Reaction kinetics

Thermodynamics says which direction a reaction would go; kinetics says whether it will get there this century. The two are independent, and the gap between them is why diamond does not turn into graphite while you watch even though doing so would release energy. The conversion is favourable and effectively infinitely slow, because rearranging the carbon lattice demands an enormous activation energy that room temperature never supplies.

Rate comes from collisions. Reactant particles must meet, hit hard enough to clear the activation barrier, and be oriented correctly when they do. Every practical way of speeding a reaction acts on one of those three. Higher concentration or pressure means more frequent collisions. Larger surface area — powder instead of a lump — exposes more particles to collide. Higher temperature does the most work, not by making collisions slightly more frequent but by pushing far more molecules into the high-energy tail of the distribution that can actually clear the barrier.

Catalysts work differently again: they open a new pathway with a lower activation barrier, so a much larger fraction of collisions succeeds. A catalyst takes part in the mechanism — it is not a bystander — but is regenerated by the end, so it is not consumed and a tiny amount can process an unlimited quantity. And note what it cannot do: lowering the barrier speeds both directions equally, so a catalyst never changes where equilibrium lies or makes an unfavourable reaction favourable. Finally, rate laws are found by experiment, not read off the balanced equation — the coefficients describe the overall stoichiometry, not the slow step that actually controls the rate.