I was also taught in biology to think of oxidation in metabolism as a release of potential energy held by C-C or C-H bonds (due to equal sharing of electrons turning into a polarized sharing of electrons, releasing chemical potential energy). However oxidation does not always result in this change, sometimes it occurs by turning a polarized bond into a nonpolar bond. Can someone explain to me exactly what oxidation is in terms of potential energy and electron localization?
Moreover, does oxidation always produce energy?
The situation is a bit more complicated than biologists would like you to believe. Recall that oxidation reactions are always coupled to reduction reactions. An oxidation reaction occurs when a compound loses electrons. Because the electrons lost by the compound can't just float away, an oxidation reaction must be coupled to the gain of electrons by a second compound. This second compound therefore undergoes a reduction reaction. Thus, one can think of redox reactions as the transfer of electrons between compounds. The compound that loses the electrons is the compound undergoing the oxidation and the compound gaining electrons is the compound undergoing reduction.
I think of a redox reaction as the movement of electrons from a compound of high potential energy to a compound of low potential energy. We quantify the potential energy of these compounds by their reduction potentials (E
o). Compounds with higher reduction potentials have a higher affinity for electrons than compounds with lower reduction potentials. Therefore, redox reactions where electrons are transferred from compounds with lower reduction potentials to compounds with higher reduction potentials will be favorable (i.e. they will occur spontaneously and release free energy). To put it mathematically, if ΔE > 0 then ΔG < 0.
Consider for example the reaction occurring at complex I of the electron transport chain. Here the oxidation of NADH to NAD
+ is coupled to the reduction of coenzyme Q. Because NAD
+ has a reduction potential of -0.320V and coenzyme Q has a reduction potential of +0.060V, the reaction occurs spontaneously. Thus, the free energy of the reaction can be used to perform work (here, pumping protons against their concentration gradient).
Even though this reaction oxidizes NADH and releases free energy, it's incorrect to say that the oxidation of NADH produces energy. It's the specific combination of the oxidation of NADH and reduction of coenzyme Q that causes energy to be released. If one pairs the oxidation of NADH with a reduction reaction whose reduction potential is less than -0.320V, the redox reaction would require an input of energy to proceed.
How can an exothermic reaction be endergonic? Conversely, how can an endothermic reaction be exergonic
Simply put there are two driving forces in chemistry. The first is the desire to reduce the potential energy of the system. The second is the desire to increase the entropy of the system. Often, the balance between these two, sometimes opposing, goals will define whether a reaction is favorable or not. An exothermic reaction can be energonic if the reduction of entropy of the system is enough to outweigh the loss of potential energy of the system.
For example, consider the formation of ice from liquid water. The formation of ice is exothermic: when water freezes, the intramolecular bonds between water molecules becomes stronger, thus lowering the potential energy of the system. However, ice is much more ordered than liquid water. Thus, at room temperature, the formation of ice is an endergonic process. The energy released by the formation of ice is not enough to counter the loss of entropy.
It turns out that temperature is the factor that controls the balance between the desire to lower potential energy and the desire to increase entropy. At higher temperatures, entropy plays a bigger role and at lower temperatures potential energy plays a bigger role. Thus, below a certain temperature (the freezing point), the formation of ice becomes exergonic and occurs spontaneously. Indeed, below 0
oC, the loss of potential energy due to freezing outweighs the loss of entropy.
I'm not understanding here is the difference between delta-G and delta-H and delta-E for a reaction. Please explain, if you can, in terms of a reaction diagram (the one that shows energy state of reactants and products, with the energy hill between them which is the activation energy). I am a very visual learner.
A really good explanation of free energy can be found at the following blog: http://gravityandlevity.wordpress.com/2009/05/02/when-nature-plays-skee-ball-the-meaning-of-free-energy/
On that note, delta-E in my book = sum of bond energy of bonds broken - sum of bond energy of bonds formed. But then it also says that delta-E = sum of bond energies of reactants - sum of bond energies of products. First of all, aren't these two different values? Secondly, isn't the second difference the same thing as delta-G?
The two definitions of ΔE are indeed equivalent. Try comparing them with a simple reaction H
2 + C
2H
4 --> C
2H
6. The second definition is not the same thing as ΔG because ΔG takes the change in entropy of the system into account.
I'll leave the other questions until later/for someone else as these answers should give you something to think about.
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Topic: Bond energy, redox reactions, and thermodynamics (Read 7603 times)