Sec 4 Chemistry Rate of Reactions notes

Rate is always how much changes, divided by how long it takes: rate = Δquantity ÷ Δtime. For a gas-producing reaction (marble in acid), collect the gas in a gas syringe and plot volume against time; for a reaction that loses gas to the air (carbonate + acid on an open balance), plot mass loss against time; for the sodium thiosulfate “disappearing cross”, time how long the precipitate takes to hide a cross beneath the beaker.

Reactions happen when reactant particles collide — but only collisions carrying at least the activation energy succeed. Every rate factor therefore works through just two levers: how often particles collide (collision frequency) and what fraction of collisions clear Eₐ (collision energy).

This is the explanation template for every factor: name the factor’s physical effect, convert it to collision frequency or collision energy, then conclude the rate change. More frequent successful collisions each second mean a faster reaction.

A more concentrated solution packs more particles into each dm³, so collisions per second rise and the rate climbs. The same logic compressed: for gases, raising the pressure squeezes the same particles into less volume — same result, faster reaction. (A classic hospital-rate analogy: a crowded corridor produces more shoulder-bumps per second than an empty one.)

Particle size is surface area in disguise: one big lump hides most of its particles inside, while grinding it into powder exposes nearly all of them. Marble chips fizzle politely in hydrochloric acid; the same mass of marble powder froths violently — then both curves level at the same final volume, because the amount of CaCO₃ never changed.

A 10 °C rise can roughly double a reaction rate. Temperature works through both levers: particles move faster (slightly more frequent collisions) and — the big one — a much larger fraction of collisions now exceeds Eₐ. The energy shift beats the frequency shift, which is why heating is the chemist’s favourite accelerator.

A catalyst offers a lower-energy route: Eₐ drops, so a bigger share of collisions succeed at the same temperature. Manganese(IV) oxide, MnO₂, decomposes hydrogen peroxide at room temperature: 2H₂O₂ → 2H₂O + O₂ — the splint test confirms oxygen roaring out, and the black powder is recovered unchanged afterwards.

Light is the rarest rate factor: photochemical reactions use photons as their energy source. Silver bromide in film darkens in light (2AgBr → 2Ag + Br₂), which is why film is sold in black canisters; photosynthesis runs on sunlight and stops in the dark.

On a gas volume–time curve, the steeper the gradient, the faster the reaction. Every curve starts steep and flattens as reactants are used up — collisions keep happening but successful ones become rare. The flat region means the reaction is over: the limiting reactant is exhausted.

Compare curves like a forensic scientist: same plateau, steeper start same amount, faster conditions (powder vs chips, hotter vs colder, concentrated vs dilute). Different plateaus different amounts of limiting reactant. To turn a plateau into moles, use the r.t.p. molar gas volume: 24 dm³ = 24 000 cm³ per mole.