Reaction profiles — AQA GCSE Chemistry
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Reaction profiles explained
Reaction profiles covers 4 specification points, including Chemical reactions can occur only when… and The minimum amount of energy that particles….
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Open a point below for its explanation and revision guidance.
Chemical reactions can occur only when reacting particles collide with each other and with sufficient energy.
This is the collision theory explanation of reaction rates. For a reaction to happen, reactant particles must collide, and the collision must involve at least a certain minimum energy called the activation energy. A collision with less energy than the activation energy does not lead to reaction; the particles simply bounce apart. For example, in the reaction between hydrogen and oxygen, most collisions at room temperature lack sufficient energy, which is why the mixture can remain unreactive until a spark provides activation energy. Reaction profiles show the energy of reactants and products and the activation energy as the height from the reactants' energy level to the peak of the curve. Increasing temperature, concentration, pressure or surface area, or adding a catalyst, increases the frequency or success of collisions and so increases the rate.
The minimum amount of energy that particles must have to react is called the activation energy.
Particles collide constantly, but most collisions do nothing because the particles simply bounce apart. For a reaction to happen, the colliding particles must have at least a certain amount of kinetic energy so that existing bonds can begin to break and new arrangements can form. This threshold is the activation energy. Picture two molecules approaching: if their combined energy is below the threshold, they rebound unchanged; if it equals or exceeds the threshold, they can react. Activation energy is measured in joules, commonly kJ/mol, and is shown on a reaction profile as the vertical gap from the reactant energy level up to the peak of the curve. A high activation energy means few collisions succeed, so the reaction is slow at a given temperature; adding a catalyst provides an alternative pathway with a lower activation energy.
Reaction profiles can be used to show the relative energies of reactants and products, the activation energy and the overall energy change of a reaction.
A reaction profile, also called an energy level diagram, is a graph of energy against the progress of a reaction. The reactants start at one energy level and the products finish at another, joined by a curved line that rises to a peak as bonds break and rearrange. Three features carry the chemistry. The relative heights of the start and finish levels show whether the products are lower or higher in energy than the reactants. The vertical gap from the reactants up to the peak is the activation energy. The difference between the reactant and product levels is the overall energy change. If products sit lower, energy is released to the surroundings and the reaction is exothermic; if products sit higher, energy is taken in and the reaction is endothermic.
Students should be able to: draw simple reaction profiles (energy level diagrams) for exothermic and endothermic reactions showing the relative energies of reactants and products, the activation energy and the overall energy change, with a curved line to show the energy as the reaction proceeds use reaction profiles to identify reactions as exothermic or endothermic explain that the activation energy is the energy needed for a reaction to occur.
This practical skill combines drawing and interpreting reaction profiles. To draw one, mark energy on the vertical axis and reaction progress on the horizontal axis, place the reactant level, then draw a curved line rising to a peak before falling to the product level. For an exothermic reaction the product level sits below the reactants; for an endothermic reaction it sits above, but the curve still descends from the peak to reach it. Label the activation energy as the gap from the reactants to the peak, and the overall energy change as the gap between reactant and product levels. To identify a reaction type, compare the two levels: lower products mean energy is released, so the reaction is exothermic; higher products mean energy is absorbed, so it is endothermic. Activation energy is the minimum energy colliding particles need for a reaction to occur.
Your focus
- State the two conditions needed for a chemical reaction to occur.
- Define activation energy and identify it on a reaction profile.
- Explain how changes in conditions affect the rate of reaction using collision theory.
Show all 12 objectives
- Define activation energy as the minimum energy particles need to react.
- Identify the activation energy on a drawn reaction profile.
- Explain why most collisions do not lead to reaction.
- Interpret a reaction profile to identify relative reactant and product energies.
- Determine the activation energy and overall energy change from a profile.
- Classify a reaction as exothermic or endothermic using its profile.
- Draw labelled reaction profiles for exothermic and endothermic reactions.
- Use a reaction profile to classify a reaction as exothermic or endothermic.
- Explain activation energy as the energy needed for a reaction to occur.
Reaction profiles exam tips
Marking Points
- States that particles must collide for a reaction to occur, and that not every collision is successful.
- Explains that a collision must have at least the activation energy for a reaction to take place.
- Describes activation energy as the minimum energy that colliding particles need to react.
- Uses reaction profiles to identify the energy of reactants, the energy of products and the activation energy.
