Summary — What This Chapter Covers

• Collision Theory: reactions only happen when particles collide with enough energy (≥ activation energy) AND the correct orientation.
• Rates of Reaction: how to measure and calculate rate from graphs (mass loss, gas volume, tangents).
• Maxwell-Boltzmann Distributions: a graph showing the spread of particle energies, and how temperature reshapes it.
• Catalysts: speed up reactions by lowering activation energy via an alternative pathway, without being used up.
• Dynamic Equilibrium: in a closed system, reversible reactions reach a state where forward rate = backward rate.
• Le Chatelier's Principle: a system at equilibrium shifts to counteract any change made to it (concentration, pressure, temperature).
• Industrial Compromises: real factories (like the Haber Process) balance rate vs. yield vs. cost vs. safety.

1. Collision Theory

Why do some collisions "work" and others don't?

Imagine two people trying to shake hands while running past each other in a crowded street. If they're moving too slowly, or their hands miss each other completely, nothing happens — they just bump and carry on. But if they're moving fast enough and their hands meet just right, a proper handshake happens.

That's exactly what's going on at the molecular level. When reactant particles collide, most of these collisions are unsuccessful (ineffective) — the particles just bounce off each other without reacting. Only a small fraction are successful (effective) collisions, where a chemical reaction actually takes place.

For a collision to be effective, two conditions must both be met:

1. The particles must collide with the correct orientation (the reactive parts of the molecules need to actually meet).
2. The particles must collide with enough energy — at least equal to the activation energy (E_a) of the reaction.

Collision Frequency — what makes collisions happen more often?

Collision frequency is simply how many collisions happen per unit time. It can be increased by:

• Increasing concentration — more particles squeezed into the same volume means they bump into each other more often.
• Increasing pressure (for gases) — same number of particles in a smaller volume = more frequent collisions.
• Increasing temperature — particles move faster (more kinetic energy), so they collide more often and harder.
• Increasing surface area (smaller particle size) — more particles are exposed and available to react.

Activation Energy — the "hill" reactions must climb

Activation energy (E_a) is the minimum energy that colliding particles need in order to break the bonds necessary to start a reaction. Think of it as a hill: even if your ball (the reactants) is sitting at the top of a slope leading down to a lower energy valley (the products), it still needs a little push to get over the small bump in front of it before it can roll all the way down.

2. Rates of Reaction — Graphs & Calculations

As a reaction proceeds, reactant concentration falls and product concentration rises. The rate of reaction is how fast this change happens, with units mol dm⁻³ s⁻¹.

Method 1 — Measuring Mass Loss

If a gas escapes from the reaction vessel (e.g. CaCO₃ + HCl → CO₂), the total mass on a balance decreases over time. This mass loss is proportional to the amount of reactant used up, so plotting mass vs. time gives a curve you can analyse just like a concentration-vs-time graph.

Method 2 — Measuring Gas Volume

Alternatively, you can trap the gas produced (e.g. using a gas syringe, or by collecting it over water in an inverted measuring cylinder) and record its volume at regular time intervals. This gives a graph of "amount of product formed" against time — the classic rising curve that flattens off once the reaction finishes.

Finding Rate From a Graph — the Tangent Method

Since these graphs are curves (not straight lines), you can't just use "rise over run" directly — the rate keeps changing throughout the reaction. Instead, to find the rate at a specific moment in time, you draw a tangent (a straight line that just touches the curve at that one point) and calculate its gradient.

Method 3 — "Stopping the Clock" (Quenching)

Measuring concentration directly using titration is tricky, because sampling and titrating takes time — and the reaction keeps going while you do it, unless you deliberately stop it (this is called quenching). A simpler alternative is to time how long it takes to reach a visible endpoint — for example, timing until a magnesium ribbon fully dissolves, or the classic "disappearing cross" experiment with sodium thiosulfate and hydrochloric acid, where a precipitate of sulfur gradually obscures a cross viewed through the flask.

Limitation: this method only gives you one single data point per experiment (the total time taken), rather than a full curve of data throughout the reaction.

3. Maxwell-Boltzmann Distributions

Not every particle in a sample has the same energy. Picture a school full of students running a race — a few are exhausted and barely moving, a few are sprinting flat out, but most are running at some "medium" pace in between. The Maxwell-Boltzmann distribution is exactly this idea applied to particle energies: it's a graph showing how many particles have each possible amount of energy at a given temperature.

Key features of the curve you must be able to describe:

• The curve starts at the origin (0,0) — no particles have zero energy.
• It does not touch the x-axis at the high-energy end — the curve approaches it but never quite reaches zero, meaning a very small number of particles can have extremely high energy.
• The area under the whole curve represents the total number of particles in the sample.
• Only the small area to the right of E_a represents particles with enough energy to react on collision.

What happens when you increase the temperature?

Raising the temperature gives particles more kinetic energy on average. This has two visible effects on the curve:

1. The peak shifts to the right (higher average energy) and gets lower (flatter, more spread out).
2. The curve becomes broader — energies are spread across a wider range.

Crucially, the shaded area beyond E_a gets bigger at the higher temperature — meaning a greater proportion of particles now have enough energy to react. Combined with the fact that particles are also moving faster (more frequent collisions), this is why a small rise in temperature causes a surprisingly large increase in reaction rate.

4. Catalysts & Energy

A catalyst speeds up a chemical reaction without itself being used up in the process. It achieves this not by giving particles more energy, but by offering them an easier route to the same destination.

