kinetics, not thermodynamics

Chemistry draws a hard line between whether a reaction can happen and how fast it will, and a lot of confused arguments are really a collision between those two questions. Something can be entirely favorable and still almost never occur, because whether a thing happens and how fast it happens run on different machinery.

\[ \begin{aligned} k &\propto e^{-E_a / RT} \\ RT &\approx 0.6 \ \tfrac{\text{kcal}}{\text{mol}} \quad (\text{near room temperature}) \\ \text{lower } E_a \text{ by } 6 \ \tfrac{\text{kcal}}{\text{mol}} = 10\,RT \;&\Rightarrow\; \frac{k_{\text{new}}}{k} = e^{10} \approx 22{,}000 \end{aligned} \]

k is the reaction rate, E-a the activation energy, meaning the height of the hill between start and finish. R is the gas constant and T the absolute temperature, so RT is the thermal energy available to climb that hill, about 0.6 kcal/mol at room temperature. lowering the barrier by 6 kcal/mol, which is ten times RT, makes the catalyzed rate k-new about e to the tenth, roughly 22,000 times the original rate k, while the start and end points don't move at all.

This is what a catalyst does. It doesn't change where the reaction ends up, only the size of the hill in the middle -- and the rate shoots up. A good mentor or a good tool works the same way. It doesn't make something happen that was never going to, it just lowers the cost of the thing that was already worth doing. And being right isn't worth much if you're too slow, because by the time you get there it can look the same as having been wrong.

same start, same product. the dashed hill is the uncatalyzed reaction, the solid lower hill is with a catalyst. dropping the barrier about 6 kcal/mol speeds the reaction roughly 22,000-fold, and the two endpoints never move.