HookEvery unopened cola bottle is a finished chemistry experiment
A sealed bottle of cola holds roughly three to four times its own volume of carbon dioxide, forced into the liquid at a pressure of a few atmospheres. On the shelf it looks like the least interesting object in the shop — nothing moves, nothing changes. But at the molecular level it is frantic: CO₂ molecules are constantly escaping from the liquid into the headspace, and CO₂ molecules in the headspace are constantly dissolving back in, at exactly equal rates. The reaction \(\mathrm{CO_2(g)} \rightleftharpoons \mathrm{CO_2(aq)}\) has reached dynamic equilibrium — a standstill made of two opposite processes running flat out. Crack the cap and you hear the experiment being interrupted: the pressure drops, the balance breaks, and the system responds exactly the way this section teaches you to predict.
C6 asks two separate questions, and keeping them separate is half the topic. How fast does a reaction go — that is rate, and it is explained by collision theory. How far does it go — that is extent, and for reversible reactions the answer is an equilibrium position that shifts when you change the conditions. Warm cola demonstrates both at once: warm it and the gas escapes faster (rate); leave it open and the gas keeps leaving until the drink is flat (extent — an open bottle is an open system, so equilibrium is never re-established). Every mark in this section is one of those two dials.
ModelMeasuring rate — mean rate, tangents and the steep start
The rate of reaction is how quickly a reactant is used up or a product is formed. You cannot see molecules reacting, so you track something measurable that changes as they do: the mass a flask loses as gas escapes (on a balance), the volume of gas collected (gas syringe or upturned measuring cylinder), or how quickly a solution turns cloudy.
The mean rate over any period is simply \(\text{mean rate} = \dfrac{\text{quantity changed}}{\text{time taken}}\), with units that name what you measured: g/s if you followed mass, cm³/s if you followed gas volume — and on the Higher tier, mol/s once you convert using relative formula mass or molar gas volume. On a graph of product against time, the curve is steepest at the start, when reactant particles are at their most concentrated, and it flattens to horizontal the moment one reactant runs out. Because the rate changes from second to second, the rate at a particular instant is the gradient of a tangent — a ruler-drawn straight line that just touches the curve at that time. Higher-tier questions ask you to calculate that gradient explicitly.
Marble chips (calcium carbonate) react with dilute hydrochloric acid in a flask on a balance, with a cotton-wool plug to let gas out but keep acid spray in. The mass falls by 0.66 g in the first 120 s as CO₂ escapes. Mean rate = 0.66 ÷ 120 = 0.0055 g/s. Higher tier: the relative formula mass of CO₂ is 44, so 0.66 g is 0.66 ÷ 44 = 0.015 mol, giving 0.015 ÷ 120 = \(1.25 \times 10^{-4}\) mol/s. For the rate at exactly 60 s, draw a tangent touching the curve at 60 s; if it passes through (20 s, 0.30 g) and (100 s, 0.62 g), the gradient is (0.62 − 0.30) ÷ (100 − 20) = 0.32 ÷ 80 = 0.004 g/s. State the units at every step — they are usually a mark on their own.
ModelCollision theory — one sentence that explains every factor
Here is the sentence: particles only react when they collide with energy equal to or greater than the activation energy — the minimum energy needed for a collision to lead to a reaction. Everything that changes a rate changes either how often particles collide or how energetic those collisions are.
Raise the concentration of a solution (or the pressure of a gas) and you pack more particles into the same volume, so collisions happen more frequently — roughly double the concentration, double the rate. Increase the surface area of a solid by crushing it and you expose vastly more particles to attack: powdered marble fizzes furiously where a single lump merely bubbles, because rate scales with the surface-area-to-volume ratio. The same physics has a body count — on 2 May 1878 the Washburn 'A' Mill in Minneapolis, the largest flour mill in America of its day, was destroyed when airborne flour dust ignited; finely divided, an everyday foodstuff becomes an explosive because almost every particle can burn at once.
