Effect of Surface Area on Rate of Reaction

When a solid reacts with a liquid (like an acid) or a gas, the reaction can only happen where the two reactants actually touch — at the surface of the solid.

The particles inside the solid are surrounded on all sides by other solid particles. The acid cannot reach them, so they cannot react until the particles covering them have reacted away first.

This single idea is the key to the whole topic: more surface = more particles exposed to the acid = more places where a reaction can happen at once.

In the diagram, the grey particles inside the lump are locked away from the acid. When the solid is broken into smaller pieces, those once-buried particles end up on a new surface — so many more particles can react at the same time.

For a reaction to happen, particles of the two reactants must collide. The rate depends on how often successful collisions happen each second.

The mark-scheme explanation:

• Increasing the surface area exposes more solid particles to the acid.
• This means there are more frequent collisions between the acid particles and the solid particles each second.
• More collisions per second → more successful collisions per second → faster rate.

Notice what does not change: the temperature is the same and the acid concentration is the same, so the proportion of collisions that succeed is unchanged. The surface area only changes the number of collisions happening at the surface each second.

Powder is the extreme case. Grinding a solid to a fine powder exposes a huge number of particles, so a powder reacts dramatically faster than a single large lump of the same mass — sometimes finishing in seconds rather than minutes.

Three-step answer to learn: more surface area → more particles exposed → more frequent collisions per second → faster rate.

The reason powder reacts faster is that smaller pieces have a larger total surface area for the same mass (and same volume). This is the surface-area-to-volume ratio getting bigger.

A worked number to fix the idea. Compare one large cube with eight small cubes of the same total volume:

	Surface area	Volume
One 2 cm cube	6 × (2 × 2) = 24 cm²	2 × 2 × 2 = 8 cm³
Eight 1 cm cubes	8 × 6 × (1 × 1) = 48 cm²	8 × 1 = 8 cm³

Cutting the cube into eight pieces doubled the surface area without changing the mass or volume. Keep subdividing and the surface area grows and grows — a fine powder of the same mass can have thousands of times the surface area of a single lump.

So although large chips, small chips and powder all have the same mass of calcium carbonate, the powder has by far the largest surface area — and therefore reacts fastest.

The standard 4WSD0/1C investigation uses marble chips (calcium carbonate, CaCO₃) reacting with dilute hydrochloric acid:

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

The carbon dioxide gas given off lets you follow the rate. You compare the same mass of CaCO₃ in three forms: large chips, small chips and powder.

Two ways to measure the rate:

• Gas volume — fit the flask with a bung and delivery tube to a gas syringe; record the volume of CO₂ every 10 seconds.
• Mass loss — stand the flask on an electronic balance and plug the neck with cotton wool (lets CO₂ escape, stops acid spray); record the falling mass every 10 seconds as CO₂ escapes.

Control variables (to make it a fair test) — change ONLY the surface area:

• Same mass of calcium carbonate (e.g. 5.0 g).
• Same volume of acid (e.g. 50 cm³).
• Same concentration of acid (e.g. 1.0 mol/dm³).
• Same temperature.
• Same size and shape of flask.

Reading the graph:

• The steepest initial gradient = the fastest rate. Powder is steepest, large chips shallowest.
• Powder reaches the plateau first (finishes soonest); large chips finish last.
• Crucially, all three level off at the SAME final volume of CO₂ (and the same total mass lost). The same mass of CaCO₃ produces the same total amount of gas no matter how it is broken up — surface area changes the speed, never the amount.