Reaction rates (intro)
The idea
Some reactions finish in a flash — an airbag inflates in milliseconds — while others, like iron rusting, take years. Reaction rate measures the speed: how fast a reactant's concentration falls or a product's concentration rises per unit time, typically in M/s. Collision theory explains what sets the pace, building on your particle model: particles must collide, with enough energy to rearrange bonds (the activation energy) and in a workable orientation, before they can react.
Every rate-boosting factor works by improving collisions. Higher concentration packs more particles into the same space, so collisions come more often. Higher temperature makes particles move faster, producing both more frequent and more energetic collisions — a double effect, which is why temperature changes rate so steeply. Grinding a solid raises surface area, exposing more particles to attack. A catalyst is subtler: it provides an alternative pathway with a lower activation energy, so a larger fraction of the existing collisions succeed.
Two corrections about catalysts: they are not consumed — the same catalyst molecules emerge unchanged and keep working — and they do not create more product; they only deliver the same product sooner. Also expect rates to change over time: reactions typically start fast and slow down as reactants are used up and collisions become rarer.
Worked example
Hydrogen peroxide decomposes by 2 H₂O₂ → 2 H₂O + O₂. In a flask, the H₂O₂ concentration falls from 0.800 M to 0.200 M over 150 s. What is the average rate of disappearance of H₂O₂ over this interval?
- Find the change in concentration: 0.800 − 0.200 = 0.600 M of H₂O₂ was consumed during the interval.
- Divide the change by the elapsed time to get the average rate: 0.600 M ÷ 150 s = 0.00400 M/s, or 4.00 × 10⁻³ M/s.
- Mind the convention: the concentration change is negative for a reactant (it is disappearing), but rates are reported as positive numbers, so the sign is dropped when stating a rate of disappearance.
- Interpret what 'average' hides: the reaction ran fastest near the start, when the concentration — and so the collision frequency — was highest, and slower near the end; the instantaneous rate at 150 s is below this 0.00400 M/s average.
Answer. The average rate of disappearance of H₂O₂ is 4.00 × 10⁻³ M/s over the 150 s interval.
Check your understanding
- Why does a modest temperature increase speed a reaction so dramatically, given the two effects it has on collisions?
- Why does crushing a solid reactant into powder accelerate its reaction, in collision terms?
- Why do most reactions slow down as they proceed, and what would a graph of concentration versus time look like because of it?
- How can a catalyst make a reaction faster without being used up or changing how much product forms?
Build the foundations first
Reaction rates (intro) builds on these concepts. If any feel shaky, start there.