Activation Energy & Reaction Heat Simulator (2D)

Reaction Type

Temperature T 400 K

Activation energy E_a 60 kJ/mol

Reaction enthalpy |ΔH| 50 kJ/mol

Fraction E ≥ E_a

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Rate constant k

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ΔH (this reaction)

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Swarm success rate

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How it works

Collision theory: molecules only react when they collide with enough kinetic energy to clear the activation barrier E_a. In two dimensions the Maxwell–Boltzmann kinetic-energy distribution is exactly exponential, so the fraction of collisions energetic enough to react is given cleanly by the Boltzmann factor below — that same factor is the exponential term in the Arrhenius equation for the rate constant k.

f(E ≥ Ea) = exp(-Ea / RT)
k = A · exp(-Ea / RT)   (Arrhenius equation)
R = 8.314 J/(mol·K)

Temperature T — raising T shifts more of the energy distribution above E_a, so more particles cross the hump per second and k rises steeply (not linearly).
Activation energy E_a — the height of the barrier the reaction coordinate must climb; a taller barrier exponentially suppresses the successful-collision fraction.
Catalyst — opens a lower-E_a pathway without changing the energy of reactants or products, so it speeds the reaction up without touching ΔH — the barrier height drops, but the two plateaus stay put.
ΔH (exothermic vs endothermic) — whether products sit lower (heat released, particles flash hot orange crossing the hump) or higher (heat absorbed from the surroundings) than the reactants. Only the plateau heights change; the shape of the barrier is independent of ΔH.

Real-world relevance: this is why refrigerated food spoils slower (lower T → far fewer collisions clear E_a) and why catalytic converters and enzymes speed reactions up by lowering E_a rather than by heating anything.

Reactant molecule Transition state Product (exothermic) Product (endothermic)