Three Big Advantages of Catalysts

1. Increased rate of reaction — more product made in a given time.
2. Reduced energy costs — the reaction can be run at lower temperature/pressure and still proceed quickly, saving energy and money.
3. Improved atom economy — the alternative pathway can reduce unwanted side-reactions and by-products (e.g. zeolites used in phenol synthesis).

Homogeneous vs. Heterogeneous Catalysts

On the Maxwell-Boltzmann curve, a catalyst doesn't move the curve at all (temperature hasn't changed) — instead it lowers the E_a threshold itself, which shifts more of the existing area under the curve into the "successful collision" zone.

5. Dynamic Equilibrium in Reversible Reactions

Some reactions go to completion — the reactants are entirely used up and the reaction just stops. But many reactions are reversible: the products can react with each other to reform the original reactants. We show this using two half-arrows (⇌) instead of one.

What Actually Is "Dynamic Equilibrium"?

This is one of the most commonly misunderstood ideas in the whole topic, so let's be really precise about it. Picture two identical escalators — one going up, one going down — right next to each other, both moving at exactly the same speed, with people constantly stepping on and off both. From a distance, if you just count the number of people on each escalator at any moment, that number stays constant. But that doesn't mean nobody is moving — people are still riding both escalators the whole time. That's dynamic equilibrium.

At dynamic equilibrium:

• The rate of the forward reaction equals the rate of the backward reaction.
• Both reactions are still happening constantly — nothing has "stopped."
• The concentrations of reactants and products remain constant (but they are almost never equal to each other).
• This can only happen in a closed system, where nothing escapes.

Closed System vs. Open System

Reactions entirely in solution can often reach equilibrium even in an "open" flask, because a negligible amount of liquid is lost to evaporation. But any reaction involving a gas escaping needs a genuinely sealed, closed container to reach equilibrium.

6. Le Chatelier's Principle

Think of equilibrium like a see-saw that's perfectly balanced. If you push down on one side, the see-saw doesn't just stay tilted forever — it shifts to try to restore some balance, though it won't return to exactly where it started. That's Le Chatelier's principle in a nutshell: the system "fights back" against whatever you do to it, though it never fully cancels the change out.

• Equilibrium shifting left → concentration of reactants increases.
• Equilibrium shifting right → concentration of products increases.

Effect of Concentration Changes

Effect of Pressure Changes (Gas Reactions Only)

This one only matters when reactants or products are gases — and it's all about counting the number of gas moles on each side of the equation.

Why? Because shifting toward the side with fewer gas molecules reduces the total number of particles bumping around, which lowers the pressure back down — directly opposing the increase you made. It's Le Chatelier's principle in action again.

Effect of Temperature Changes

This depends on which direction of the reaction is exothermic and which is endothermic. Every reversible reaction has one direction that releases heat and the opposite direction that absorbs it.

7. Industrial Compromises — The Haber Process

In a lab, you can just pick whatever conditions give the highest yield. But in a real factory, chemists have to balance rate (how fast product is made) against yield (how much product you get) against cost and safety. The Haber Process — making ammonia — is the textbook example of this compromise.

The Temperature Compromise

Because the forward reaction is exothermic, Le Chatelier's principle tells us that a lower temperature would give a higher yield of ammonia (equilibrium shifts right, toward the exothermic direction). But lower temperature also means a much slower rate of reaction — it could take forever to reach that high yield.

The industry compromise: 450 °C. This gives an acceptable yield (roughly 35%) within a reasonable time frame. Higher temperatures aren't used because yield drops too much and energy costs rise; lower temperatures aren't used because the rate becomes uneconomically slow.

The Pressure Compromise

There are 4 moles of gas on the reactant side (1 N₂ + 3 H₂) but only 2 moles on the product side (2 NH₃). So increasing pressure shifts the equilibrium right, increasing the yield of ammonia — great news for yield.

But very high pressures are extremely expensive to generate and maintain, and come with serious health and safety risks (thick, reinforced pipework and vessels are needed). In practice, around 200 atmospheres (20 MPa) is used — doubling it to 400 atmospheres only increases yield by about 7%, which isn't worth the massive extra cost and risk.

What to Memorise

Concepts Checklist

Tick each concept off once you can explain it out loud, without notes, in your own words.

• I can explain what makes a collision "successful" vs "unsuccessful."
• I can list the four ways to increase collision frequency.
• I can define activation energy and sketch exothermic/endothermic energy profiles.
• I can calculate rate of reaction from a tangent drawn on a graph.
• I know the difference between measuring rate by mass loss vs. gas volume, and their limitations.
• I can sketch a Maxwell-Boltzmann distribution curve and label E_a.
• I can explain how increasing temperature reshapes the Maxwell-Boltzmann curve.
• I can explain how a catalyst affects the Maxwell-Boltzmann curve and E_a.
• I know the difference between homogeneous and heterogeneous catalysts.
• I can define dynamic equilibrium precisely, without saying "concentrations are equal."
• I understand why equilibrium requires a closed system.
• I can state Le Chatelier's principle from memory.
• I can predict shifts in equilibrium position from changes in concentration.
• I can predict shifts in equilibrium position from changes in pressure (using mole counting).
• I can predict shifts in equilibrium position from changes in temperature (using ΔH).
• I can explain the temperature and pressure compromises used in the Haber Process.

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