Rates of Reactions¶

Syllabus mapping

CH-11-03 — Chemical Reactions. Covers: the roles of activation energy, molecular orientation and the frequency of successful collisions in collision theory; using collision theory to analyse how concentration, temperature, surface area and catalysts affect reaction rate; a laboratory experiment measuring the effect of one of these factors.

Wording above is paraphrased for teaching use — check the official syllabus before treating any line as verbatim NESA text.

You need to know: the three things collision theory says must happen for particles to react, and how to use collision theory to explain why concentration, temperature, surface area and catalysts change the rate of a reaction.

Assumed knowledge check¶

Before this page makes sense, you should be comfortable with:

The particle model — particles in liquids and gases are always moving, faster when hotter
The four factors that affect rate from Year 10
Concentration
Reading a line graph and comparing how steep two curves are — see Trends and Relationships

Core¶

What is reaction rate?¶

The rate of a reaction is how fast reactants are used up, or products are made. Some reactions are almost instant (an explosion, a precipitate forming); some take years (iron rusting).

You can't see particles, so you measure something that changes as the reaction goes:

Measure	Good for
Volume of gas produced over time (gas syringe)	Reactions that give off a gas, e.g. metal + acid
Mass lost over time (open flask on a balance)	Reactions that give off a gas
Time for a colour change, or for a cloudy precipitate to hide a cross under the flask	Reactions with a visible change

On a graph of product against time, the rate is shown by how steep the curve is:

It's steepest at the start, when there are the most reactant particles.
It gets flatter as reactants get used up.
It goes flat when the reaction stops — one reactant has run out.

🖼️ Diagram needed — volume of gas vs time graph for one reaction: steep at the start, curving over, then flat. Label "fastest here", "slowing down", "reaction finished". A second, steeper curve reaching the same final height, labelled "faster reaction, same amount of product". TODO — placeholder until sourced/drawn.

Collision theory¶

For particles to react, three things must happen:

They must collide.
They must collide with enough energy — at least the activation energy (\(E_a\)), the minimum energy needed to start breaking bonds.
They must collide with the right orientation — the right parts of each particle have to hit each other.

A collision that does all three is a successful collision. Most collisions aren't successful — the particles just bounce off.

The rate of a reaction depends on the number of successful collisions per second.

Anything that increases the number of successful collisions per second speeds up the reaction. Every factor below is explained this way.

Factor 1: Concentration (and pressure)¶

Higher concentration → faster reaction.

Why: more reactant particles in the same volume means they collide more often. More collisions per second → more successful collisions per second.

For gases, increasing the pressure does the same thing — it squashes the particles closer together.

Factor 2: Temperature¶

Higher temperature → faster reaction. As a rough guide, a 10 °C rise often about doubles the rate.

Why — two reasons:

Particles move faster, so they collide more often.
More importantly, a bigger fraction of the collisions have at least the activation energy, so more of them are successful.

Reason 2 is the bigger effect — and it's the one examiners look for.

Factor 3: Surface area¶

Smaller pieces (larger surface area) → faster reaction.

Why: only particles on the surface of a solid can collide with the other reactant. Cutting a solid into smaller pieces exposes more particles on the surface, so there are more collisions per second.

That's why powdered marble reacts with acid much faster than marble chips, why kindling catches fire faster than a log, and why chewing food helps you digest it.

Factor 4: Catalysts¶

A catalyst is a substance that speeds up a reaction without being used up. It's still there, unchanged, at the end.

Why: a catalyst provides a different pathway for the reaction with a lower activation energy. At the same temperature, many more collisions now have enough energy to succeed.

Examples:

Manganese dioxide makes hydrogen peroxide break down rapidly into water and oxygen.
Enzymes are biological catalysts — the catalase in potato or liver does the same job.
Catalytic converters in cars speed up reactions that turn toxic exhaust gases into less harmful ones.

🖼️ Diagram needed — energy profile diagram with two humps over the same reactants and products: a tall hump labelled "without catalyst" and a lower hump labelled "with catalyst", activation energy arrows for each. Same style as the energy profile diagrams on Energy Changes. TODO — placeholder until sourced/drawn.

Summary¶

Factor	Change	Effect on rate	Collision theory reason
Concentration / pressure	Increase	Faster	More frequent collisions
Temperature	Increase	Faster	More frequent collisions and more with energy ≥ \(E_a\)
Surface area	Increase (smaller pieces)	Faster	More particles exposed, so more frequent collisions
Catalyst	Add one	Faster	Alternative pathway with lower \(E_a\), so more collisions are successful

Worked example — explaining a result

Two identical masses of zinc react with excess hydrochloric acid. Experiment A uses \(1.0\ \text{mol L}^{-1}\) acid; Experiment B uses \(2.0\ \text{mol L}^{-1}\) acid. Experiment B produces gas faster, but both produce the same total volume of hydrogen. Explain both observations.

Faster: B's acid is more concentrated, so there are more acid particles in the same volume. They collide with the zinc more often, so there are more successful collisions per second.

