Activation Energy & Reaction Heat Simulator

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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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 on the diagram 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, particles cool to blue) 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)

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