HookThe reaction that feeds half the planet — and the compromise behind it
In July 1909, in a laboratory in Karlsruhe, Fritz Haber watched a few cubic centimetres of liquid ammonia drip from his apparatus — the first time anyone had made ammonia from nothing but the nitrogen in the air and hydrogen, forced together under crushing pressure over a catalyst. Four years later the company BASF opened a full industrial plant at Oppau. The scientist and historian Vaclav Smil has estimated that roughly half of the nitrogen atoms in your body today passed through a Haber process reactor, because the fertiliser it makes feeds close to half of humanity. It has a fair claim to being the most consequential reaction of the twentieth century.
And it is built on a compromise that is the whole of C6. Making ammonia is reversible and exothermic, so a low temperature would give more of it — yet a low temperature makes the reaction unbearably slow. Haber's answer was to accept a middling temperature that gives a modest yield quickly, then recycle the leftovers. Rate (how fast a reaction goes) and equilibrium (how far it goes) pull in opposite directions, and industrial chemistry is the art of settling that argument. This section is how you measure a rate, why concentration, temperature, surface area and catalysts change it, and how Le Chatelier's Principle predicts which way a reversible reaction will shift.
ModelMeasuring a rate — mean rates and tangents
The rate of reaction is how fast a reactant is used up or a product is made, measured per unit of time. You track it by following something that changes as the reaction runs: the mass lost as a gas escapes on a balance, the volume of gas collected, or a colour or cloudiness developing.
The mean rate over a period is just the change divided by the time, \(\text{mean rate} = \dfrac{\text{change in quantity}}{\text{time taken}}\), with units such as cm³/s, g/s or mol/s. On a graph of product against time the line is steepest at the very start, where the reaction is fastest because the reactant is most concentrated, and it flattens to horizontal when the reaction stops. To find the rate at one particular instant you draw a tangent to the curve at that point and measure its gradient. On the Higher tier you can then convert a rate in cm³/s or g/s into mol/s using the molar gas volume or the relative formula mass.
Magnesium ribbon reacts with excess dilute hydrochloric acid, and 48 cm³ of hydrogen is collected before the reaction finishes at 40 s.
Mean rate over the whole reaction: \(\dfrac{48}{40} = 1.2\ \text{cm}^3\text{/s}\).
To find the rate at 10 s, draw a tangent to the curve there. Suppose the tangent rises 36 cm³ over a 20 s window: its gradient is \(\dfrac{36}{20} = 1.8\ \text{cm}^3\text{/s}\).
The instantaneous rate at 10 s (1.8 cm³/s) is higher than the whole-reaction mean (1.2 cm³/s), because early on the acid is at its most concentrated, so collisions are most frequent. Explaining why the tangent is steeper than the mean line is the analysis mark examiners are looking for.
MechanismCollision theory — every rate factor, one explanation
A reaction happens only when particles collide, and only when they collide with at least the activation energy and in a suitable orientation. The rate depends on how frequently such successful collisions happen, and every rate factor on the specification works through that one idea.
Raise the concentration of a solution, or the pressure of a gas, and the particles are packed closer together, so they collide more often and the reaction speeds up. Increase the surface area — a powder instead of a lump — and more particles are exposed to collide, which is why fine powders react so fast that flour mills and custard factories have suffered dust explosions. Raise the temperature and two things happen at once: the particles move faster, so they collide more often, and, more importantly, a larger fraction of them now exceed the activation energy, so more of the collisions succeed — which is why a rise of about 10°C roughly doubles the rate of many reactions. Add a catalyst and it offers a different route with a lower activation energy, so a greater fraction of collisions succeed, all without the catalyst being used up.
MechanismCatalysts and enzymes
A catalyst speeds up a reaction by providing a different pathway with a lower activation energy — an easier route over the hill — and it is not used up, so a tiny amount keeps working over and over. Crucially it changes only the rate, not the products and not the overall energy change: on a reaction profile a catalyst lowers the hump but leaves the reactants and products exactly where they were.
Different reactions need different catalysts — iron in the Haber process, platinum and rhodium in a car's catalytic converter — and industry values them because a lower activation energy means a reaction can run at a lower temperature, which saves energy and money. Enzymes are biological catalysts: large protein molecules that catalyse the reactions inside living cells, and also do useful work in brewing, baking and the biological washing powders that break down stains at low temperatures. Each enzyme is highly specific to the particular reaction it speeds up.
