BMS100_BCH1-05_W23_Thermodynamics part 1.pdf

Definitions • System – (chemical): All reactants and products and the immediate surrounding environment § Closed system – the system exchanges energy, but not matter with the environment outside the system § Open system – the system exchanges energy and matter with the environment outside the system...

Definitions • System – (chemical): All reactants and products and the immediate surrounding environment § Closed system – the system exchanges energy, but not matter with the environment outside the system § Open system – the system exchanges energy and matter with the environment outside the system • “Universe”: the system and the environment outside the system First law of thermodynamics • In any chemical change, the total amount of energy in the universe remains constant, although the form of the energy may change § Kinetic energy • Heat, motion, pressure § Potential energy • Bonds, gradients • Biological systems depend on energy transformation Heat “Complexity” “Metabolism” Movement Gradients Internal Energy and Enthalpy • The internal energy of a system includes: § Potential energy of bonds § Kinetic energy (heat and pressure) • Enthalpy simplified - the energy content of the bonds within the system § Very difficult to measure directly… and not useful • We tend to measure the change in enthalpy that occurs when reactants form products à ΔHrxn § We can measure the change in enthalpy by measuring the change in heat within the system § ΔErxn = ΔHrxn = ∑ ΔHP - ∑ ΔHR Definitions • A reaction that releases heat is exothermic § It has a negative ΔH (-ΔH) § Tends to be spontaneous – but depends… • A reaction that absorbs heat is endothermic § Positive ΔH (+ΔH) § Often is not spontaneous – but depends… • A reaction that neither releases nor absorbs heat is isothermic § Zero ΔH Second Law of Thermodynamics • Systems move spontaneously from order to disorder – the total entropy of the universe is continually increasing • Entropy – a useful definition § How “spread out” (dispersed) energy is in a system § An increase in entropy (S) means that the energy in that system is more dispersed • Example – a hot pan cools down § The molecules in the pan that were moving rapidly represent “concentrated” energy § The pan does cool down… but its energy was dispersed widely to the atmosphere around it • The pan didn’t cool down by “concentrating” its heat elsewhere Second Law of Thermodynamics • Even more chemically-useful: § Energy disperses or spreads out unless it is hindered from doing so • Example: Why isn’t gasoline exploding all the time? Why does it need a spark? § The reaction is highly exothermic – when gasoline combusts to form water vapour and carbon dioxide, the total chemical bonds go from a higher to a lower energy state (-ΔH) § The entropy of the system increases – the gaseous carbon dioxide and water are more “random” or dispersed than the reactants (gasoline and oxygen) Using the Laws of Thermodynamics • What we really want to know: § Will a chemical reaction happen – i.e. will it be spontaneous? § Will that chemical reaction be a source of energy, or will it require energy? • Based on the laws of thermodynamics, we know that: § Reactions that release heat (-ΔH) tend to be spontaneous § Reactions that increase entropy (+S) tend to be spontaneous Gibbs Free Energy • Measures the energy difference between the products and reactants and takes into account both entropy & enthalpy Change in Entropy of the system ΔG = ΔH - TΔS Temperature (in Kelvin) Enthalpy change = heat absorbed or released • A negative ΔG = a spontaneous reaction • A positive ΔG = a nonspontaneous reaction Gibbs Free Energy • Definitions: § (-) ΔG = an exergonic reaction = spontaneous § (+) ΔG = an endergonic reaction = non-spontaneous • So, what is happening if ΔG=0? § The reaction is at equilibrium • No net transfer of heat or energy • ΔG is dependent on temperature, pH, and relative concentrations of products and reactants § Standard Gibbs free energy is calculated at 298 K (room temperature), at a pH of 7, at 1 M of each reactant and product à ΔGoʼ Exergonic? Endergonic? • So, when would the following reactions be spontaneous? 1. 2. 3. 4. A (-) ΔHrxn and a (+) ΔSsystem A (+) ΔHrxn and a (+) ΔSsystem A (-) ΔHrxn and a (-) ΔSsystem A (+) ΔHrxn and a (-) ΔSsystem • Possible answers: § Always § Sometimes § Never Standard Gibbs Free Energies Reaction Glucose + 6 O2 à 6 CO2 + 6 H2O ATP + H2O à ADP + Pi ATP + H2O à AMP + PPi Glucose 6-phosphate à fructose 6-phosphate Fructose 1,6 bisphosphate à DHAP and G3P Glucose + Pi à glucose 6-phosphate ΔGoʼ -2840 kJ/mol -30.5 kJ/mol -45.6 kJ/mol +1.67 kJ/mol +23.9 kJ/mol +13.8 kJ/mol Gibbs Free Energy – Non-Standard Conditions as a Biologic Strategy • When products and reactants are not both at 1 M, then the reaction is said to be at non-standard conditions • Impacts the Gibbs free energy: ΔG = Non-standard ΔG ΔGoʼ + RT ln 8.314 J/mol K ["#$%&'()] [#+,'(,-()] Natural log Gibbs Free Energy – Non-Standard Conditions as a Biologic Strategy • Hypothetical reaction with a ΔGoʼ of 5.0 kJ/mol: A+BàC+D • Plugging into the values into from the table into the equation: ΔG = ΔGoʼ + RT ln ! [#] % [&] [A] and [B] [C] and [D] 1 10 2.3 10.6 kJ/mol 10 10 0 5.0 kJ/mol 30 10 -1.1 2.3 kJ/mol 100 10 -2.3 -0.57 kJ/mol ln ! [#] % [&] ΔG Gibbs Free Energy – Non-Standard Conditions as a Biologic Strategy • How does the Gibbs free energy change as: § The reactants are increased compared to the products? § The reactants are reduced compared to the products? [A] and [B] [C] and [D] 1 10 2.3 10.6 kJ/mol 10 10 0 5.0 kJ/mol 30 10 -1.1 2.3 kJ/mol 100 10 -2.3 -0.57 kJ/mol ln ! [#] % [&] ΔG Gibbs Free Energy – Non-Standard Conditions as a Biologic Strategy • Even reactions with a fairly large +ΔG can develop into exergonic reactions if the [reactants] is substantially greater than the [products] • You will see this is a common strategy used to drive some of the endergonic reactions of glycolysis forward § It’s actually a common strategy used in many biochemical and cellular physiological situations Gibbs Free Energy – Coupled Reactions as a Biologic Strategy • The following is the first reaction of glycolysis: § Glucose + Pi Glucose-6-phosphate • ΔGoʼ = 12 kJ/mol § Note the positive ΔGoʼ § How is it possible that a reaction with a (+) ΔGoʼ still occurs in our bodies – all the time? Gibbs Free Energy – Coupled Reactions as a Biologic Strategy Gibbs Free Energy – Coupled Reactions as a Biologic Strategy • This is the whole “rationale” for why the body phosphorylates ATP (and dephosphorylates it) as a source of energy currency § The high-energy phosphodiester bond, when broken, has a negative ΔG § That energy can be coupled to another reaction with a positive ΔG § The net reaction is exergonic… • if the ΔG of ATP is “negative enough” to counteract the endergonic ΔG of the coupled reaction How do Enzymes help us out? • Most coupled reactions involve enzymes as catalysts • Do enzymes themselves drop the ΔG? § No – enzymes only drop the activation energy (kind of like the spark to gasoline) § Activation energy = the energy required to break chemical bonds à leads to the formation of new ones… • According to the ΔG of that reaction