Temperature is the double lever, and examiners test whether you know both halves. Hotter particles move faster, so they collide more frequently — but the far bigger effect is that a much greater proportion of collisions now carry at least the activation energy. That is why a modest temperature rise has a dramatic effect: as a rule of thumb, many reactions roughly double their rate for every 10 °C increase.
MechanismCatalysts and enzymes — lowering the bar, not pushing harder
A catalyst increases the rate of a reaction without being used up, and it does it by providing an alternative reaction pathway with a lower activation energy. It does not give particles more energy and it does not make them collide more often — it lowers the bar, so that a far larger share of the collisions already happening are now energetic enough to react. On a reaction profile the reactants and products sit at exactly the same energies as before; only the hump between them shrinks. Because the catalyst is unchanged at the end, it never appears in the chemical equation.
The catalytic converter under a petrol car is the everyday case: a honeycomb coated with platinum, palladium and rhodium turns carbon monoxide and unburnt fuel into carbon dioxide and water in the fraction of a second the exhaust gas takes to pass through — and the same few grams of metal are still there a decade later. Enzymes are biological catalysts: protein molecules that do the same job in living cells. Yeast's enzymes catalyse fermentation, which is why one topic sits behind both bread and beer. Different reactions need different catalysts — there is no universal one — which is why catalyst discovery is a serious industry.
DataRequired practical 11 — two ways to clock the same idea
The practical asks how concentration affects rate, and you should know both standard versions. Version 1 (gas volume): magnesium ribbon reacts with dilute hydrochloric acid; you record the volume of hydrogen in a gas syringe every 10 seconds, then repeat the whole run with each acid concentration. The independent variable is the acid concentration; the dependent variable is the volume of gas collected over time; the control variables are temperature, the mass and length of magnesium, and the total volume of acid. Version 2 (turbidity): sodium thiosulfate solution reacts with hydrochloric acid, producing a fine yellow precipitate of sulfur that gradually clouds the mixture. The flask sits on a printed black cross, and you time how long the cross takes to disappear from view; because a faster reaction gives a shorter time, the rate proxy is 1 ÷ time.
The errors are where the exam marks live. In the gas method, gas escapes in the second before the bung is seated — a systematic loss, reduced by having the syringe connected and starting the timer at the moment of mixing. In the cross method, the endpoint is a human judgement: the same person should watch the same cross through the same depth of solution every run. The reaction is slightly exothermic and room temperature drifts, so a water bath tightens the control. And the mixture releases sulfur dioxide, so the room must be well ventilated and solutions kept dilute. Repeat each concentration and use the mean — anomalies in timing experiments are routine, not rare.
A class runs the disappearing-cross experiment at five thiosulfate concentrations. At 8 g/dm³ the cross vanishes in 132 s; at 16 g/dm³, 66 s; at 24 g/dm³, 44 s; at 32 g/dm³, 33 s; at 40 g/dm³, 26 s. Convert to the rate proxy 1 ÷ time: 0.0076, 0.015, 0.023, 0.030 and 0.038 s⁻¹. Notice the pattern before you plot anything: doubling the concentration halves the time and doubles the rate. A graph of 1 ÷ time against concentration is a straight line through the origin, so rate is directly proportional to concentration — exactly what collision theory predicts, because doubling the particles per cm³ doubles the collision frequency.
ModelReversible reactions and the energy see-saw
Some reactions run both ways: the products can react to re-form the reactants, written \(\mathrm{A} + \mathrm{B} \rightleftharpoons \mathrm{C} + \mathrm{D}\). The classroom classic is copper(II) sulfate. Heat the blue hydrated crystals and they lose their water of crystallisation, leaving a white anhydrous powder; drip water back on and the blue returns instantly — and the tube gets noticeably hot. Ammonium chloride does the same trick with heat alone: warm it and it splits into ammonia and hydrogen chloride gases, which recombine into white ammonium chloride where the tube is cooler.