Same total volume: the amount of hydrogen depends on how much zinc there is (the zinc is the limiting reagent — the acid is in excess). Both experiments used the same mass of zinc, so they make the same amount of hydrogen. Changing the rate changes how fast you get the product, not how much.

Practical¶

Suggested prac: measuring the effect of one factor on rate. Two classic options:

Sodium thiosulfate and hydrochloric acid ("disappearing cross"). The reaction makes a cloudy yellow precipitate of sulfur. Put the flask on a paper marked with a cross, add the acid, and time how long it takes until you can't see the cross. Repeat at different temperatures (warm the thiosulfate solution first) or different concentrations. Shorter time = faster rate; you can use \(\dfrac{1}{\text{time}}\) as a measure of rate.

Marble chips and hydrochloric acid. Measure the mass lost (or gas volume) every 30 s, using the same mass of marble as large chips, small chips and powder (surface area) — or with different acid concentrations.

Variable	In the thiosulfate temperature experiment
Independent	Temperature
Dependent	Time for the cross to disappear
Controlled	Volumes and concentrations of both solutions; the same cross; the same person judging

Safety: the thiosulfate reaction gives off a small amount of sulfur dioxide gas — work in a well-ventilated room and don't heat the solutions above about 50 °C. See Planning for variables and risk assessments.

Advanced

Why temperature has such a big effect: the energy distribution. In any sample, particles have a range of energies — a few slow, most medium, a few very fast. Only the particles with energy above \(E_a\) can react. Raising the temperature shifts the whole spread towards higher energies, so the number of particles above \(E_a\) increases a lot — far more than the small increase in collision frequency. A catalyst works on the same graph from the other direction: it moves the \(E_a\) line lower, so more of the existing particles are above it.

🖼️ Diagram needed — Maxwell–Boltzmann distribution: number of particles vs kinetic energy, curves at a lower and a higher temperature, with the activation energy marked and the area beyond it shaded for each. TODO — placeholder until sourced/drawn.

Calculating rate from a graph. The average rate over a period is the change in amount ÷ time taken (e.g. \(\dfrac{40\ \text{mL}}{20\ \text{s}} = 2.0\ \text{mL s}^{-1}\)). The rate at one instant is the gradient of a tangent to the curve at that point — the same skill as finding velocity from a displacement–time graph.

What changes the final amount of product, and what doesn't. Temperature, surface area, catalysts and (excess-reagent) concentration change how fast the curve rises. Only changing the amount of the limiting reagent changes how high it finishes. Exam graphs often test exactly this.

Extension

Beyond the syllabus for this course — won't appear in the HSC, included for interest / depth study inspiration.

Dust explosions. Flour, sugar, coal dust and even custard powder can explode if a cloud of fine powder meets a spark. Each grain is tiny, so the total surface area exposed to oxygen is enormous, and the combustion happens almost instantly. Flour mills and grain silos have strict rules about dust for exactly this reason — it's the surface area factor at its most extreme.

Enzymes have a "best" temperature. Unlike most catalysts, enzymes are proteins with a precise 3D shape. Warm them up and the rate increases — until, beyond about 40 °C for most human enzymes, the shape breaks down (denatures) and the rate falls sharply.

📎 Depth study idea

Investigate how temperature affects the rate at which the enzyme catalase (from potato or liver) breaks down hydrogen peroxide, measuring the volume of oxygen produced, and find the optimum temperature. Compare it with a non-biological catalyst such as manganese dioxide over the same temperature range.

Video/visual resources¶

🖥️ PhET Simulation — Reactions & Rates — change temperature and concentration and see how many collisions succeed (older Java-based sim — runs in the browser but can be slow to load)
🎥 TODO: source a short video of the disappearing-cross experiment at two temperatures side by side.
🎥 Khan Academy — TODO: source a collision theory explainer.

Check yourself¶

State the three conditions for a successful collision.

Answer

The particles must (1) collide, (2) with at least the activation energy, and (3) with the correct orientation.
Explain, using collision theory, why increasing the concentration of an acid increases the rate at which it reacts with magnesium.

Answer

A higher concentration means more acid particles in the same volume, so they collide with the magnesium more often. More collisions per second means more successful collisions per second, so the reaction is faster.
Give the two reasons a higher temperature increases reaction rate. Which is more important?

Answer

(1) Particles move faster, so they collide more often. (2) A larger fraction of collisions have at least the activation energy, so more are successful. The second reason is more important.
Why does powdered calcium carbonate react with acid faster than the same mass of large chips?

Answer

The powder has a much larger surface area, so more particles are exposed to the acid. That means more collisions per second, and so more successful collisions per second.
What is a catalyst, and how does it speed up a reaction?

Answer

A catalyst is a substance that speeds up a reaction without being used up. It provides an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions have enough energy to be successful.
A reaction is repeated with a catalyst added. Will the catalyst change the total amount of product formed? Explain.

Answer

No. A catalyst only makes the reaction happen faster. The total amount of product depends on the amount of the limiting reagent, which hasn't changed.