate E'º = ‒0.19V NADH E'º = ‒0.32V such that to get the overall equation as written above then NADH NAD+ + H+ + 2e- E'º = +0.32V ΔE'º = 0.32 ‒ 0.19 = 0.13V ΔG'º = ‒nFΔE'º ΔG'º = ‒2 x 96,485 J V-1 mol-1 x 0.13 V ΔG'º = ‒25 kJ mol-1 at standard conditions Estimating the energy content of molecules  The energy content of a molecule may be estimated by converting the number of ‘reduced bonds’ in a molecule to an energy value  The concept of a ‘reduced bond’ does not mean a different type of bond but actually refers to the oxidation state of carbon or nitrogen 17 Oxidation and reduction of carbon   A carbon is reduced if  The number of hydrogen atoms bonded to a carbon increases  The number of bonds to more electronegative atoms (e.g. O, N, F, Cl, I or S) decreases A carbon is oxidised if  The number of hydrogen atoms bonded to a carbon decreases  The number of bonds to more electronegative atoms increases Oxidation states of carbon and ‘reduced bonds’ In cells, carbon occurs in five different oxidation states 18 Estimating the energy content of molecules  A ‘reduced bond’ is considered to be any C‒C, C‒H, C‒N or N‒H bond in a molecule  When these bonds are broken, electrons are released and the carbon is oxidised  The electrons from oxidation are transferred to electron carriers, which then transfer them to mitochondria, where they can eventually be used to produce ATP  ‘Rule of thumb’ estimate of the energy (enthalpy) from each C‒C or C‒H ‘reduced bond’ is ‒220 kJ/mole  ‘Rule of thumb’ estimate for the energy (enthalpy) from each N‒C or N‒H ‘reduced bond’ is ‒105 kJ/mole  Note these are not bond energies or energies of formation, but rather simply estimates using the concept of a ‘reduced bond’ Energy content of carbohydrates: glucose Glucose Number of reduced bonds: C-C bonds = 5 C-H bonds = 7 Total = 12 Hence the estimated change in enthalpy (ΔH): 12 x -220 kJ/mole = ‒2640 kJ/mole or ‒15 kJ/g (MW = 180 g/mole) 6% difference Observed enthalpy (ΔH) of oxidation (combustion) = ‒2808 kJ/mole Calculated free energy (ΔG) of oxidation (combustion) = ‒2864 kJ/mole 19 Oxidation of glucose C6H12O6 + 6O2 6CO2 + 6H2O ΔH for glucose oxidation = -2808 kJ mole-1 ΔS for glucose oxidation = +182.4 J mole-1 At 37°C = 310 K and with ΔG = ΔH – TΔS ΔG = -2,808,000 J mole-1 – (310 K x (+182.4 J K-1 mole-1)) ΔG = -2,808,000 J mole-1 – 56,544 J mole-1 ΔG = -2,864,544 J mole-1 ΔG = -2,864.5 kJ mole-1 Energy content of fat (lipids): palmitate Palmitate H H C H H H C H C H H H C H C H H H C H C H H H C H C H H H H C H C H H C H C H H H C H O C H C O H Number of reduced bonds: C-C bonds = 15 C-H bonds = 31 Total = 46 Hence the estimated change in enthalpy (ΔH): 46 x -220 kJ/mole = -10,120 kJ/mole or -39 kJ/g (MW = 256 g/mole) Observed enthalpy (ΔH) of oxidation = -10,040 kJ/mole 20 Energy content of amino acids: alanine Number of reduced bonds: C‒C bonds = 2 C‒H bonds = 4 Total = 6 Hence estimated energy content: 6 x -220 = -1320 kJ/mole Number of reduced bonds: N‒C bonds = 1 N‒H bonds = 2 Total = 3 Hence estimated energy content: 3 x -105 = -315 kJ/mole Total estimated enthalpy (ΔH) -315 + -1320 = -1635 kJ/mole or -18 kJ/g (MW = 89 g/mole) Observed enthalpy (ΔH) of oxidation = -1623 kJ/mole Oxidation of glucose and production of ATP C6H12O6 + 6O2 6CO2 + 6H2O 12 reduced bonds x -220 kJ/mole/reduced bond = -2640 kJ/mole To convert 1 ADP + Pi 1 ATP need +31 kJ/mol Therefore, glucose should be capable of producing 2640 kJ/mol / 31 kJ/mol/ATP molecule ≈ 85 ATP molecules But only about 36-38 ATP molecules are produced…why? Because oxidation occurs under biochemical conditions 21 Questions for learning  Explain how the concept of ‘reduced bonds’ can be used to estimate the energy content of molecules  Estimate the energy content of glycine using the concept of ‘reduced bonds’  Where does most of the free energy in ATP come from?  The hydrolysis of ATP is an exothermic reaction with a large negative free energy. What prevents solid (crystal) ATP in an open bottle from reacting with water in the air and forming products? Further reading  Gary, R.K. (2004) The Concentration Dependence of the ΔS Term in the Gibbs Free Energy Function: Application to Reversible Reactions in Biochemistry. Journal of Chemical Education, 81:1599-1604. https://pubs.acs.org/doi/abs/10.1021/ed081p1599  Feinman, R.D. (2006). Oxidation-reduction calculations in the biochemistry course. Biochemistry and Molecular Biology Education, 32: 161-166. https://iubmb.onlinelibrary.wiley.com/doi/full/10.1002/bmb.2004.494032030347  Stokes, G.B. (1988). Estimating the energy content of nutrients. Trends in Biochemical Sciences, 13: 422-424. https://www.sciencedirect.com/science/article/abs/pii/0968000488902101 22