Red (A) and blue (B) reactant molecules bounce around a sealed chamber at a speed set by the temperature. Every 0.1 s "chemistry tick", any A and B pair close enough to collide gets a chance to react — computed straight from the Arrhenius equation — and turns into green product C. C can also thermally decompose back into A + B, so the mixture settles toward a dynamic equilibrium rather than fully converting. Gold nanoparticle clusters mark catalyst sites: any collision happening in their vicinity reacts against a lower activation energy, in both the forward and reverse direction — which is why a catalyst speeds up how fast equilibrium is reached without changing where that equilibrium sits.
k(T) = A·e^(−Eₐ/RT)
near a catalyst site: Eₐ,eff = Eₐ − ΔEₐ_cat (ΔEₐ_cat = 18 kJ/mol)
reverse barrier: Eₐ,rev = Eₐ + |ΔH| (ΔH = −25 kJ/mol, exothermic)
per-tick reaction chance: p = 1 − e^(−k·Δt)
- Temperature — sets molecular speed (and collision frequency) and appears directly in the Boltzmann factor e^(−Eₐ/RT); higher T raises the rate constant exponentially, not linearly.
- Activation energy Eₐ — the energy barrier a colliding pair must clear to react. Small changes here swing the rate constant by orders of magnitude because it sits in the exponent.
- Catalyst sites — how many active-site clusters are seeded in the chamber; more sites means a larger fraction of collisions happen inside the lowered-barrier zone, raising the overall conversion rate even though each individual site lowers Eₐ by the same fixed amount.
- Catalyst toggle — switches the lowered-barrier effect on or off so you can directly compare catalyzed vs. uncatalyzed kinetics at the same temperature.
Real-world relevance: this is exactly how catalytic converters, industrial ammonia synthesis (Haber-Bosch) and enzymes work — none of them change the thermodynamics of a reaction, they all open a lower-energy pathway so the same equilibrium is reached far faster.