3bRates of reaction

Five factors, and all five make a reaction faster.

Knowing the direction earns little on its own. The marks are in explaining why, which is the next note.

Pressure only counts for gases

If a reaction is between a solid and a solution, changing the pressure of the air above it does nothing.

And check whether the gas is a reactant. A reaction that produces a gas is not affected by pressure in the way a reaction between gases is.

More is not the same as more concentrated

Using a larger volume of the same acid gives more product in total, but does not change the rate. Rate depends on how crowded the particles are, not how many there are altogether.

That distinction turns up in questions regularly.

Rate and yield are different things

This is worth fixing early, because it runs through the whole topic.

Two flasks with identical amounts of reactant, one at 20 °C and one at 40 °C, produce the same total volume of gas. The hot one just gets there sooner.

• Rate — how fast the product appears.
• Yield — how much product forms in total, fixed by how much reactant there was.

Temperature, surface area and catalysts all change the rate without changing the yield — the same amount of product forms, just sooner.

Concentration is the exception worth care. Raising it speeds the reaction up, and if that reactant is the limiting one you have also added more of it, so more product forms. Rate and yield both change. If the reactant is in excess, only the rate changes.

So the safe statement is: for a fixed amount of reactants, these factors change how fast, not how much.

Reactions happen when particles collide with enough energy. Every rate explanation comes back to one of those two things: how often particles collide, or how energetic the collisions are.

Activation energy

The activation energy is the minimum energy a collision needs for a reaction to occur.

That is why colliding is not enough on its own. Many collisions are too gentle, and the particles simply bounce apart.

Match each factor to its mechanism

Getting the right mechanism for the right factor is what separates a full answer from a partial one.

Concentration and pressure

Both put more particles into the same volume, so the particles are closer together and collide more often.

They do not make collisions more energetic — only temperature does that.

Surface area

Only particles on the surface of a solid can react. Breaking a lump into powder exposes far more of them, so collisions with the liquid or gas happen more often.

This is why flour dust can explode while a bag of flour cannot. Enormous surface area makes the reaction fast enough to release its energy almost instantly.

Temperature does two things

This is the only factor with two effects, and full marks usually need both:

1. Particles gain energy and move faster, so they collide more often.
2. A greater proportion of collisions exceed the activation energy, so more of them succeed.

The second effect is the larger one, and it is why a modest temperature rise can change the rate so sharply. An answer giving only the frequency effect is incomplete.

Catalysts

A catalyst increases the rate without being chemically changed.

It works by providing an alternative pathway with a lower activation energy, so a greater proportion of the existing collisions have enough energy to react.

It does not give the particles more energy. Examiners report this confusion specifically, and it is worth saying the correct version aloud: the barrier drops, the particles are unchanged.

Because it is not used up, a small amount works indefinitely and can be recovered and reused — which is what makes catalysts economic in industry, alongside letting a plant run at a lower temperature.

Two ways to follow a reaction that makes a gas

1. Collect the gas in a syringe and record the volume at regular intervals.
2. Measure the loss in mass on a balance as the gas escapes.

The mass falls because carbon dioxide leaves the flask. Mass is not destroyed — it leaves the apparatus, which is a different statement and the one that earns the mark.

Reading a volume-time graph

The gradient is the rate. Steeper means faster.

A typical graph is steepest at the start and gradually levels off:

• Steepest at the start, because the reactants are at their highest concentration.
• Getting shallower, because the reactants are being used up, so collisions become less frequent.
• Horizontal at the end, because the reaction has stopped — a reactant has run out.

A horizontal line means the rate is zero, not that it is steady.

The disappearing cross method

A cross drawn under a flask disappears as a precipitate clouds the mixture. A shorter time means a faster reaction.

Its weakness is that deciding exactly when the cross has vanished is a judgement, so different people record different times. Using a light sensor and data logger removes that subjectivity — "be more careful" does not.

Investigating surface area

Use the same mass of marble each time, with different chip sizes. Using the same number of chips changes two variables at once.

Keep the acid volume, acid concentration and temperature constant.

The powder and the large chips give the same final volume of gas, because the same mass of marble reacts. Only the time differs.

Investigating catalysts

To compare solids as catalysts for hydrogen peroxide decomposition, keep the volume and concentration of peroxide and the mass of solid the same, and collect the oxygen produced.

2H₂O₂ → 2H₂O + O₂

Include a run with no solid added as a control. Without it, nothing shows the solids made any difference.

To prove a solid acted as a catalyst, filter, dry and reweigh it. An unchanged mass shows it was not used up.