The energy bookkeeping is beautifully strict: if the forward reaction is exothermic, the reverse is endothermic, and the same amount of energy is transferred each way. Turning blue copper sulfate white absorbs exactly as much energy as turning white copper sulfate blue releases. Energy is conserved across the double arrow — a see-saw, not a leak.
ModelDynamic equilibrium — the busiest kind of standstill
Run a reversible reaction in a closed system — one where nothing can enter or leave — and it settles into equilibrium: the forward and reverse reactions both keep running, but at equal rates, so the amounts of every substance stop changing. Nothing has stopped. The concentrations are constant because the two rates cancel, not because the chemistry has finished — which is why the word dynamic matters.
Be precise about what is equal at equilibrium: the rates, not the amounts. The position of equilibrium can sit far to one side, with the flask overwhelmingly full of products and a whisper of reactants, and it is still a perfectly good equilibrium. The sealed cola bottle from the intro is the everyday model: a closed system with dissolved and gaseous CO₂ swapping places at matched rates. The moment you open it you have an open system — gas escapes, the balance can never be restored, and the drink chases an equilibrium it will only reach when it is flat.
ModelLe Chatelier's Principle (Higher tier) — the system fights back
Change any condition of a system at equilibrium and the position of equilibrium shifts to counteract the change. That single sentence — Le Chatelier's Principle — makes three separate predictions, one per dial.
Concentration: add more of a reactant and the equilibrium shifts right, using it up and making more product; remove a product as it forms and the system keeps shifting right to replace it — the trick industrial chemists use to squeeze yield out of a reluctant reaction. Temperature: raise it and the equilibrium shifts in the endothermic direction, absorbing the added heat; cool it and the exothermic direction is favoured. Pressure (gases only): raise it and the equilibrium shifts towards the side with fewer gas molecules, relieving the squeeze — so your first move in any pressure question is to count the moles of gas on each side of the equation and write the numbers down.
The nitrogen dioxide demo makes all three dials visible. \(2\mathrm{NO_2} \rightleftharpoons \mathrm{N_2O_4}\) — NO₂ is brown, N₂O₄ is colourless, and the forward reaction is exothermic. Count the gas molecules: two on the left, one on the right. Plunge a sealed tube into ice water: cooling favours the exothermic (forward) direction, so the mixture visibly pales. Into hot water: the equilibrium shifts backward in the endothermic direction and the brown deepens. Squeeze the mixture in a sealed gas syringe: the colour first darkens (everything is compressed) and then fades, because the higher pressure pushes the equilibrium towards the single-molecule side. Three correct predictions from one principle — and in each written answer, the mark sits on the reason: 'shifts to the side with fewer gas molecules', 'shifts in the endothermic direction to absorb the added heat'.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
C6 is examined on Chemistry Paper 2 (8464/C/2F or 2H — 1 hour 15 minutes, 70 marks), and it leans harder on AO2 and AO3 than almost any other chemistry section: graph work, practical design and data analysis, with roughly a fifth of chemistry-paper marks resting on maths skills. Expect to calculate, not just recall.
On rate graphs, use a ruler for tangents, leave your construction lines visible, and give the gradient as change in y ÷ change in x with units stated — an unlabelled 0.004 drops the unit mark. Whenever a question says 'explain using collision theory', the mark scheme wants the frequency of collisions and, for temperature, the proportion of particles exceeding the activation energy. Writing 'more collisions' instead of 'more frequent collisions' is the classic half-answer.
For required practical 11, be ready to name the independent, dependent and control variables in one line each, and to state one specific error with its fix: gas lost before the bung is seated, or the subjective disappearing-cross endpoint judged by different observers. On Higher tier, every Le Chatelier answer needs a direction AND a reason tied to the change: quote the counted gas molecules for pressure, name the endothermic or exothermic direction for temperature. Direction without reason is half marks; reason without direction is none.