- What is dynamic equilibrium
- Reaction mechanisms · coming soon
- Catalysis by transition metals · coming soon
A powder that is never used up
A bottle of hydrogen peroxide solution from a pharmacy can sit in a cupboard for a year and still work. It is decomposing all the time, into water and oxygen, but so slowly that you would never notice:
Now add a pinch of black manganese(IV) oxide powder, MnO₂. The liquid fizzes at once. A glowing splint held above it bursts into flame: the gas is oxygen. In a few minutes the reaction is over.
Here is the strange part. Filter off the black powder, dry it and weigh it. It has the same mass as before, and it looks the same. Add it to fresh hydrogen peroxide and it works again, as well as the first time.
So something made the reaction very much faster, and came out unchanged. How can a substance speed up a reaction without being used up? And if it is not used up, how can it be doing anything at all?
What a catalyst is
A catalyst is a substance that increases the rate of a reaction and is chemically unchanged at the end.
It works by providing a different route for the reaction, with a lower activation energy.
Read the first sentence carefully. "Chemically unchanged at the end" does not mean "takes no part". A catalyst does take part. It reacts during the reaction, and then it is made again before the reaction finishes. Because it is made again, a small amount can be used over and over. That is why the MnO₂ weighed the same.
The Maxwell–Boltzmann lecture ended with a promise: "The next lecture explains how a catalyst provides its lower route." Here is how.
The lower route
MnO₂ works on its surface, and the details there are hard to see. A dissolved catalyst shows the idea more clearly. Iodide ions, I⁻, also speed up the decomposition of hydrogen peroxide. The reaction is thought to go in two steps:
Follow the iodide ion. In step 1 it is used: it takes an oxygen atom from a peroxide molecule and becomes IO⁻. In step 2 the IO⁻ gives that oxygen atom to a second peroxide molecule, and the iodide ion comes back. Add the two steps, and I⁻ and IO⁻ cancel. The overall equation is exactly the uncatalysed one.
So the catalyst is used and remade in every cycle. The IO⁻ ion has a different role. It is made in one step and used up in the next, so it never appears in the overall equation. A substance like this is called an intermediate. A catalyst is there at the start and at the end; an intermediate exists only in the middle.
Why is this route faster? Because each of its steps has a lower energy barrier than the one big step of the uncatalysed reaction. The energy profile shows this.
Without a catalyst, the activation energy is about 75 kJ/mol. With iodide ions, the highest barrier on the new route is about 56 kJ/mol. Look at what has not changed. The reactants start at the same level, and the products end at the same level. So the enthalpy change, , is the same on both routes. The catalyst changes the road, not the start or the finish.
Now connect this to the last lecture. On the Maxwell–Boltzmann distribution, a catalyst moves the line to the left and leaves the curve alone. A lower barrier means a larger fraction of particles have enough energy. Using the Pre-University formula (the Arrhenius equation), lowering from 75 to 56 kJ/mol at 25 °C makes that fraction about 2000 times larger. A drop of 19 kJ/mol, which looks small, makes the reaction about two thousand times faster.
That is it. That is the whole idea. A catalyst does not give particles more energy. It gives them an easier road, so that more of the particles that already exist have enough energy for it.
Heating also speeds up a reaction. How hot would hydrogen peroxide need to be, with no catalyst, to decompose as fast as it does with iodide at 25 °C?
Show the answer
The energy fraction depends on divided by the temperature in kelvin. To match, , so K, about 124 °C. (This ignores the small extra effect of faster collisions.) Water boils long before that. This is why living things depend on catalysts. Your body cannot heat itself to 124 °C to speed up a reaction, so it lowers the barrier instead. The enzyme catalase lowers the barrier for this very reaction much further than iodide does; measured values vary, from under 10 to about 20 kJ/mol.
What a catalyst does not change
Students often give a catalyst more power than it has. Test yourself first.
Two flasks each contain the same volume of the same hydrogen peroxide solution. MnO₂ is added to flask A only. Both flasks are left until the reaction is completely over. Which flask gives more oxygen in total?
The answer: both give the same total volume of oxygen. The catalyst makes flask A reach that volume much sooner. The amount of oxygen is fixed by the amount of hydrogen peroxide, through the balanced equation.
The same reasoning gives the whole list. A catalyst does not change the reactants or the products, so it does not change or the amount of product. It does not change either, so it cannot make a reaction that is not feasible happen. The Gibbs lecture showed that "feasible" and "fast" are different questions; a catalyst answers only the second one.
Many reactions can also go backwards. A catalyst lowers the barrier for the reverse reaction by exactly the same amount, because the reverse route climbs the same hill from the other side. So it speeds up the forward and reverse reactions by the same factor. The equilibrium lectures show what this means: a catalyst helps a reaction reach equilibrium sooner, but does not change where the equilibrium lies.
Put each quantity in a box. Ask: is this about the route, or about the start and the finish?
What a catalyst changes, and what it does not.
