Energy and Redox Reactions BIOL 160 Sphack Lecture 8
Energy and Redox Reactions BIOL 160 Sphack Lecture 8 Biol 160
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This 4 page Class Notes was uploaded by Anushree Kumar on Monday February 29, 2016. The Class Notes belongs to Biol 160 at University of Tennessee - Knoxville taught by Dr. Shpak in Spring 2016. Since its upload, it has received 27 views. For similar materials see Cellular and Molecular Biology in Biology at University of Tennessee - Knoxville.
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If Anushree isn't already a tutor, they should be. Haven't had any of this stuff explained to me as clearly as this was. I appreciate the help!
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Date Created: 02/29/16
BIOL 160 Sphack Energy and Redox Reactions Important point key term learning objective Laws of Thermodynamics - 1 law – energy is conserved. o Can’t be created/destroyed. o But it CAN be transformed and transferred nd - 2 law – total entropy always increases in an isolated system Entropy (S) – amount of disorder in a group of molecules Energy – the capacity to do work or supply heat - Energy exists in 2 ways o Potential energy (aka chemical energy) – hasn’t done anything yet, but has the ability to do. Example – water particles at the top of waterfalls have a lot of potential energy. They haven’t fallen yet, but they have the ability to do so Potential energy can be expressed as height o Kinetic energy – active motion. Example – water that’s moving down a waterfall has kinetic energy. They’re actively moving Potential Energy - also called chemical energy - Strong bonds have low chemical/potential energy o Strong bonds aren’t as easy to break for their energy to be used - Weak bonds have high chemical/potential energy o These bonds can break easily for energy to be used - It takes more to break a strong bond than a weak bond - For atoms, electrons further away have a higher potential energy o Electrons closer to nucleus have a stronger affinity to stay there o Electrons farther away can be easily “persuaded” to leave the atom Kinetic energy – energy of molecular movement - Measured in temperature o Low temperature – molecules move slower. “cold” o High temperature – molecules move faster “hot” Enthalpy (H) – total energy of the molecule - Potential energy + kinetic energy = total energy (enthalpy) - Symbol is “H” Entropy (S) – the amount of disorder in a group of molecules - 2ndlaw of thermodynamics - Example: diffusion o As the particles spread out, they become less organized Chemical Reactions – have reactants and products - Chemical equilibrium o Quantities/amount of reactants is same as products o The reaction can happen at the same rate if it happens forwards or backwards o Amount of molecules in the system stays the same - Spontaneous o They can happen on their own Don’t need to be facilitated with enzyme Don’t need to put energy into the reaction to make it happen o Depends on 2 things Change in potential energy Products usually have less potential energy than reactants Degree of order/entropy Products have more entropy than reactants/ o Another way to say it: Products are less ordered than reactants o Reactants High potential energy More order – lower entropy rate o Products Low potential energy Less order – low entropy rate, but reactants have a higher entropy than products o Enthalpy stays the same, because it’s the total of potential and kinetic energy - Nonspontaneous reactions o Would need to have energy put into them Gibbs Free Energy : ΔG = ΔH – TΔS - The change of Gibbs Free Energy = change in enthalpy – ((temperature) (change in entropy)) o ΔH : difference of enthalpy o ΔS : difference of entropy (disorder) o T : temperature - Exergonic (aka spontaneous) : ΔG<0 - Endergonic (aka nonspontaneous): ΔG>0 Exergonic vs. Endergonic - Exergonic o If ΔG<0, then it’s spontaneous o Happens naturally, and reaction does not need energy supplied to it Example: think of a boulder on the top of the hill. If it starts rolling down the hill, you don’t need to keep pushing it in order for it to keep rolling down the hill - Endergonic o If ΔG>0, it’s nonspontaneous o Can only happen with another reaction ATP in the system Exothermic vs. Endothermic - Exothermic – reaction that releases heat energy o ΔH<0 but entropy decreases o Temperature increases o Example: log burning o High energy reactants -> low energy products - Endothermic – reaction that absorbs heat energy o ΔH>0 and entropy increases o Temperature decreases o Example: ice melting o Low energy reactants -> high energy products - Combinations of ΔH and ΔS o Negative ΔH and positive ΔS = reaction is spontaneous at energy temperature o Positive ΔH and positive ΔS = not spontaneous at low temps, spontaneous at high temps o Negative ΔH and negative ΔS = spontaneous at low temperatures o Positive ΔH and negative ΔS = not spontaneous at all Remember that ΔG will tell you if it’s spontaneous or not. It’s all about mathematical relationships in the equation ΔG = ΔH – TΔS Redox Reactions – reduction-oxidation reactions that involve electron transfer - Reduction – gain of one or more electrons and a hydrogen ion (H+) o Atom is reduced when it gains an electron The charge of the atom is “reduced” (made negative) - Oxidation – when an atom/molecule loses an electron o Loss of one or more electrons - Oxidation and reduction reactions are coupled o If an atom loses/gain an electron, a different atom gains/loses it. Electron donors are paired with electron acceptors - Electrons can o Be transferred completely to another atom o Shift their position in covalent bonds
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