Effects of changes in surface area of a solid, concentration of a solution

Some reactions are fast (a firework explodes in a fraction of a second); some are slow (an iron nail rusts over weeks). The rate of reaction tells us how fast a reaction goes.

Definition (learn this): the rate of reaction is the change in the amount (or concentration) of a reactant or product per unit time.

In practice you measure something that changes as the reaction happens, and divide by the time taken:

• volume of gas produced per second (cm³/s),
• mass lost per second as a gas escapes (g/s),
• time for a fixed change to happen (e.g. for a cross to disappear) — here rate ∝ 1/time.

$\text{rate} = \frac{\text{change in amount of reactant or product}}{\text{time taken}}$

A reaction is fast if a lot changes in a short time, and slow if little changes over a long time. Everything in this subtopic is about how to make a reaction go faster — and why, using collision theory.

Reactions happen because particles collide with one another. But not every collision causes a reaction — most collisions just bounce apart unchanged.

Three requirements for a reaction to happen:

1. The particles must COLLIDE — they have to physically meet.
2. They must collide with enough energy — at least the activation energy (the minimum energy needed to start the reaction).
3. They must collide with the correct orientation (the right way round).

A collision that meets all three is a successful collision — only these make products.

The golden rule: the rate of reaction depends on how often successful collisions happen — the number of successful collisions per second.

So to make a reaction faster you must make successful collisions happen more frequently. Increasing concentration, surface area and temperature all do exactly this — that is why they speed reactions up.

Increasing the concentration of a solution increases the rate of reaction.

Explanation (collision theory): a more concentrated solution has more reactant particles in the same volume. The particles are therefore closer together, so they collide more often — more collisions per second means more successful collisions per second, so the reaction is faster.

Note: the temperature is unchanged, so the proportion of collisions that succeed is the same — there are simply more collisions happening.

A classic example — magnesium with hydrochloric acid:

$\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$

Using a more concentrated acid makes the magnesium fizz faster and the hydrogen gas is given off more quickly.

Gases — pressure works the same way. For reactions between gases, increasing the pressure squashes the same particles into a smaller volume, so they are closer together and collide more often. Higher pressure → faster rate, for the same reason as concentration.

Watch out: pressure has no effect on reactions of solids or solutions — only on gases.

When a solid reacts with a solution or a gas, the reaction can only happen at the surface of the solid — the particles deep inside cannot be reached.

Breaking a solid into smaller pieces increases its surface area, so more particles are exposed. This means more frequent collisions with the other reactant → more successful collisions per second → faster rate. A fine powder reacts fastest of all.

Why smaller pieces have more surface area (same mass): picture a 1 cm cube cut into eight smaller 0.5 cm cubes:

• 1 cm cube → surface area = 6 × (1 × 1) = 6 cm².
• eight 0.5 cm cubes → surface area = 8 × 6 × (0.5 × 0.5) = 8 × 1.5 = 12 cm².

Same total volume and mass, but double the surface area — and a real powder has thousands of times more surface area than a single lump.

Classic example — marble chips (calcium carbonate) with hydrochloric acid:

$\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2$

Large chips react slowly; small chips react faster; powder reacts almost instantly — because the surface area exposed to the acid keeps increasing.

Fair test (control variables): when investigating surface area, keep the mass of solid, volume and concentration of acid, and temperature the same. Change only the size of the pieces.

In experiments you often plot volume of gas (or mass lost) against time. The shape of the curve tells you everything about the rate.

How to read it:

• The steepness (gradient) of the line = the rate. A steeper line means a faster reaction.
• The steepest part is at the start — there is the most reactant, so collisions are most frequent.
• The line gets less steep over time because the reactant is being used up → fewer particles → fewer collisions per second.
• The line levels off (plateaus) when a reactant is completely used up → the reaction has stopped (rate = 0).
• If two runs use the same amounts of reactant, they level off at the same height — increasing concentration or surface area makes the reaction faster (steeper, finishes sooner) but does not change the final amount of product.