Reaction profiles (Separate Chemistry)

The peak height above the reactants is the activation energy.

Adding a catalyst lowers the peak only. The reactant and product levels stay exactly where they were, which is the diagram's way of showing that a catalyst changes the rate but never ΔH.

Marble chips and dilute hydrochloric acid produce carbon dioxide, and the rate is followed by measuring the gas.

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

Two ways to follow it

Gas syringe — collect the carbon dioxide and record the volume every 30 seconds.

Mass loss — stand the flask on a balance and record the mass every 30 seconds. The mass falls as gas escapes. Plug the neck with cotton wool so acid spray cannot escape while gas can.

Investigating surface area

Use the same mass of marble in different sizes: large chips, small chips, powder. Keep the acid volume and concentration the same.

Powder reacts fastest. The same mass broken smaller has more surface area exposed, so more collisions happen each second.

Investigating concentration

Use the same chips and vary the acid concentration.

More concentrated acid reacts faster: more acid particles in the same volume means more frequent collisions.

Reading the graphs

Volume of gas against time gives a curve that is steepest at the start and levels off.

• Gradient = rate. Steeper means faster.
• It flattens because a reactant is being used up, so collisions become less frequent.
• Flat means the reaction has stopped.

Compare rates using the initial gradient, where the conditions are still the ones you set.

Does the total volume change?

This is the distinction examiners test most.

Surface area changes the rate, not the amount. Powder and chips of the same mass reach the same final volume, just at different times. Breaking a solid up does not create more of it.

Concentration depends on what is limiting. With the marble in excess and the acid limiting, doubling the concentration doubles the acid available, so the final volume also doubles. If the acid is in excess, the curves rise at different rates but finish at the same volume.

So "faster but the same total" is right for surface area, and right for concentration only when the varied reactant is in excess — read the quantities before deciding.

Fair test

Change one variable and hold the rest: mass and size of marble, volume and concentration of acid, and temperature. Temperature affects rate strongly, so an experiment run on a warm bench and a cold one is not comparing what it claims to.

Where marks are lost

• Saying powder produces more gas rather than producing it faster.
• Measuring the gradient at the flat end, where every rate looks the same.
• Forgetting the cotton wool, so lost acid spray is counted as lost gas.

Hydrogen peroxide decomposes slowly on its own. Different solids are tested to see which speed it up.

2H₂O₂ → 2H₂O + O₂

Apparatus

Conical flask with a bung and delivery tube, gas syringe (or measuring cylinder inverted over water), stopwatch, balance.

Eye protection throughout. Hydrogen peroxide is an irritant and an oxidiser — it bleaches skin and clothing and stings the eyes. Use dilute solution only.

The reaction with manganese(IV) oxide is fast and vigorous, so gas is produced suddenly in a stoppered flask. Make sure the delivery tube is clear before adding the catalyst, and never wedge the bung tight.

Method

1. Measure a fixed volume and concentration of hydrogen peroxide into the flask.
2. Add a fixed mass of the solid being tested.
3. Stopper immediately and start the stopwatch.
4. Record the volume of oxygen every 15 or 30 seconds.
5. Repeat with each solid, and with no solid at all as a control.

Comparing the solids

Plot all the runs on one set of axes and compare the initial gradients. The steepest curve is the best catalyst.

Manganese(IV) oxide is dramatically the most effective. Copper(II) oxide and lead(IV) oxide have some effect; sand has essentially none and behaves like the control.

The no-catalyst run is what makes the others mean anything — without it there is no baseline to say the solid did anything at all.

Testing the gas

A glowing splint relights: oxygen.

Showing the catalyst is unchanged

This is the part that turns the experiment into evidence for what a catalyst is.

1. Weigh the solid before adding it.
2. At the end, filter the mixture to recover it.
3. Wash and dry the residue.
4. Reweigh.

The mass is the same. The catalyst speeded the reaction up without being used up, which is why a small amount keeps working indefinitely.

A fuller demonstration adds fresh hydrogen peroxide to the recovered solid and shows it still works.

Fair test

Only the solid may change:

• Same volume and concentration of hydrogen peroxide
• Same mass of solid
• Same particle size — a powder has more surface area than lumps, and would appear to be the better catalyst on that ground alone
• Same temperature

Particle size is the one most often missed, and it is the one that most easily produces a false conclusion.

Where marks are lost

• Saying the catalyst is "used up slowly" — the recovered mass is the evidence that it is not.
• Omitting the control run.
• Comparing final volumes rather than initial rates. Every run ends at the same volume, because the catalyst changes the speed and not the amount of oxygen the peroxide can give.