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Module 3 · L11 of 12 35 min ⚡ +50 XP in Learn · +25 to complete

Collision Theory & Reaction Rate

Snap a glow stick and drop one in iced water and another in warm water: the warm one glows brightly but fades quickly, while the cold one glows dimly but lasts much longer. The chemistry is the same in both, and each stick starts with the same amount of reactant, so what changes is the reaction rate. Temperature is the variable doing the work here.

Today's hook, Snap two identical glow sticks and put one in iced water and one in warm water. The warm stick glows brightly but fades quickly; the cold stick glows dimly but lasts much longer. Same chemistry, same starting amount of reactant, only the rate differs, and temperature is the variable controlling it.
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Warm up and recall

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions
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Recall, your gut answer first
+5 XP warm-up

You’ve probably snapped a glow stick and seen it start glowing. If you put it in a cup of hot water, it glows intensely for a short time then fades quickly. If you put it in the freezer, it glows dimly for a very long time.

Safety, glow sticks and warm water. Keep glow sticks intact, never cut, puncture or microwave one, as the contents can irritate skin and eyes. To compare temperatures use warm (not boiling) tap water and iced water; do not use boiling water or a naked flame. If liquid leaks, avoid skin and eye contact and rinse with water.
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What you'll master, and the words for it

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What you'll master
Know

Key facts

  • The two conditions for an effective collision: sufficient energy (kinetic energy ≥ activation energy Eₐ) AND correct orientation of the colliding particles
  • The definition of activation energy (Eₐ), the minimum energy colliding particles must have for the reaction to occur
  • How to read and interpret a Maxwell-Boltzmann distribution diagram: peak position, curve shape, and the area to the right of Eₐ
Understand

Concepts

  • Why only a small fraction of collisions are effective, most particles do not have sufficient energy or correct orientation simultaneously
  • Why reaction rate decreases over time as reactant concentrations fall, fewer particles available → fewer collisions per second
  • How increasing temperature shifts the Maxwell-Boltzmann distribution to higher energies, increasing the proportion of particles that exceed Eₐ
Can do

Skills

  • Calculate average reaction rate from experimental data over a given time interval (Δ[quantity] / Δt)
  • Draw a labelled energy diagram identifying reactants, products, transition state, Eₐ (measured from reactant level to peak), and ΔH
  • Explain why a glow stick glows more brightly in hot water using collision theory language: kinetic energy → proportion exceeding Eₐ → effective collision frequency → rate
Beyond the syllabus. Collision theory and activation energy are Module 3 Core and both are examined. You need to be able to say that an effective collision requires energy at or above Eₐ and the correct orientation, and to explain a rate change in those terms. Reading a Maxwell-Boltzmann distribution as a named diagram, peak position, curve shape and the area beyond Eₐ, is supporting depth rather than a prescribed Year 11 endpoint. The idea the curve carries is not optional, though: a temperature answer only scores full marks if it says the proportion of particles with energy at or above Eₐ rises, not merely that collisions become more frequent. Learn that reasoning, and treat the curve as a picture of it rather than as something to memorise.
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Key terms
Collision theory
A model stating that reactions occur when particles collide with sufficient energy and correct orientation.
Activation energy (Ea)
The minimum energy required for a collision to result in a chemical reaction; must be exceeded for effective collisions.
Effective collision
A collision between particles that results in a reaction; requires sufficient energy (≥ Ea) and correct geometry.
Reaction rate
The change in concentration of a reactant or product per unit time; measured in mol L⁻¹ s⁻¹.
Maxwell-Boltzmann distribution
A graph showing the distribution of particle kinetic energies; area under the curve to the right of Ea represents the fraction of particles (or collisions) with at least the activation energy. Whether such a collision actually reacts also depends on correct orientation.
Transition state (activated complex)
The high-energy, unstable intermediate at the peak of the energy profile diagram; formed briefly during the collision.
Cross-lesson links: In L01–L10 you studied what reactions occur and which direction they proceed, precipitation, combustion, redox, and galvanic cells. L11–L12 shift the question: given that a reaction can occur, how fast? The collision theory model you build here is the particle-level foundation for every rate calculation and experimental design question in HSC Chemistry.
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Defining Reaction Rate