- Changes: the rate of reaction · the route (mechanism) · the activation energy · the fraction of particles with enough energy · the time taken to reach equilibrium.
- Does not change: ΔH · ΔG (whether the reaction is feasible) · the total amount of product · the Maxwell–Boltzmann curve · the position of equilibrium.
Two kinds of catalyst
Catalysts are sorted by their state compared with the reactants.
A homogeneous catalyst is in the same state as the reactants. Iodide ions in hydrogen peroxide solution are an example: both are dissolved in water. Homogeneous catalysts usually work as the iodide does, by forming an intermediate and then being remade.
A heterogeneous catalyst is in a different state from the reactants, almost always a solid with gases or liquids reacting on it. MnO₂ in hydrogen peroxide is an example. So is the iron in the Haber process, which makes ammonia from nitrogen and hydrogen. A heterogeneous catalyst works on its surface, in four steps.
The key step is adsorption. When a molecule sticks to the surface, it forms weak bonds with the surface atoms, and its own bonds become weaker. In the Haber process, this matters a great deal. The N≡N triple bond is one of the strongest bonds there is, with an average bond enthalpy of 946 kJ/mol. On the iron surface, it is weakened enough to break at about 400–500 °C.
Since the reaction happens only on the surface, a solid catalyst is made with as much surface as possible: a fine powder, a porous solid, or a thin coating spread over a support. The surface can also be blocked. In a car's catalytic converter, a thin layer of platinum, palladium and rhodium is spread over a ceramic honeycomb. Lead from leaded petrol coats this surface and stops it working, a process called poisoning. When catalytic converters were fitted to most new petrol cars in the United States from the 1975 model year, those cars had to use unleaded petrol.
Enzymes are the catalysts of living things. They are proteins, each with an active site shaped to fit particular molecules. Most work in solution in cells, so they are often counted as homogeneous; some books treat them as a third kind. Catalase, which decomposes hydrogen peroxide in your blood and cells, is one of the fastest enzymes known.
A catalyst that fed the world
On 2 July 1909, Fritz Haber showed chemists from the company BASF a laboratory apparatus that made ammonia from nitrogen and hydrogen. Carl Bosch then led the engineering work, and BASF's chemists searched for a cheaper catalyst than the rare metal Haber had used, finally choosing iron with small amounts of other oxides added. Commercial production began at Oppau, in Germany, on 9 September 1913.
Ammonia made this way is the starting point for most nitrogen fertilisers. In 1999 the scientist Vaclav Smil estimated that, without the Haber–Bosch process, almost two-fifths of the world's population would not be alive. Without its iron catalyst, the reaction would be far too slow at any temperature where ammonia survives.
Worked examples
Question: "Explain, using the Maxwell–Boltzmann distribution, how a catalyst increases the rate of a reaction."
- 1. The catalyst provides an alternative route with a lower activation energy.
- 2. The temperature is unchanged, so the Maxwell–Boltzmann distribution is unchanged.
- 3. The line moves to the left, so the area beyond it is larger: a greater fraction of particles have at least the activation energy.
- 4. So a greater fraction of collisions are successful, and the rate increases.
Check: the answer names the route, says the curve is unchanged, and links the larger area to more successful collisions. "The catalyst gives the particles more energy" would lose marks.
Iron ions catalyse the reaction between peroxodisulfate ions and iodide ions. One proposed route is:
Find the overall equation. Which species is the catalyst, and which is the intermediate? Is the catalysis homogeneous or heterogeneous?
Add the steps and cancel what appears on both sides. 2Fe²⁺ is on the left of step 1 and the right of step 2. 2Fe³⁺ is on the right of step 1 and the left of step 2. Both cancel:
Fe²⁺ is present at the start and remade at the end, so it is the catalyst. Fe³⁺ is made in step 1 and used in step 2, so it is the intermediate. Everything is dissolved, so the catalysis is homogeneous.
Check: charges balance in the overall equation: on the left, on the right. ✓ There is also a reason this helps: S₂O₈²⁻ and I⁻ are both negative and repel each other, while each step joins a negative ion with a positive one.
For an invented reaction, kJ/mol. The activation energy is 120 kJ/mol without a catalyst and 70 kJ/mol with one. Find the activation energy of the reverse reaction on each route.
The reverse reaction starts at the products and climbs to the same peak. The collision theory lecture gave the rule: reverse = forward − .
With the catalyst: 70 − (−40) = 110 kJ/mol
Both barriers fall by 50 kJ/mol. The catalyst lowers the forward and reverse barriers by the same amount.
Check: the reverse of an exothermic reaction is endothermic, so its barrier must be larger than the forward one on each route: 160 > 120 and 110 > 70. ✓
Three traps
A catalyst gives no energy and does not change the Maxwell–Boltzmann curve. It lowers the barrier, so that more of the particles that already have energy can react.
