Reaction Rate and Activation Energy Lab

Molecules do not react just because they collide. They must collide hard enough, with energy exceeding the activation energy (Ea) — the barrier that must be crossed to break existing bonds.

The energy distribution

In any sample, molecules share out energy unevenly — the familiar bell curve. Only the shaded tail beyond Ea has enough to react. This is why two of the three main ways to speed up a reaction work the way they do:

• Raise the temperature — the whole curve shifts slightly and flattens, but the tail beyond Ea grows enormously. A 10 K rise can double a rate.
• Add a catalyst — it does not give molecules more energy. It lowers the barrier, so a much larger fraction of the unchanged curve already qualifies.
• Raise concentration — more collisions per second, but the fraction that succeed is unchanged.

The Arrhenius equation

k = A · e^−Ea/RT

The exponential is where the sensitivity lives. Because Ea sits in the exponent with R (8.314 J/mol·K) and T below it, a reaction with a high barrier is far more temperature-sensitive than one with a low barrier. That is why a bump of ten degrees barely changes some reactions and doubles others.

Try this

• Note the tiny fraction of molecules that qualify at room temperature — often well under one percent.
• Raise T by 10 K and watch that fraction multiply, even though the curve barely moved.
• Apply the catalyst: the curve stays put, the barrier drops, and the qualifying fraction jumps.
• Set ΔH positive (endothermic) and check that the reverse barrier is now larger than the forward one.

Why it matters

Activation energy explains why petrol needs a spark despite being surrounded by oxygen, why food lasts far longer in a freezer, why enzymes are so effective (they are catalysts tuned by evolution), and why the catalytic converter in a car needs to heat up before it starts working.