Collision Theory & Reaction Rate
Cyalume chemiluminescent light sticks — developed by American Cyanamid in 1973 and adopted by the US military by 1978 — were tested from −20°C in Arctic deployments to +40°C in desert operations. The same stick that glowed steadily for 12 hours in the cold exhausted in under 90 minutes in the heat. Identical chemistry, identical total energy — only the rate changed.
Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
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.
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
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
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
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: the change in concentration of a reactant or product per unit time. Rate can be measured by any observable quantity that tracks how much reaction has occurred — gas volume collected, mass lost, colour change.
Concentration change
Mass change
Gas volume
Colour change
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.
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
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:
- 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.
- 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.
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?
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:
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).
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.
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:
| Variable | Effect on Collision Frequency | Effect on Proportion Exceeding Eₐ | Net Effect on Rate |
|---|---|---|---|
| Increase temperature | Increases | Increases (more exceed Eₐ) | Rate increases |
| Increase concentration | Increases | No change | Rate increases |
| Increase surface area | Increases (at surface) | No change | Rate increases |
| Add catalyst | No significant change | Increases (lower Eₐ) | Rate increases |
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
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 faster; in cold water, dimmer but longer — the total light output is the same.
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.
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?
Worked examples · reveal as you go
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.
Rate = Δvolume ÷ Δtime = (32 − 0) mL ÷ (60 − 0) s = 32 ÷ 60 = 0.53 mL/s
Rate = Δvolume ÷ Δtime = (52 − 48) mL ÷ (180 − 120) s = 4 ÷ 60 = 0.067 mL/s
Between 120 and 180 seconds, the reaction is nearly complete — the concentration of HCl has decreased significantly.
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.
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.
The reaction releases 40 kJ/mol, so ΔH = −40 kJ/mol → EXOTHERMIC
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.
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.
Drag these statements into the correct order to explain why increasing concentration increases reaction rate (using collision theory).
Key Concepts — This Lesson
Common errors · the 3 traps that cost marks
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.
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."
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.
Quick-fire practice · 5 reps +2 XP per reveal
(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)
(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)
(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)
Experimental data for CaCO₃ + HCl: 0 s = 0 mL CO₂; 30 s = 18 mL CO₂. Calculate the average rate between 0 and 30 s.
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?
Go back to your Think First response. You've now seen exactly why Cyalume light sticks (American Cyanamid, 1973) perform so differently at −20°C vs +40°C. Using activation energy (Ea), the Maxwell-Boltzmann distribution, and effective collision frequency, write a 4-sentence explanation of the rate difference. Then explain why both sticks produce exactly the same total light — just at different rates.
Pick your answer, then rate your confidence — that tells the system what to drill next.
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 .
Q1. 8. (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)
Q2. 9. (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)
Q3. 10. (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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