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Defining Reaction Rate
core concept

Drop marble chips into dilute hydrochloric acid, fizzing starts fast and slows as the reaction proceeds. Warm the acid first and the same marble chips fizz violently from the first second. Something is different: not the reactants, not the products, but how quickly the reaction proceeds. That measurable "how quickly" is called the reaction rate. In practice you measure the average rate of change of any quantity that tracks how much reaction has occurred, gas volume collected, mass lost, or colour change. The standard concentration rate (in mol L⁻¹ s⁻¹) is the change in concentration of a reactant or product per unit time, obtained from those measurements using amounts and volume.

Concentration change

What Is Monitored: Reactant or product concentration
Units: mol/L/s
Example Reaction: Coloured ion reactions, titration

Mass change

What Is Monitored: Decrease in mass as gas escapes
Units: g/s
Example Reaction: CaCO₃ + HCl → CO₂

Gas volume

What Is Monitored: Volume of gas collected
Units: mL/s
Example Reaction: Zn + H₂SO₄ → H₂

Colour change

What Is Monitored: Absorbance or light transmission
Units: Arbitrary units/s
Example Reaction: Permanganate decolourisation

Reaction rate is not constant throughout a reaction, it is typically fastest at the start (when reactant concentrations are highest) and slows progressively as reactants are consumed. A rate versus time graph shows a decreasing curve that approaches zero as the reaction nears completion.

Must Know: When describing how to measure reaction rate in an investigation, specify both what is being measured and the time interval. “Measuring the volume of CO₂ gas collected every 30 seconds using a gas syringe” is a complete method. “Measuring the gas” is not.
Common Error: Students state that reaction rate is constant throughout a reaction. It is not, rate decreases over time as reactant concentration falls. If asked for the rate at a specific time from a graph, calculate the gradient of the tangent at that point, not the overall gradient.

Reaction rate = Δ[concentration] / Δtime (mol L⁻¹ s⁻¹). Can also be measured as g/s (mass loss) or mL/s (gas volume). Rate is fastest at the start when reactant concentration is highest, and decreases as reactants are consumed, it is not constant.

Pause, copy the highlighted definition into your book before moving on.

Match it: Match each measurement method to the unit used to express reaction rate.

  • Volume of gas collected
  • Change in concentration
  • Loss of mass as gas escapes
  • Colour change via absorbance
  • Arbitrary units/s
  • g/s
  • mol L⁻¹ s⁻¹
  • mL/s
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Collision Theory, The Particle-Level Explanation

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Collision Theory, The Particle-Level Explanation
core concept

We just saw that reaction rate is measurable and decreases over time. That raises a question: why don’t all collisions between reactant particles result in a reaction, what makes some collisions effective and others not? This card answers it → two conditions must be met simultaneously: sufficient kinetic energy (≥ Eₐ) AND correct orientation of the colliding particles.

Chemical reactions don’t happen just because reactant particles are present, they happen only when particles collide in exactly the right way with exactly enough energy.

Collision theory states: for a chemical reaction to occur, reactant particles must collide. However, not every collision results in a reaction. An effective collision is one that results in the formation of products. For a collision to be effective, two conditions must both be met simultaneously:

  1. Sufficient energy: the colliding particles must have kinetic energy equal to or greater than the activation energy (Eₐ). If below Eₐ, the particles simply bounce apart.
  2. Correct orientation: the particles must collide with the reactive parts of their molecules facing each other. Even with sufficient energy, a collision between the wrong parts of two molecules will not produce the desired products.

The proportion of collisions that are effective is small for most reactions at room temperature, which is why reactions that are thermodynamically favourable can still be slow.

Must Know: Always state BOTH conditions for an effective collision in HSC answers, sufficient energy (≥ Eₐ) AND correct orientation. (Orientation matters most for collisions between molecules; for simple ions or single atoms, geometry is much less relevant.) A response that mentions only energy will not usually score full marks.
Common Error: Students confuse collision frequency with collision effectiveness. Increasing temperature increases both, but they are different things. Collision frequency is how often particles collide; collision effectiveness is whether those collisions have sufficient energy AND correct orientation.
Insight: For a typical reaction at room temperature, only about 1 in 10⁹ collisions is effective. This seems impossibly low, yet reactions still proceed at measurable rates because the total number of collisions per second in a sample is astronomical (on the order of 10³⁰ per second in a typical solution). Even a tiny fraction of an enormous number gives a measurable reaction rate.

