CC15 · Heat energy changes in chemical reactions · Topic 7 — Rates of reaction and energy changes | GCSE Edexcel Combined Science

Revise the key ideas

An exothermic change transfers heat energy from the reacting system to the surroundings, so the surrounding temperature can rise. The system loses energy overall.
An endothermic change takes in heat from surroundings, which can cool. The system gains energy overall.
Temperature observations depend on heat exchange and insulation.
Common exothermic examples include combustion and many acid–alkali neutralisations and displacement reactions. Do not infer every reaction's energy change from its category alone without evidence.
Dissolving salts in water can release or absorb heat depending on the salt. Precipitation reactions can also show measurable temperature changes.
The same energy transfer can be described from two perspectives: an exothermic system loses heat while surroundings gain it. A temperature change in a solution reflects exchange with that solution and apparatus.
For the reverse of a reaction, the energy change has the opposite sign. An exothermic forward reaction has an endothermic reverse reaction.

For a solution reaction, measure suitable reagent volumes and starting temperatures, mix in an insulated container and record the highest or lowest temperature reached.
Use a polystyrene cup with a lid where appropriate to reduce heat exchange with the room. Support the cup securely in a beaker and use an appropriate thermometer or temperature probe.
Insulation and a lid reduce exchange with the room; stir and record the extreme temperature.
Temperature change ΔT = final relevant temperature − initial temperature. A rise gives positive ΔT for the measured surroundings; a fall gives negative ΔT.
For a comparison, control reagent amounts and concentrations, initial temperature, container and mixing. Repeat runs and use comparable recording intervals.
Heat loss during an exothermic reaction can make the observed maximum rise too small. Heat entering from the room can make an endothermic cooling appear smaller.
Wear eye protection and follow the specified dilute-reagent method. A rise alone does not identify every product, and a small change may be limited by instrument resolution.

A reaction profile plots energy vertically against reaction progress horizontally. Reaction progress is not a time axis.
Reactants begin at one energy level, the curve reaches a peak, and products end at another. The activation energy is the difference from the reactant level to the peak for the forward reaction.
Activation energy is the minimum energy barrier required for successful reaction. Exothermic reactions still need an initial energy input; a fuel does not necessarily ignite spontaneously.
For an exothermic reaction the product level is below reactants. Overall energy change ΔH = energy of products − energy of reactants is negative.
Products lower than reactants; activation energy remains positive.
For an endothermic reaction the product level is above reactants and ΔH is positive. The activation-energy arrow must still start at the reactant level.
Products higher than reactants; distinguish overall ΔH from the activation barrier.
A catalyst gives an alternative pathway with a lower peak, reducing activation energy. Reactant/product energy levels and overall energy change remain unchanged.
A catalyst changes the pathway and activation energy, not ΔH.
Forward and reverse activation energies are measured from different starting levels to the relevant peak. They are not necessarily equal.

Breaking chemical bonds requires energy and is endothermic. It is incorrect to say breaking a fuel's bonds itself releases energy.
Forming chemical bonds releases energy and is exothermic. Overall reaction energy depends on both bond breaking and bond making.
An exothermic reaction releases more energy in forming product bonds than it takes in to break reactant bonds. An endothermic reaction takes in more than it releases.
An exothermic result releases more in making bonds than it needs for breaking them.
Average bond energies are usually given in kJ mol⁻¹ of bonds. They are average values, so calculations using them may differ from the actual measured energy change of a reaction (its enthalpy change).
Count all bonds in every molecule, multiplying by the balanced-equation coefficients. A double bond has its given double-bond energy, not necessarily twice a single-bond value.

Estimated reaction energy change = total energy needed to break reactant bonds − total energy released forming product bonds. Use the sign to identify endothermic or exothermic.
For H₂ + Cl₂ → 2HCl, using H–H 436, Cl–Cl 243 and H–Cl 431 kJ mol⁻¹: break 436 + 243 = 679; form 2 × 431 = 862; ΔH = 679 − 862 = −183 kJ mol⁻¹ for the equation as written.
Equation: H₂ + Cl₂ → 2HCl; negative ΔH is exothermic.
For 2H₂ + O₂ → 2H₂O, using H–H 436, O=O 498 and O–H 463: break 2 × 436 + 498 = 1370; form 4 × 463 = 1852; ΔH = −482 kJ mol⁻¹ for two moles of water as written.
For CH₄ + 2O₂ → CO₂ + 2H₂O, count four C–H bonds and two O=O bonds broken; two C=O bonds and four O–H bonds formed. Coefficients are essential.
If a question asks per mole of one product rather than per balanced reaction, adjust accordingly. For the water equation above, per mole of water is −241 kJ mol⁻¹ using those average values.
Show the bond counts and both totals before subtracting. A negative answer indicates heat released overall; do not drop a minus sign just because the released energy has a positive magnitude.