Why does a chemical reaction sometimes happen in a flash?
Ever wonder why mixing baking soda and vinegar fizzes instantly, while rusting takes months? That speed difference is what chemical kinetics studies.
💡 In Simple Words: The rate of reaction tells us how fast reactants turn into products. Think of it like how quickly water flows through a pipe – the faster the flow, the quicker the bucket fills.
What is reaction rate?
Reaction rate is the change in concentration of a reactant or product per unit time. Concentration means amount of substance in a given volume, usually expressed in moles per litre (M). So, if the concentration of a reactant drops by 0.2 M in 10 seconds, the rate is 0.02 M s⁻¹.
How to calculate the rate of a reaction
For a simple reaction aA + bB → cC + dD, the rate can be written using any species:
- Rate = –(1/a) Δ[A]/Δt (negative because reactant concentration falls)
- Rate = +(1/c) Δ[C]/Δt (positive for product formation)
Pick the one that’s easiest to measure in the lab.
Units of reaction rate
Because concentration is in moles per litre (M) and time in seconds, the standard unit is mol L⁻¹ s⁻¹. Some textbooks write it as M s⁻¹ – they’re the same thing.
Factors that speed up or slow down a reaction
Just like a recipe, the outcome changes if you tweak the ingredients or cooking conditions. The main knobs you can turn are:
- Concentration – More particles mean more collisions, so the reaction usually goes faster.
- Temperature – Heating gives particles extra energy, making collisions more forceful and frequent.
- Surface area – If a solid is broken into powder, each tiny piece offers a fresh surface for collisions.
- Catalyst – A substance that provides an alternative pathway with a lower activation energy (the hill particles must climb to react).
- Nature of reactants – Some bonds are easier to break; for example, reactions involving gases often proceed quicker than those involving solids.
Rate law and order of reaction
The rate law connects the measurable rate to the concentrations of the reactants. It looks like:
Rate = k [A]^m [B]^n
Here, k is the rate constant – a number that depends on temperature and the specific reaction. The exponents m and n are called reaction orders; they tell us how the rate changes when a concentration changes. If m = 1, doubling [A] doubles the rate. If m = 0, the rate doesn’t care about [A] at all.
How to find the rate law experimentally
Usually you can’t guess the orders from the balanced equation. You have to do experiments:
Quick comparison table
| Aspect | What it means |
|---|---|
| Reaction rate | Δ(concentration)/Δtime |
| Rate constant (k) | Proportionality factor, depends on temperature |
| Order (m, n) | How strongly rate depends on each reactant |
| Activation energy | Energy hill particles must climb; lowered by a catalyst |
Worked example
Consider the reaction 2NO₂ → 2NO + O₂. An experiment gives the following initial rates:
| [NO₂] (M) | Initial rate (M s⁻¹) |
|---|---|
| 0.10 | 2.0 × 10⁻⁴ |
| 0.20 | 8.0 × 10⁻⁴ |
When the concentration doubles, the rate quadruples. That tells us the reaction is second order in NO₂ (rate ∝ [NO₂]²). So the rate law is Rate = k[NO₂]². Plug in one data point to find k:
2.0 × 10⁻⁴ = k(0.10)² → k = 2.0 × 10⁻⁴ / 0.01 = 0.020 M⁻¹ s⁻¹.
Key take‑aways
- Rate = change in concentration over time, usually expressed in M s⁻¹.
- Higher concentration, temperature, surface area, or a catalyst generally increase the rate.
- The rate law (Rate = k[A]^m[B]^n) is discovered experimentally, not from the balanced equation.
- Reaction order tells you how the rate responds to concentration changes.
- Rate constant k grows with temperature; the Arrhenius equation (k = A e^(-Ea/RT)) describes this relationship.
📝 Likely Exam Questions
- Define reaction rate and write its unit.
- Explain how temperature influences the rate constant.
- For the reaction 2A + B → C, the initial rates are: [A]=0.1 M, [B]=0.1 M, rate=1.0 × 10⁻⁴ M s⁻¹; [A]=0.2 M, [B]=0.1 M, rate=4.0 × 10⁻⁴ M s⁻¹. Determine the order with respect to A and write the rate law.
- State two ways a catalyst speeds up a reaction.
- Why can’t the stoichiometric coefficients be directly used as reaction orders?