Collision theory: an effective collision requires (1) kinetic energy ≥ activation energy (Eₐ) AND (2) correct orientation of reacting particles. Most collisions are ineffective. Collision frequency = how often particles meet; effectiveness = whether they overcome Eₐ with correct geometry.

Add the highlighted definition to your notes before the check below.

Odd one out: Three of these are conditions or features of an effective collision. Which one does NOT belong?

Activation Energy and Energy Diagrams

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Activation Energy and Energy Diagrams
core concept

We just saw that effective collisions require kinetic energy ≥ Eₐ. That raises a question: how do we represent activation energy and the overall energy change of a reaction visually, and how do we read Eₐ and ΔH from an energy diagram? This card answers it → on an energy profile, Eₐ is the height from reactants to the transition state peak; ΔH is the height difference between reactants and products.

The activation energy (Eₐ) is the minimum kinetic energy that colliding particles must have for the collision to be effective. It represents the energy required to break the bonds in the reactants sufficiently to allow atoms to rearrange into products.

An energy diagram (reaction progress diagram) plots the energy of the reacting system against the progress of the reaction from reactants to products:

Energy rises to the activation-energy peak, then falls to products below the reactants (ΔH < 0).

Key features: reactants at their initial energy level (left); the transition state (activated complex) at the peak; products at their final energy level (right). Eₐ is measured from the reactant energy level to the peak. ΔH is the difference between reactant and product energy levels. For an exothermic reaction, products are lower than reactants (ΔH < 0). For endothermic, products are higher (ΔH > 0).

Must Know: Eₐ is always measured from the reactant energy level to the peak, not from zero and not from the product level. In exothermic reactions, students sometimes confuse ΔH with Eₐ. They are different quantities on the same Eₐ is the barrier height; ΔH is the net energy change.
Common Error: Students draw the activation energy arrow starting from zero on the y-axis, or from the product level. It must start from the reactant energy level and end at the transition state peak. Practise labelling the diagram correctly until it is automatic.

Activation energy (Eₐ) is measured from the reactant energy level to the transition state peak, not from zero. ΔH is measured from reactant to product energy levels. Exothermic: products lower than reactants (ΔH < 0). Endothermic: products higher (ΔH > 0). Eₐ and ΔH are independent.

Pause, write the highlighted definitions into your book.

Fill the gap: On an energy diagram, the activation energy (Eₐ) is measured from the [___] energy level to the transition state peak.

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What Variables Affect Reaction Rate, Preview

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What Variables Affect Reaction Rate, Preview
core concept

We just saw that Eₐ is the energy barrier particles must overcome. That raises a question: what experimental variables can we change to alter how often particles successfully overcome that barrier, and which mechanism does each one exploit? This card answers it → temperature affects both collision frequency and proportion exceeding Eₐ; concentration/surface area affect frequency only; catalysts lower Eₐ.

Every variable that affects reaction rate does so by changing either the frequency of collisions, the proportion of collisions with sufficient energy, or both. These four factors are explored in full in L12:

VariableEffect on Collision FrequencyEffect on Proportion Exceeding EₐNet Effect on Rate
Increase temperatureIncreasesIncreases (more exceed Eₐ)Rate increases
Increase concentrationIncreasesNo changeRate increases
Increase surface areaIncreases (at surface)No changeRate increases
Add catalystNo significant changeIncreases (lower Eₐ)Rate increases
Must Know: Every explanation of a rate factor in an HSC answer must use collision theory language: activation energy, collision frequency, effective collision, proportion of particles exceeding Eₐ. A response that says “increasing temperature makes particles move faster and react more” without using these terms will not score full marks.

Four rate factors: Temperature, increases collision frequency AND proportion exceeding Eₐ. Concentration/Surface area, increase collision frequency only. Catalysts, lower Eₐ, increasing the proportion that exceed it. Use collision theory language in every HSC explanation.