CaseThe disappearing cross and the gas syringe (Required practical 5)
Required practical 5 measures how concentration affects rate, in two classic ways. The first follows turbidity — cloudiness. Sodium thiosulfate reacts with hydrochloric acid to make a fine pale precipitate of sulfur, \(\text{Na}_2\text{S}_2\text{O}_3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{S} + \text{SO}_2 + \text{H}_2\text{O}\). You stand the flask on a paper cross and time how long the cross takes to disappear as the solution clouds over; a higher thiosulfate concentration clouds faster, so the time is shorter.
The second follows gas volume: magnesium reacts with hydrochloric acid, and you collect the hydrogen in a gas syringe, recording the volume every 10 s. A steeper graph means a faster reaction. In both methods the quantity you plot against concentration is \(\dfrac{1}{\text{time}}\), because a shorter time means a faster reaction — rate is inversely proportional to the time taken. Sensible controls are the same temperature, the same total volume and, for the cross, the same person judging when it has vanished — that end-point is subjective, and saying so is a genuine evaluation point.
ModelReversible reactions and dynamic equilibrium
Some reactions run both ways: the products can react to reform the reactants, shown by the double arrow. Heating hydrated copper sulfate drives off its water and turns the blue crystals white, an endothermic change, and adding the water back turns them blue again and releases heat, an exothermic change: \(\text{CuSO}_4 \cdot 5\text{H}_2\text{O} \rightleftharpoons \text{CuSO}_4 + 5\text{H}_2\text{O}\). Whatever the energy change is one way, it is exactly equal and opposite the other — if the forward reaction is exothermic, the reverse is endothermic by the very same amount.
In a closed system, where nothing can escape, a reversible reaction reaches equilibrium: the forward and backward reactions are going at the same rate, so the amounts of reactants and products stay constant. It is called a dynamic equilibrium precisely because both reactions are still happening — they are balanced, not stopped. Nothing looks like it is changing, but at the level of the particles both directions are running flat out.
MechanismLe Chatelier's Principle — predicting the shift (Higher tier)
Le Chatelier's Principle is the tool for predicting equilibria: if you change a condition on a system at equilibrium, the equilibrium shifts in the direction that opposes — partly cancels — the change you made.
Change the concentration: add more of a reactant and the equilibrium shifts forward to use it up, making more product; remove a product and it also shifts forward. Change the temperature: raise it and the equilibrium shifts in the endothermic direction to absorb the extra heat; lower it and it shifts in the exothermic direction — so heating a reaction whose forward direction is exothermic actually reduces its yield. Change the pressure (for gases only): raise it and the equilibrium shifts to the side with fewer gas molecules to relieve the pressure; lower it and it shifts to the side with more.
The Haber process is \(\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3\), with the forward reaction exothermic and 4 moles of gas on the left but only 2 on the right.
Le Chatelier predicts the best yield of ammonia from high pressure (a shift towards the 2-mole side) and low temperature (a shift in the exothermic forward direction).
But a low temperature makes the rate hopelessly slow. So industry settles on a compromise: about 450°C — sacrificing some yield to gain a workable rate — a pressure of about 200 atmospheres — as high as is safe and affordable — and an iron catalyst, which speeds both directions equally and so reaches equilibrium faster without changing its position. The unreacted nitrogen and hydrogen are recycled, so the roughly 15% converted on each pass is not wasted. That trade — theory says cold, reality says warm enough to be fast — is the compromise from the introduction, and the single most examined idea in C6.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
For rate calculations, always divide a change in quantity by a time and attach the unit; for an instantaneous rate, draw a tangent and take its gradient; and be ready to explain a graph's shape through collision frequency — steepest at the start because the reactant is most concentrated, flattening as it runs out.
When a question asks why a factor changes the rate, route every answer through collision theory: more frequent successful collisions. For temperature, add the second, larger effect — more particles now exceed the activation energy — because that is the point examiners most want to see. In Required practical 5, remember that the measured rate is one over the time, name your control variables, and admit that the disappearing-cross end-point is subjective.
On Higher-tier equilibrium questions, use a three-part structure: state the change, say 'the equilibrium shifts to oppose it', then name the direction and its effect on the yield — and for pressure, count the gas molecules on each side first. Above all, keep rate and yield separate: examiners deliberately set up situations, like the Haber process, where the condition that is fast and the condition that gives the best yield disagree, and the marks go to the candidate who explains the compromise rather than picking one side.