- Links successful collisions to reaction rate, explaining that more frequent or more energetic collisions increase the rate.
- Activation energy is defined as the minimum energy that colliding particles must possess for a reaction to occur.
- It is the energy barrier that must be overcome before reactant bonds can break and new bonds can form.
- On a reaction profile it is represented by the vertical distance from the reactants' energy level to the highest point of the curve.
- Collisions with energy below the activation energy do not lead to reaction; the particles separate unchanged.
- Activation energy is usually expressed in kJ/mol and is not the same as the overall energy change of the reaction.
- A catalyst lowers the activation energy by providing an alternative reaction pathway.
- A reaction profile plots energy on the vertical axis against reaction progress on the horizontal axis.
- The reactant and product levels show relative energy: products lower than reactants indicates an exothermic change, products higher indicates an endothermic change.
- The activation energy is the vertical distance from the reactant energy level to the peak of the curve.
- The overall energy change is the vertical difference between the reactant level and the product level.
- The curved line represents how energy varies as bonds break and new bonds form during the reaction.
- Profiles allow comparison of different reactions, including the effect of a catalyst, which lowers the peak.
- Draw axes with energy vertical and reaction progress horizontal, and place reactant and product levels at appropriate heights.
- Use a single curved line that rises from the reactants to a peak and then descends to the products.
- For an exothermic profile, place the product level below the reactant level; for an endothermic profile, place it above.
- Label the activation energy as the vertical gap from the reactant level to the peak.
- Label the overall energy change as the vertical difference between reactant and product levels.
- Identify a reaction as exothermic when products are lower in energy and endothermic when products are higher.
- Explain activation energy as the minimum energy that colliding particles must have for a reaction to occur.
Examiner Tips
- 💡Use the phrase 'collide with sufficient energy' when explaining why a reaction occurs, and refer to activation energy by name.
- 💡On a reaction profile, mark the activation energy as the vertical distance from the reactants' energy level to the peak of the curve.
- 💡When explaining rate changes, state whether collision frequency, collision energy or both are affected.
- 💡When labelling a profile, draw the activation energy arrow from the reactant level to the top of the curve, not from the products.
- 💡Use the phrase 'minimum energy' in definitions; answers that say only 'the energy needed' may be treated as incomplete.
- 💡Check the units: if the axis is in kJ/mol, give activation energy values in kJ/mol rather than J.
- 💡Label both axes and mark the reactant level, product level, peak, activation energy and overall energy change clearly.
- 💡Use dashed horizontal lines from the reactant and product levels to the energy axis to make comparisons easier to read.
- 💡State whether the reaction is exothermic or endothermic and justify it by referring to the relative heights of the two levels.
- 💡Sketch the curve first, then add labels and arrows so the diagram stays clear and uncluttered.
- 💡Use a ruler for the vertical energy arrows but keep the reaction pathway line curved.
- 💡When asked to identify a reaction type, quote the relative heights of the reactant and product levels as your evidence.
Common Mistakes
- Claiming that any collision causes a reaction; the correction is that only collisions with at least the activation energy are successful.
- Confusing activation energy with the overall energy change of the reaction; the correction is that activation energy is the energy needed to start the reaction, while overall energy change is the difference between reactants and products.
- Saying that a catalyst increases the energy of the particles; the correction is that a catalyst provides an alternative pathway with a lower activation energy.
- Confusing activation energy with the overall energy change: the overall change is the difference between reactant and product energy levels, whereas activation energy is measured from the reactants up to the peak.
- Thinking that every collision causes a reaction: only collisions with energy at or above the activation energy are successful.
- Believing a catalyst increases the energy of the particles: a catalyst lowers the activation energy instead, leaving particle energies unchanged.
- Measuring the overall energy change from the peak to the products instead of from the reactants to the products.
- Drawing the activation energy arrow starting at the products when the reaction is endothermic; it must always start at the reactants.
- Treating the peak as a stable intermediate that can be isolated, when it represents the highest-energy arrangement along the pathway.
- Drawing the product level below the reactants for an endothermic reaction: endothermic products must be higher in energy.
- Omitting the curved line and joining levels with straight lines, which fails to show the energy barrier.
- Labelling the overall energy change as the distance from the peak to the products rather than from reactants to products.
- Drawing the curve rising from the peak to the products for an endothermic reaction; correct by ensuring the curve always falls from the peak to the product level.