Chemical Energetics is about the energy change that comes with every chemical reaction. As bonds are broken and formed, energy is either given out to the surroundings or taken in from them. These notes cover how to classify reactions as exothermic or endothermic, how to read and draw energy profile diagrams, and how to calculate the enthalpy change ΔH from bond energies.
Examiners link this topic to combustion, neutralisation and fuels questions across the paper, and bond-energy calculations appear in the O-Level paper. Most lost marks come from vague temperature explanations, missing ΔH signs and miscounted bonds. Learn the definitions word for word and write out the working in full.
What these notes cover
- •Exothermic and endothermic reactions
- •Classifying common processes
- •Sign convention for ΔH
- •Energy profile diagrams
- •Activation energy, Eₐ
- •Effect of a catalyst
- •Bond breaking and bond forming
- •Calculating ΔH from bond energies
- •Worked examples
- •Hydrogen as a fuel and the hydrogen–oxygen fuel cell
- •Common mistakes and exam technique
- •Summary
Key definitions
- Exothermic reaction
- A reaction in which energy is transferred to the surroundings, often as heat. The temperature of the reaction mixture and its surroundings increases, and ΔH is negative.
- Endothermic reaction
- A reaction in which energy is taken in from the surroundings. The temperature of the reaction mixture and its surroundings decreases, and ΔH is positive.
- Activation energy (Eₐ)
- The minimum energy that colliding particles must have for a reaction to occur.
- Surroundings
- Everything around the reaction mixture, such as the water the chemicals are dissolved in, the beaker, the thermometer and the air. A thermometer in the mixture measures the temperature of the surroundings.
- Fuel cell
- A device that produces electricity directly from a chemical reaction. In the hydrogen–oxygen fuel cell, hydrogen and oxygen react to form water, and the energy is released as electrical energy rather than as heat.
Key formulas
Activation energy from a profile
Eₐ = energy at peak − energy of reactants
Always measured from the reactant level up to the peak, never from the x-axis and never from the product level.
Enthalpy change from a profile
ΔH = energy of products − energy of reactants
Products below reactants gives a negative ΔH (exothermic). Products above reactants gives a positive ΔH (endothermic).
Enthalpy change from bond energies
ΔH = total energy absorbed breaking bonds − total energy released forming bonds
Bond energies are given in the question in kJ/mol. Multiply each bond count by the coefficients in the balanced equation.
Sign convention
Exothermic: ΔH < 0; Endothermic: ΔH > 0
Units are kJ/mol. The sign is part of the answer, so write ΔH = −57 kJ/mol for a neutralisation, not 57 kJ/mol.
Neutralisation (ionic equation)
H⁺(aq) + OH⁻(aq) → H₂O(l)
This is the energy-releasing change in every neutralisation between a dilute acid and an alkali. In NaOH + HCl, the Na⁺ and Cl⁻ ions are spectator ions.
Combustion of hydrogen / fuel cell reaction
2H₂(g) + O₂(g) → 2H₂O(l), ΔH < 0
This is the overall reaction in the hydrogen–oxygen fuel cell. The only product is water.
Worked example
Methane burns in oxygen: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g). Bond energies: C–H = 410 kJ/mol, O=O = 496 kJ/mol, C=O = 805 kJ/mol, O–H = 460 kJ/mol. Calculate ΔH and state whether the reaction is exothermic or endothermic.
- 1Bonds broken: one CH₄ has 4 C–H bonds and two O₂ molecules have 2 O=O bonds, so energy absorbed = (4 × 410) + (2 × 496) = 1640 + 992 = 2632 kJ.
- 2Bonds formed: one CO₂ has 2 C=O bonds and two H₂O molecules have 2 × 2 = 4 O–H bonds, so energy released = (2 × 805) + (4 × 460) = 1610 + 1840 = 3450 kJ.
- 3Subtract: ΔH = energy absorbed − energy released = 2632 − 3450 = −818 kJ/mol.
- 4ΔH is negative, so more energy is released forming bonds than is absorbed breaking them, and the reaction is exothermic.
Watch out: The most common error in this question type is forgetting the coefficients. The 2 in front of O₂ gives two O=O bonds, and the 2 in front of H₂O gives four O–H bonds.
Common mistakes
Thinking bond breaking releases energy
Wrong: Breaking bonds releases energy.
Right: Bond breaking always absorbs energy (endothermic) and bond forming always releases energy (exothermic). A reaction is exothermic overall only because the energy released forming new bonds is greater than the energy absorbed breaking old bonds.
Saying a catalyst changes ΔH
Wrong: A catalyst lowers the energy change or makes the reaction more exothermic.
Right: A catalyst lowers Eₐ only, by providing an alternative reaction pathway. The reactant and product levels are unchanged, so ΔH is unchanged.
Misexplaining why an endothermic mixture gets cold
Wrong: The mixture gets cold in an endothermic reaction because the chemicals are losing energy.
Right: The chemicals gain energy because they absorb thermal energy from the surroundings. It is the temperature of the surroundings (the water, the beaker, the thermometer) that falls.
Questions students ask
What is the difference between exothermic and endothermic reactions?
An exothermic reaction transfers energy to the surroundings, so the temperature of the surroundings rises and ΔH is negative. An endothermic reaction takes in energy from the surroundings, so the temperature falls and ΔH is positive. Combustion and neutralisation are exothermic; thermal decomposition and dissolving ammonium nitrate in water are endothermic.
Does a catalyst change the enthalpy change ΔH?
No. A catalyst provides an alternative pathway with a lower activation energy, which is drawn as a dashed curve with a lower peak on the same axes. The reactant and product levels stay the same, so ΔH is unchanged.
How do you calculate ΔH from bond energies?
First total the energy absorbed to break all the bonds in the reactants and the energy released to form all the bonds in the products, multiplying by the coefficients in the balanced equation. Then calculate ΔH = energy absorbed − energy released. A negative answer means the reaction is exothermic and a positive answer means it is endothermic.
What are the advantages and disadvantages of hydrogen as a fuel?
Hydrogen is clean at the point of use because the only product is water, and a fuel cell converts chemical energy to electrical energy directly, which is more efficient than burning the fuel. However, hydrogen is highly flammable, bulky and difficult to store as a gas, and expensive to produce. Producing it may also use energy from fossil fuels, so carbon dioxide can still be released at the power station.
In the full PDF
- •Labelled energy profile diagrams for exothermic, endothermic and catalysed reactions
- •Further worked examples on classifying reactions from temperature data and reading Eₐ and ΔH off a profile
- •More bond-energy calculations: H₂ + Cl₂, the decomposition of HI, and the endothermic N₂ + O₂ reaction
- •The remaining misconception and model exam phrasing for temperature-change answers
- •Checklists for drawing energy profiles and setting out bond-energy tables
- •One-glance recap table and the five key takeaways for Topic 9