Thermodynamics in Human Biology - Part 1 PDF

Thermodynamics in Human Biology - Part 1 Pre-learning Dr. Andrew Vargo BMS 100 Week 2 Video Links Laws of thermodynamics https://ccnm.ca.panopto.com/Panopto/Pages/Viewer.a spx?id=e8f4a0c4-f749-4474-8602-af81002db639 Spontaneity and Gibb’s Free Energy https://ccnm.ca.panopto.com/Panopto/Pages/Vi...

Thermodynamics in Human Biology - Part 1 Pre-learning Dr. Andrew Vargo BMS 100 Week 2 Video Links Laws of thermodynamics https://ccnm.ca.panopto.com/Panopto/Pages/Viewer.a spx?id=e8f4a0c4-f749-4474-8602-af81002db639 Spontaneity and Gibb’s Free Energy https://ccnm.ca.panopto.com/Panopto/Pages/Viewer.a spx?id=4db98dda-e6c7-4670-97f7-af81002dddb9 Gibb’s Free Energy and biology https://ccnm.ca.panopto.com/Panopto/Pages/Viewer.a spx?id=1f8c0d5d-d014-44d3-99d2-af81002dddd1 Overview Laws of thermodynamics First and second laws Enthalpy and entropy Spontaneity and Gibb’s Free Energy Exothermic and endothermic reactions Exergonic and endergonic reactions Catalysis and activation energy Gibb’s Free Energy and biology Coupled reactions Non-standard conditions 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