Chemical Kinetics: Collision Theory & Arrhenius Reaction

Reaction Vessel

Temperature 400 K

Activation energy 60 kJ/mol

Concentration 30 / species

[A] remaining

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[B] remaining

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[C] formed

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Reaction rate

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Collision freq.

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Rate constant k

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k = A·e^−Ea/RT. Not every collision reacts — only ones with enough energy to cross the activation barrier, drawn probabilistically from that same exponential each time A meets B.

How it works

Blue spheres are reactant A, orange spheres are reactant B; both bounce elastically inside the vessel, their speed set by temperature (kinetic theory: mean speed scales with √T). Whenever an A and a B molecule touch, that's a "collision" in the collision-theory sense — but only a fraction of collisions actually cross the activation-energy barrier and react. That fraction is exactly the Arrhenius factor: on every A–B contact, the sim draws a random number and reacts only if it falls under e^−Ea/RT. Every reacted pair turns green (product C) and stops reacting further.

rate = k·[A]·[B]
k = A_factor · exp(−Ea / R·T)   (Arrhenius equation)
R = 8.314 J/(mol·K)

Temperature — raises molecular speed (more collisions per second) and raises the reacting fraction of those collisions (bigger exponential factor) — the two effects Arrhenius bundles into one equation.
Activation energy (Ea) — the energy barrier a colliding pair must clear to react; raise it and almost every collision bounces off unreacted instead.
Concentration — more A and B molecules in the same volume means more collisions per second, raising the rate even at fixed temperature (rate ∝ [A][B]).
Catalyst — provides an alternative reaction pathway with a lower activation energy (transition-state theory: a lower-energy transition state), without being consumed itself.

The measured rate constant k below is back-calculated from the actual reaction count, collision count, [A] and [B] each frame, so you can watch it track the theoretical Arrhenius curve as you drag the sliders.

Reactant A Reactant B Product C

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