Add the highlighted table to your notes before the check below.

Match it: Match each factor to its effect on collision frequency and the proportion of particles exceeding Eₐ.

  • Increasing temperature
  • Increasing concentration
  • Adding a catalyst
  • Increasing surface area
  • Increases collision frequency at reaction interface only; no change to proportion exceeding Eₐ
  • Lowers Eₐ, increasing proportion exceeding it; no significant change to collision frequency
  • Increases collision frequency only; no change to proportion exceeding Eₐ
  • Increases both collision frequency AND proportion exceeding Eₐ

Glow Sticks, Activation Energy and Temperature in Action

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Glow Sticks, Activation Energy and Temperature in Action
core concept

We just saw that temperature affects both collision frequency and the proportion exceeding Eₐ. That raises a question: what does this look like in a tangible, observable experiment, where increasing temperature visibly changes reaction rate in real time? This card answers it → a glow stick in warm water glows brighter and fades faster; in cold water, dimmer but for longer. Because each stick starts with the same amount of reactant, the brightness-and-duration trade-off is the intended simplified model (in reality, temperature can also affect efficiency and side reactions).

A glow stick contains two chemical compartments: an outer tube containing hydrogen peroxide (H₂O₂) and a fluorescent dye, and an inner glass vial containing a phenyl oxalate ester. When snapped, the two solutions mix and a chemiluminescent reaction produces light.

At higher temperature (warm water): particles have greater average kinetic energy; a larger proportion of collisions exceed Eₐ; more effective collisions occur per second; reaction rate is faster; the glow is brighter but shorter (reactants consumed more quickly).

At lower temperature (freezer): particles have lower average kinetic energy; fewer collisions exceed Eₐ; fewer effective collisions per second; reaction rate is slower; the glow is dimmer but lasts longer.

Crucially, the total amount of light produced is approximately the same at both temperatures (same amount of reactant), only the rate of production differs.

Must Know: The short answer question for this lesson will ask you to explain, using collision theory, why a glow stick glows more brightly in hot water. Structure your answer: temperature effect on kinetic energy → proportion of particles exceeding Eₐ → effective collision frequency → reaction rate → brightness.
Common Error: Students state that hot water “gives energy to the reaction” or “adds energy to the glow stick.” This is imprecise. The correct explanation is that hot water increases the average kinetic energy of the particles inside the glow stick, increasing the proportion that exceed the activation energy barrier, it does not change Eₐ.

Hot water → higher average KE → more particles exceed Eₐ → more effective collisions/s → faster rate → brighter but shorter glow. Cold water → fewer particles exceed Eₐ → slower rate → dimmer but longer glow. Total light is the same, only the rate of consumption differs.

Pause, write the highlighted explanation into your book.

Odd one out: Three of these are true statements about a glow stick in hot water compared to cold water. Which one does NOT belong?

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Worked examples

Worked examples · reveal as you go

Worked example +5 XP on full reveal

A student monitors the reaction between marble chips (CaCO3) and hydrochloric acid by measuring CO2 gas produced over time: 0 s = 0 mL; 30 s = 18 mL; 60 s = 32 mL; 90 s = 42 mL; 120 s = 48 mL; 150 s = 51 mL; 180 s = 52 mL. (a) Calculate the average rate between 0 and 60 s. (b) Calculate the average rate between 120 and 180 s. (c) Explain why the rate in (b) is lower than in (a) using collision theory.

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(a) Rate from 0 to 60 s
Rate = Δvolume ÷ Δtime = (32 − 0) mL ÷ (60 − 0) s = 32 ÷ 60 = 0.53 mL/s
Use the rate formula: change in volume divided by change in time over the given interval.
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(b) Rate from 120 to 180 s
Rate = Δvolume ÷ Δtime = (52 − 48) mL ÷ (180 − 120) s = 4 ÷ 60 = 0.067 mL/s
Notice the rate is much lower in the second interval because less gas is being produced in the same time span.
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(c) Why the rate decreased
Between 120 and 180 seconds, the reaction is nearly complete, the concentration of HCl has decreased significantly.
As reactants are consumed, fewer particles remain in the solution to collide.
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Lower concentration → decreased collision frequency
Fewer H+ ions per unit volume means fewer H+ and CaCO3 surface particles collide per second. Since temperature is constant, the proportion of effective collisions is unchanged, only the total number of collisions decreases.
Collision frequency (how often particles meet) is different from collision effectiveness (whether they react). Concentration mainly changes the frequency; temperature mainly changes the effectiveness (and also raises the frequency a little). These are dominant effects, not exclusive ones.
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Worked example +5 XP on full reveal

