Question

How does lowering activation energy affect temperature required to start a chemical reaction

Answer 1

Lowering the activation energy (Ea) makes a reaction proceed faster at a given temperature — so you can reach a given reaction rate (or "start" the reaction) at a lower temperature.

Why: reaction rates follow the Arrhenius law
k = A · exp(−Ea/(R T)),
where k is the rate constant, A the pre‑exponential factor, R the gas constant, and T the temperature (K). Because Ea is in the exponential, even modest reductions in Ea give exponentially larger rates at the same T.

Two useful ways to see the effect

- For a fixed temperature, the rate increases by the factor
k2/k1 = exp[(Ea1 − Ea2)/(R T)].
Example: lowering Ea by 20 kJ/mol at 300 K increases the rate by exp(20000/(8.314·300)) ≈ 3.0×10^3.

- To get the same rate with a lower Ea, you can solve for the new temperature (assuming A doesn’t change):
Ea1/(R T1) = Ea2/(R T2) ⇒ T2 = (Ea2/Ea1) · T1.
Example: if Ea is reduced from 100 kJ/mol to 80 kJ/mol, the temperature required for the same rate drops to 0.8·T1 (so 600 K → 480 K).

Molecular picture: the fraction of molecules with enough energy to react scales roughly as exp(−Ea/(R T)), so lowering Ea greatly increases that fraction at a given T.

Practical note: catalysts work by lowering Ea, allowing reactions to run at much lower temperatures. This doesn’t change thermodynamic equilibrium (ΔG) — it only changes how quickly the equilibrium is reached. Also, in real systems A and multi‑step mechanisms can matter, but the exponential dependence above gives the dominant effect.