Why does a refrigerator keep food fresh?
Food spoils more slowly in the fridge. A lump of coal burns slowly, but coal dust in a mine can explode. Hydrogen peroxide fizzes violently as soon as manganese dioxide is added (1.7.2). The same reaction can run at very different speeds under different conditions.
Factors that affect reaction rate
The rate depends first on the nature of the reactants: sodium reacts violently with cold water, while magnesium barely reacts. For a given reaction, these conditions also matter:
- Concentration: raising the concentration of a reactant speeds up the reaction.
- Pressure: for reactions involving gases, compressing the gas raises the pressure and the concentration, so the reaction speeds up. Pressure has almost no effect on solids and liquids.
- Temperature: raising the temperature speeds up the reaction. A rough rule of thumb: for many reactions, a 10 °C rise makes the rate 2 to 4 times faster.
- Catalysts: a catalyst can speed up a reaction enormously. The enzymes in living things are very efficient catalysts.
- Surface area: the smaller the pieces of a solid, the more surface is exposed to the other reactants and the faster the reaction. Light and stirring speed up some reactions too.
Explaining the factors with collisions
A reaction’s rate depends on how many effective collisions happen per second:
- Higher concentration or pressure: more molecules per liter means more collisions per second, so more effective collisions. The fraction of collisions with enough energy doesn’t change.
- Higher temperature: molecules move faster, so they collide more often. More importantly, a larger fraction of them have energy above Ea, so a larger fraction of collisions are effective.
- Catalyst: a catalyst provides a different pathway with a lower activation energy, so a larger fraction of collisions have enough energy. It speeds up the forward and reverse reactions equally.
- Surface area: for a solid, only particles at the surface can collide with the other reactant; more surface means more collisions.
Compare the energy distributions at two temperatures. The shaded area is the fraction of molecules with energy above Ea. Raise the temperature or add a catalyst and watch it change:
Schematic: the activation energies are far smaller than in real reactions so the shaded areas are visible. The fractions are calculated exactly from the energy distribution.
Controlling variables
To study one factor, keep every other factor the same: this is a controlled experiment. The lab below uses sodium thiosulfate and dilute sulfuric acid:
The sulfur makes the solution cloudy. Draw an X on paper under the beaker and time how long until you can’t see it from above: the shorter the time, the faster the reaction. Change the thiosulfate concentration or the temperature, one at a time, and record your results:
| [Na₂S₂O₃] | Temperature | Time for the X to vanish | 1/t × 1000 |
|---|
Model data: calculated from a simple model (t inversely proportional to [Na₂S₂O₃], activation energy about 50 kJ/mol). The sulfuric acid concentration stays the same.
How Chinese textbooks differ
- Chinese textbooks list pressure as its own factor and explain every factor in terms of the number and the percentage of “activated molecules” (活化分子) per unit volume. A raised concentration increases their number but not their percentage; a higher temperature or a catalyst increases both.
- Chinese classes run the same disappearing-X experiment. They also compare FeCl₃ and CuSO₄ as catalysts for H₂O₂ decomposition, and time how fast acidified KMnO₄ loses its color with oxalic acid.
Example 1: designing controlled experiments
Three trials of the thiosulfate reaction are run:
| Trial | [Na₂S₂O₃] (M) | [H₂SO₄] (M) | Temperature (°C) |
|---|---|---|---|
| 1 | 0.10 | 0.10 | 20 |
| 2 | 0.20 | 0.10 | 20 |
| 3 | 0.10 | 0.10 | 40 |
(a) Which two trials show the effect of concentration? (b) Of temperature? (c) Predict which trial hides the X first.
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Why watch an X? The reaction forms solid sulfur, which clouds the solution until an X drawn under the flask disappears:
The sooner the X disappears, the faster the reaction.
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Control variables. Change one condition at a time and keep the others the same; only then can you tell what that condition does.
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Concentration. Compare trials 1 and 2: only [Na₂S₂O₃] differs.
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Temperature. Compare trials 1 and 3: only the temperature differs. Trials 2 and 3 differ in two ways, so they can’t be compared.