A reaction has an activation energy of 85 kJ/mol and releases 40 kJ/mol of energy. (a) Is the reaction exothermic or endothermic? (b) Describe the energy diagram, labelling Ea, ΔH, and the transition state. (c) Show the effect of adding a catalyst on the same diagram.

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(a) Is the reaction exothermic or endothermic?
The reaction releases 40 kJ/mol, so ΔH = −40 kJ/mol → EXOTHERMIC
Negative ΔH means energy is released. Products are lower in energy than reactants.
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(b) Energy diagram description
Set reactant energy level as the reference (baseline). The energy curve rises to the transition state peak at 85 kJ above the reactant level. This is Ea = 85 kJ/mol. The curve then falls to the product level, which is 40 kJ below the reactant level.
Remember: Ea is measured from the reactant level to the peak, NOT from zero. ΔH is measured from reactants to products.
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(c) Effect of adding a catalyst
A catalyst lowers the activation energy. On the same diagram, draw a second reaction curve with a lower peak (e.g., 55 kJ above reactant level instead of 85 kJ). Label the lower peak "Ea (catalysed)" and the original peak "Ea (uncatalysed)".

Key point: Reactant and product energy levels are UNCHANGED, the catalyst does not alter ΔH, only the barrier height.
A catalyst provides an alternative reaction pathway with lower Ea. It speeds up both forward and reverse reactions equally.
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Put the method in order

Sort the steps +7 XP

Drag these statements into the correct order to explain why increasing concentration increases reaction rate (using collision theory).

More effective collisions occur per second → reaction rate increases.
Increasing concentration means more solute particles per unit volume.
The proportion of collisions that are effective is unchanged, temperature and Eₐ are constant.
More particles per unit volume → more frequent collisions between reactant particles.
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What you'll master, and the words for it

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Formula reference · this lesson
core formula
📐

Key Concepts, This Lesson

$\text{Reaction rate} = \dfrac{\Delta[\text{concentration}]}{\Delta t} \;\left(\text{mol L}^{-1}\text{s}^{-1}\right)$
Also expressed as $\Delta\text{mass}/\Delta t$ (g/s) or $\Delta\text{volume}/\Delta t$ (mL/s). Rate decreases over time as reactant concentration falls.
Effective collision: energy $\geq E_a$ AND correct orientation
Both conditions must be met simultaneously, one alone is not sufficient
$T\uparrow \;\Rightarrow\; \overline{KE}\uparrow \;\Rightarrow\; \text{more particles exceed }E_a \;\Rightarrow\; \text{rate}\uparrow$
Temperature is the only factor that changes the proportion of particles exceeding $E_a$

Common misconception

Common errors · the 3 traps that cost marks

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Common misconception

A catalyst increases the amount of product formed at equilibrium.

Fix: A catalyst speeds up both forward and reverse reactions equally, so equilibrium is reached faster but the equilibrium position and product yield do not change. Catalysts lower activation energy but do not alter ΔH or the thermodynamic favourability of a reaction.

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Temperature increases rate only because molecules collide more frequently

A student explains: "Increasing temperature makes molecules move faster, so they collide more often, so the rate increases." This is incomplete and will not score full marks.

Fix: Collision frequency does increase slightly with temperature, but the dominant effect is that a much larger proportion of molecules now have energy ≥ Eₐ. The Maxwell-Boltzmann distribution shifts to higher energies, the area to the right of the Eₐ threshold increases significantly. HSC answers must reference activation energy, kinetic energy distribution, and the proportion of effective collisions, not just "more frequent collisions."

