🧪 Activation Energy and Collision Theory

About this simulation

Collision theory says a chemical reaction only happens when reactant molecules collide with enough kinetic energy to break existing bonds and reach the transition state — the energy barrier called the activation energy (Ea). Most collisions in a gas or liquid are gentle glancing bounces that fail to clear this barrier; only the energetic minority in the high-speed tail of the molecular speed distribution react.

How it works

• Each sphere is a molecule; its color shifts from blue (slow) to red (fast) following a Maxwell–Boltzmann-like speed spread that widens as temperature rises.
• When two molecules touch, the simulation computes the kinetic energy of that collision along the line of centres and compares it with Ea.
• A gold flash marks a successful, energy-sufficient collision; a dim grey flash marks an ordinary elastic bounce.
• The energy diagram plots each recent collision's energy against the Ea barrier as a scrolling strip of dots.

Controls

• Temperature — raises average molecular speed, widening the energy distribution and pushing more collisions above Ea.
• Concentration — number of molecules in the box; more molecules means more collisions per second, but not a higher success rate.
• Activation energy (Ea) — raises or lowers the barrier collisions must clear to react.
• Catalyst toggle — opens an alternative, lower-energy pathway, shown as a dashed lower hump on the diagram.
• ↺ Reset molecules — respawns the box. Drag to rotate, scroll to zoom.

Did you know?

A 10°C rise in temperature roughly doubles many reaction rates near room temperature — not because molecules collide much more often, but because a much larger fraction of collisions now carry enough energy to clear Ea, exactly as predicted by the Arrhenius equation.