A catalyst takes part in the reaction and is remade. Do not confuse it with an intermediate, which is made in one step and used in a later one. Catalyst: there at the start and the end. Intermediate: only in the middle.
A catalyst gives the same products in the same amounts, with the same and the same equilibrium position. It only gets there faster. If an answer says a catalyst "increases the yield", it has confused rate with amount.
Work each one on paper first. Open a hint only when you are stuck.
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High in the atmosphere, chlorine atoms destroy ozone. One route is: Cl(g) + O₃(g) → ClO(g) + O₂(g), then ClO(g) + O(g) → Cl(g) + O₂(g). Write the overall equation, name the catalyst and the intermediate, and say whether the catalysis is homogeneous or heterogeneous.
Hint 1
Add the two steps and cancel anything that appears on both sides.
Hint 2
Which species is there at the start and remade at the end? Which is made and then used?
Full answer
Overall: O₃(g) + O(g) → 2O₂(g). Cl is the catalyst: used in step 1 and remade in step 2. ClO is the intermediate: made in step 1 and used in step 2. Everything is a gas, so the catalysis is homogeneous. Check: atoms balance (4 O on each side), and because Cl is remade, one chlorine atom can destroy many ozone molecules, which is why small amounts of chlorine do so much damage.
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Explain why a catalyst does not change the total volume of oxygen made from 50 cm³ of hydrogen peroxide solution, but does change the shape of the graph of volume against time.
Hint 1
What decides how much oxygen can form? What decides how fast it forms?
Hint 2
Think about the gradient at the start and the final height of the curve.
Full answer
The total volume is set by the amount of hydrogen peroxide and the equation (2 mol H₂O₂ gives 1 mol O₂), which the catalyst does not change. The catalyst provides a route with a lower activation energy, so the initial rate is greater: the curve starts much steeper and levels off sooner, at the same final volume. Check: the rate experiments lecture showed that the initial rate is the gradient of the tangent at the start, and the final volume is a stoichiometric amount.
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A student has 0.0050 mol of H₂O₂ in solution and adds MnO₂. What volume of oxygen forms at room temperature and pressure (24 dm³/mol)? What mass of MnO₂ is left if 0.50 g was added?
Hint 1
Use the mole ratio from 2H₂O₂ → 2H₂O + O₂.
Hint 2
Moles of O₂ = 0.0050 ÷ 2. Multiply by 24 000 cm³/mol.
Full answer
Moles of O₂ = 0.0050 ÷ 2 = 0.0025 mol. Volume = 0.0025 × 24 000 = 60 cm³. The MnO₂ is a catalyst, so 0.50 g remains. Check: the volume does not depend on how much MnO₂ was added; more MnO₂ (more surface) would only make the 60 cm³ arrive sooner.
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An invented reaction in the body has kJ/mol without a catalyst. An enzyme lowers it to 60 kJ/mol. At body temperature (310 K), the Pre-University formula says this makes the reaction about 110 000 times faster. Find the temperature the uncatalysed reaction would need to match it, and explain why the body uses enzymes instead. Then compare the size of this effect with the iodide catalyst for hydrogen peroxide.
Hint 1
The Stop-and-think used the idea that the energy fraction depends on divided by the temperature in kelvin.
Hint 2
Set and solve for T. Then convert to °C.
Full answer
To match without the enzyme: K, about 192 °C. The body cannot reach this: proteins lose their shape far below it, and water would boil. Lowering the barrier is the only way to make such reactions fast at 37 °C. Compare: iodide lowers the barrier by about 19 kJ/mol and gives a factor of about 2000; this enzyme lowers it by 30 kJ/mol and gives about 110 000. Each extra few kJ/mol multiplies the effect. Check: 465 ÷ 310 = 1.5 = 90 ÷ 60. ✓
Every way of speeding up a reaction in this topic works through the same sentence from collision theory: more collisions, more of them with enough energy, more of them the right way round. Concentration and surface area change the first part. Temperature changes the first two. A catalyst is the only common way to change the barrier itself.
That is why catalysts sit behind most of modern chemistry and all of biology. Industry uses them to run reactions at lower temperatures, which saves energy and often gives fewer unwanted products. Living things run thousands of reactions at about 37 °C, each with its own enzyme. And because a catalyst does not change where equilibrium lies, the next topic, equilibrium, asks the question this lecture leaves open: when a reaction can go both ways, where does it stop?
- A catalyst increases the rate of a reaction by providing a different route with a lower activation energy, and it is chemically unchanged at the end.
- A catalyst takes part in the reaction and is remade, while an intermediate is made in one step and used up in a later one.
- A catalyst does not change the Maxwell–Boltzmann curve; it moves the Eₐ line to the left, so a larger fraction of particles have enough energy.
- A catalyst does not change ΔH, ΔG, the amount of product or the position of equilibrium, because these depend only on the start and the finish.
- Homogeneous catalysts are in the same state as the reactants; heterogeneous catalysts are usually solids that work on their surface by adsorption, reaction and desorption.