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Prediction. Doubling the concentration roughly doubles the rate in trial 2. Trial 3 is 20 °C warmer, which is two 10 °C steps; at 2 to 4 times per step:
so it is 4 to 16 times faster. Trial 3 hides the X first.
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A common mistake. Adding the factors (2 + 2 = 4 to 4 + 4 = 8 times). Each 10 °C step multiplies the rate again.
Example 2: explaining with collision theory
Use collision theory to explain why (a) coal dust burns faster than a lump of coal and (b) a little manganese dioxide makes hydrogen peroxide decompose much faster.
- The key idea. The rate depends on the number of effective collisions per second. You can raise it two ways: more collisions in total, or a larger fraction of collisions with enough energy to react.
- Coal dust. Fine particles expose far more surface to oxygen, so more carbon atoms can collide with O₂ molecules each second. More effective collisions means faster burning. The activation energy hasn’t changed; the number of collisions has.
- Manganese dioxide. MnO₂ is a catalyst. It provides a pathway with a lower activation energy, so a much larger fraction of collisions have enough energy to react.
- The catalyst itself. At the end, the MnO₂ has the same mass and chemical properties as before (1.7.2), even though it took part in the reaction pathway.
- Compare. More surface area or higher concentration means more collisions; a catalyst or a higher temperature means a larger fraction of successful collisions (warming also makes collisions a little more frequent).
- Using powdered marble instead of marble chips with hydrochloric acid
More surface area.
- Using 1 M HCl instead of 0.1 M HCl with zinc
Higher [H⁺].
- Putting food in the refrigerator
A lower temperature slows the reactions that spoil food.
- Adding a little MnO₂ to hydrogen peroxide
MnO₂ is a catalyst that lowers the activation energy.
- Adding helium to a rigid container where N₂ + 3H₂ ⇌ 2NH₃ is running
The volume is fixed, so [N₂] and [H₂] don’t change.
- Compressing the container for N₂ + 3H₂ ⇌ 2NH₃
[N₂] and [H₂] increase.
- Adding water to zinc reacting with dilute sulfuric acid
[H⁺] decreases.
- Adding a little solid NaCl to iron reacting with hydrochloric acid
NaCl doesn’t take part, and [H⁺] doesn’t change.
- Warming the thiosulfate reaction in a hot-water bath
A larger fraction of molecules have energy above Ea.
Common mistakes
Mistake 1: “Higher concentration makes each collision more likely to succeed.” It only makes collisions more frequent. The fraction of collisions with enough energy depends on temperature and Ea, not on concentration.
Mistake 2: “A catalyst doesn’t take part in the reaction.” A catalyst does take part, changing the pathway. It is simply unchanged at the end. It also speeds up the reverse reaction just as much.
Mistake 3: “Raising the pressure always speeds up a reaction.” Only if it raises the concentration of a gaseous reactant. Adding an unreactive gas to a rigid container raises the total pressure but leaves every reactant concentration, and so the rate, unchanged.
Quick check
Which change increases the fraction of collisions that have enough energy to react?
A higher temperature means more molecules have energy above Ea. The other changes only make collisions more frequent.
At 20 °C the X disappears after 40 s. If the rate roughly doubles for every 10 °C, about how long will it take at 30 °C?
Doubling the rate halves the time: 40 s ÷ 2 = 20 s.
Summary
- Higher concentration, higher gas pressure, higher temperature, a catalyst, and greater surface area all speed up a reaction.
- Concentration, pressure, and surface area raise the number of collisions; temperature and catalysts raise the fraction of collisions with enough energy.
- To study one factor, change only that factor and keep everything else the same.
On to AP (for later)
AP topic 5.1 asks you to name the factors that affect rate and compare rates from experimental data. Topic 5.5 uses the collision model and the energy-distribution graph to explain temperature: heating increases the fraction of molecules with energy above Ea. Topic 5.11 covers catalysts, which provide a new pathway (mechanism) with a lower activation energy and are regenerated at the end.