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Correct orientation is not needed if particles have enough energy

A student states: "As long as particles have energy ≥ Eₐ, the reaction will occur."

Fix: Both conditions must be met simultaneously for an effective collision. Sufficient energy alone is not enough, particles must also collide with the correct orientation (reactive parts facing each other). Even with kinetic energy ≥ Eₐ, a collision between the wrong parts of two molecules will not produce products. In HSC responses, always state BOTH: sufficient energy (≥ Eₐ) AND correct orientation.

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Drill, then revisit

Quick-fire practice · 5 reps +2 XP per reveal

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(4 marks) State the two conditions required for an effective collision. For each condition, explain what happens when that condition is not met (i.e., what happens to the particles). (2 marks each)

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(4 marks) A reaction has an activation energy of 60 kJ/mol and produces 35 kJ/mol of energy. (a) Is the reaction exothermic or endothermic? (b) On a labelled energy diagram, identify and give numerical values for Eₐ and ΔH. (c) Explain whether the reaction would proceed faster or slower if the activation energy were 30 kJ/mol instead of 60 kJ/mol (assume the same temperature). (1 + 2 + 1 marks)

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(5 marks) A glow stick is cracked and placed in a water bath at 5°C. An identical glow stick is placed in a water bath at 45°C. (a) Predict which glow stick will be brighter and which will last longer. (b) Using collision theory, explain why the glow stick at 45°C glows more brightly. Your explanation must reference activation energy, kinetic energy of particles, and the proportion of effective collisions. (c) The total amount of light produced by both glow sticks is the same. Explain this observation in terms of the amount of reactant and reaction rate. (1 + 3 + 1 marks)

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Experimental data for CaCO₃ + HCl: 0 s = 0 mL CO₂; 30 s = 18 mL CO₂. Calculate the average rate between 0 and 30 s.

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A reaction has activation energy 75 kJ/mol and releases 45 kJ/mol of energy. (a) Is it exothermic or endothermic? (b) What is the energy of the transition state above the reactant energy level?

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Revisit your thinking

Go back to your Think First response. You've now seen why a glow stick performs so differently in warm water vs iced water. Using activation energy (Ea), the Maxwell-Boltzmann distribution, and effective collision frequency, write a 4-sentence explanation of the rate difference. Then explain why two identical sticks, starting with the same amount of reactant, produce a similar total amount of light overall, just released at different rates (the brightness-and-duration trade-off is the intended simplified model).

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Interactive Tool, Rates of Reaction Open fullscreen ↗
The Reaction Rates tool shows that increasing temperature increases the rate because…
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Multiple choice

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Multiple choice
+5 XP per correct · +25 XP all-correct

Pick your answer, then rate your confidence. That tells the system what to drill next.

Fill the blanks +4 XP

Complete this model answer about temperature and reaction rate:

Increasing temperature increases the of the particles. The Maxwell-Boltzmann distribution shifts to the , meaning a greater proportion of particles have energy greater than or equal to the . This increases the frequency of collisions per unit time, so the reaction rate .

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Short answer
UnderstandBand 3

Q1. (4 marks) State the two conditions required for an effective collision. For each condition, explain what happens when that condition is not met (i.e., what happens to the particles). (2 marks each)

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ApplyBand 4

Q2. (4 marks) A reaction has an activation energy of 60 kJ/mol and produces 35 kJ/mol of energy. (a) Is the reaction exothermic or endothermic? (b) On a labelled energy diagram, identify and give numerical values for Eₐ and ΔH. (c) Explain whether the reaction would proceed faster or slower if the activation energy were 30 kJ/mol instead of 60 kJ/mol (assume the same temperature). (1 + 2 + 1 marks)

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EvaluateBand 5

Q3. (5 marks) A glow stick is cracked and placed in a water bath at 5°C. An identical glow stick is placed in a water bath at 45°C. (a) Predict which glow stick will be brighter and which will last longer. (b) Using collision theory, explain why the glow stick at 45°C glows more brightly. Your explanation must reference activation energy, kinetic energy of particles, and the proportion of effective collisions. (c) The total amount of light produced by both glow sticks is the same. Explain this observation in terms of the amount of reactant and reaction rate. (1 + 3 + 1 marks)

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