Sec 4 Combined Science (Chemistry) Speed of Reactions notes

The rate of reaction is how quickly reactants turn into products: the amount of product formed, or reactant used up, per unit time. In the lab you usually follow one easy-to-measure quantity: the volume of gas collected (gas syringe), the mass of the mixture as gas escapes (balance), or the time for a solid to disappear or a colour to settle.

Rates can be quoted as a mean: if 48 cm³ of gas collects in the first 30 seconds, the mean rate over that stretch is 48 ÷ 30 = 1.6 cm³/s. A full curve tells you more: fast at the start, slowing as reactants run low.

For a reaction to happen, particles must collide with each other with at least enough energy to react. That is collision theory in one sentence. It gives two levers, and only two: increase the number of collisions per second, or increase the energy the collisions carry. Every rate factor in this chapter is one of these two levers.

The energy part explains why some mixtures just sit there: most collisions are too gentle to react. Warming the mixture makes particles move faster, so collisions happen more often and each one hits harder — which is why temperature is the strongest lever of the four.

Particle size only matters when a solid reacts with a liquid or gas. Only particles on the surface can be hit, so breaking a solid into smaller pieces increases the surface area — the same number of particles, but far more of them exposed and reachable. Powder has the largest surface area and reacts fastest; one large lump reacts slowest. (The total amount of solid is unchanged — only its exposure changes.)

Temperature raises the average speed of all particles, solution or gas. Faster particles collide more frequently, and a bigger share of collisions now carry enough energy to react — two effects pulling the same way. As a rough classroom guide, many reactions about double in rate for every 10 °C rise — from 20 °C to 50 °C is three 10-degree steps, so roughly eight times faster.

Concentration counts particles per volume. A more concentrated acid solution packs more acid particles into every cm³, so the solid's surface particles are hit more often each second — more frequent collisions, faster rate. The same logic in the gas phase is called pressure: compressing a gas into a smaller volume crowds its particles together, so gas–gas collisions become more frequent and the reaction speeds up.

• Concentration is the solution lever: 2.0 mol/dm³ acid reacts faster than 1.0 mol/dm³ acid (same volume, same temperature).
• Pressure is the gas lever: relevant to gas-phase reactions and in engines and industrial plant.
• Neither factor changes how energetic single collisions are — only how often they happen.

A fair-test detail examiners love: when comparing concentrations, keep the total volume, temperature, and solid surface area identical — otherwise you have changed two things at once and the comparison is meaningless.

A gas-volume graph (volume collected against time) has three zones: a steep start (many collisions, plenty of reactant — fast), a bend as reactants are used up (rate slows), and a flat line when the limiting reactant is gone — no more gas, reaction over. The steeper the rise, the faster the rate; the flat height tells you the total gas made (how much reactant there was).

A mass-loss graph works in mirror image: the balance reading falls as gas escapes, steepest at the start, then levels at a constant final mass. And comparison graphs — one curve per concentration — let you judge factors at a glance: the higher concentration curve is steeper and finishes sooner, but if the same amount of marble reacts it flattens at the same final volume.