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The product distribution of any chemical reaction at any instant in time is governed by two factors: thermodynamics and kinetics. The "thermodynamic" product distribution of a reaction is derived from the free energy difference between the products themselves. The "kinetic" product distribution, on the other hand, is derived from the energy difference between the transition states leading to the various products. It is a general principle of chemistry that lower-energy species will predominate over higher-energy species that come from the same starting material. As a result, the "kinetic product" of a reaction is the one with the lower energy transition state, while the "thermodynamic product" is the one with the lower energy ground state (the energy of the product itself). These two are not necessarily the same.

The question is, how do you know when your real-life reaction on the benchtop is under thermodynamic or kinetic control? A good starting point is estimation of the so-called "delta delta G" values (two deltas--not a typo), which represent the difference in free energy change between two competing reaction pathways. The difference in overall energy change we might call the "thermodynamic delta delta G," while the difference in activation energy we might call the "kinetic delta delta G." To a first approximation, the larger of the two delta delta G's will determine how your reaction is controlled. So for example, if the difference in ground state energy of two possible products (thermodynamic DDG) is large compared to the difference in the energies of activation required to reach these products (kinetic DDG), then the rates of both reaction pathways will be approximately the same, the products will form according to their ground-state energies, and the reaction will be under thermodynamic control. If, on the other hand, the kinetic DDG is large compared to the thermodynamic DDG, then one product will form much more quickly than the other, and for short time scales the reaction will be under kinetic control.

Of course if you let your reaction run too long, thermodynamic control will eventually take over. One of the most common ways to effect thermodynamic control, in fact, is to just let your reaction run for a long enough time for the rate difference between two pathways to become negligibly small with respect to the time you let the reaction run. Eventually the products will be forced to equilibrate thermodynamically! How long this takes depends on the magnitude of the kinetic DDG...the smaller it is, the longer it takes. Thermodynamic control is encouraged by doing anything that lowers kinetic DDG or increases thermodynamic DDG. To lower kinetic DDG, you can add a catalyst to bring down the higher-energy transition state closer to the lower-energy one, stabilize the higher-energy transition state through solvent effects, or use a reagent for which the reaction pathway is highly reversible (this is the same thing as saying that, for that reagent, kinetic DDG is relatively small). Increasing the temperature also effectively lowers kinetic DDG, because reverse reaction rates start becoming non-negligible as the temperature goes up.

Kinetic control is usually effected simply by lowering temperature and/or employing reagents that irreversibly perform their task (i.e., kinetic DDG is extremely large, such that the rate of the reverse reaction for the more kinetically favorable pathway is nearly zero).

As an example, consider the deprotonation of an asymmetric ketone, say (isopropyl)CO(methyl). Thermodynamically, deprotonation on the isopropyl side is favored, because this leads to a more substituted double bond. But kinetically, deprotonation on the methyl side is favored because the three hydrogens on that side are much more accessible than the one on the iPr side.

So how do we effect kinetic control? in theory we could use almost any base at super-low temperatures and get mostly the kinetic product out. The problem with this is that lowering the temperature of the reaction lowers the rates of both possible deprotonations, so it would take forever to get anything substantial to happen. We would end up isolating very little product. What we need is a reagent that will grab a proton quickly from the less crowded side and hold onto it, so that the reverse process of protonation on this side has a rate of essentially zero. The solution is a sterically hindered base like LDA, lithium diisopropylamide, at moderately low temperatures...say -78 C. It deprotonates more quickly at the methyl group due to steric effects, then holds onto its proton for dear life.

How then could we effect thermodynamic control? We need a base that can reversibly take off and put on a proton, so that the forward and reverse rates of reaction are comparable, rate differences for the two kinds of protonation are small, and thermodynamic equilibrium is quickly reached. This can be found in the alkoxide bases: sodium tert-butoxide, sodium ethoxide, etc. Kinetic DDG for these reagents is much lower than for LDA (particularly in the case of the unhindered ethoxide and methoxide), and the reaction can be run at room temperature to exaggerate thermodynamic DDG. This leads to predominant deprotonation of the more sterically hindered isopropyl side of the ketone, leading to the more thermodynamically stable enolate.

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