Cambridge IGCSE0620

Exothermic and endothermic reactions

Chemistry 0620 Chapter Notes

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Exothermic and endothermic reactions
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In this chapter

  1. 1Exothermic and Endothermic Reactions
  2. 2Energy Level Diagrams
  3. 3Energy and Chemical Bonds
  4. 4Calculating Reaction Energy
  5. 5Fuels and Combustion
  6. 6Reversible Reactions and Energy

1. Exothermic and Endothermic Reactions

Every chemical reaction involves an energy change. This energy is usually transferred as heat. Reactions are classified into two types based on this energy transfer. An exothermic reaction is one that releases energy into the surroundings, causing the temperature of the surroundings to increase. Common examples include all combustion reactions (like burning a fuel), neutralisation of an acid with an alkali, and respiration in cells. An endothermic reaction is one that absorbs energy from the surroundings, causing the temperature of the surroundings to decrease. Examples include thermal decomposition (like heating calcium carbonate), photosynthesis, and the reaction in some cold packs. We measure this heat energy change as the 'enthalpy change', represented by the symbol ΔH. For an exothermic reaction, ΔH is negative because the system loses energy. For an endothermic reaction, ΔH is positive because the system gains energy.

Exothermic: Reactants → Products + Energy (ΔH is negative)

Endothermic: Reactants + Energy → Products (ΔH is positive)

Key term

Enthalpy Change (ΔH): The overall heat energy change in a chemical reaction at constant pressure, measured in kilojoules per mole (kJ/mol).

Examiner insight

Examiners look for a clear link between the observed temperature change and the direction of heat transfer (into or out of the reaction).

Common pitfall

Confusing the energy of the chemicals with the temperature of the surroundings. In an exothermic reaction, the chemicals lose energy (so products have lower energy than reactants), but they release this energy, making the surroundings hotter.

Fun fact

Instant cold packs use an endothermic reaction. Squeezing the pack breaks an inner pouch of water, which then dissolves a salt like ammonium nitrate, absorbing heat from the surroundings and making the pack feel icy cold.

Worked example 12 marks

A student adds 25 cm³ of dilute hydrochloric acid to 25 cm³ of dilute sodium hydroxide solution in a polystyrene cup. The initial temperature of both solutions was 20.0 °C. The highest temperature reached after mixing was 26.5 °C. Is the reaction exothermic or endothermic? Explain your answer.

  1. 1
    1. Identify the temperature change: The temperature increased from 20.0 °C to 26.5 °C.
  2. 2
    1. Relate temperature change to energy transfer: An increase in the temperature of the surroundings (the solution) means that heat energy has been released by the reaction.
  3. 3
    1. Define the reaction type: A reaction that releases heat energy into the surroundings is called an exothermic reaction.
  4. 4
    1. Conclusion: The reaction is exothermic because the temperature of the solution increased, indicating that heat was given out.

Recap

  • Exothermic reactions release energy, causing the surroundings to get hotter.
  • Endothermic reactions absorb energy, causing the surroundings to get colder.
  • The enthalpy change (ΔH) for an exothermic reaction is negative.
  • The enthalpy change (ΔH) for an endothermic reaction is positive.
  • Combustion and neutralisation are common examples of exothermic reactions.
  • Thermal decomposition and photosynthesis are common examples of endothermic reactions.

Quick check

  1. Is burning wood an exothermic or endothermic process?1 mark
  2. A reaction has an enthalpy change (ΔH) of +55 kJ/mol. Is it exothermic or endothermic?1 mark

2. Energy Level Diagrams

Energy level diagrams, also known as reaction profiles, are a visual way to represent the energy changes during a chemical reaction. The vertical axis represents potential energy, and the horizontal axis represents the 'reaction progress' (the course of the reaction from reactants to products). For an exothermic reaction, the reactants have more energy than the products, so the line for reactants is drawn higher than the line for products. The overall enthalpy change (ΔH) is shown as a downward arrow from the reactant level to the product level, and it is negative. For an endothermic reaction, the reactants have less energy than the products, so the reactants are drawn lower than the products. The ΔH is shown as an upward arrow and is positive. All reactions also need a certain amount of energy to get started, called the activation energy (Ea). This is shown as a 'hump' on the diagram that the reactants must overcome to turn into products.

Key term

Activation Energy (Ea): The minimum amount of energy that reacting particles must possess for a successful collision to occur and a reaction to start.

Examiner insight

Full marks for diagrams require all key labels: both axes, reactants, products, ΔH (with the correct sign or arrow direction), and Ea.

Common pitfall

Drawing the arrow for ΔH pointing in the wrong direction. The arrow must always start at the reactant's energy level and end at the product's energy level.

Worked example 14 marks

The combustion of methane is an exothermic reaction with an enthalpy change of -890 kJ/mol. Draw a fully labelled energy level diagram for this reaction.

  1. 1
    1. Draw the axes: Label the y-axis 'Energy' and the x-axis 'Reaction Progress'.
  2. 2
    1. Position reactants and products: Since the reaction is exothermic, the reactants (CH₄ + 2O₂) have higher energy than the products (CO₂ + 2H₂O). Draw a horizontal line for reactants near the top and another for products lower down.
  3. 3
    1. Draw the reaction path: Connect the reactant line to the product line with a curve that goes up first (the activation energy hump) and then down to the product level.
  4. 4
    1. Label ΔH: Draw a vertical arrow pointing downwards from the reactant level to the product level. Label it 'ΔH = -890 kJ/mol'.
  5. 5
    1. Label Ea: Draw a vertical arrow from the reactant level up to the peak of the curve. Label it 'Activation Energy (Ea)'.

Recap

  • Energy level diagrams show the energy of reactants and products over the course of a reaction.
  • For exothermic reactions, products are at a lower energy level than reactants.
  • For endothermic reactions, products are at a higher energy level than reactants.
  • The enthalpy change (ΔH) is the vertical difference between the energy of reactants and products.
  • Activation energy (Ea) is the energy 'hump' from the reactants to the transition state.

Quick check

  1. On an energy level diagram for an endothermic reaction, are the reactants shown higher or lower than the products?1 mark

3. Energy and Chemical Bonds

The energy changes in reactions are all about breaking and making chemical bonds. A fundamental rule to remember is: bond breaking requires energy, and bond making releases energy. In any reaction, the chemical bonds in the reactant molecules must first be broken. This step always requires an input of energy from the surroundings, so it is an endothermic process. Then, new chemical bonds are formed as the atoms rearrange to make the product molecules. This step always releases energy, so it is an exothermic process. The overall energy change of the reaction (ΔH) is the net result of these two processes. If more energy is released when making new bonds than was absorbed to break the old ones, the overall reaction is exothermic. If less energy is released making bonds than was absorbed to break them, the overall reaction is endothermic.

Key term

Bond Energy: The amount of energy required to break one mole of a specific covalent bond, measured in kilojoules per mole (kJ/mol).

Examiner insight

To get full marks on explanation questions, you must explicitly mention both bond breaking (requiring energy) and bond making (releasing energy).

Common pitfall

The most common mistake is thinking that 'strong bonds store a lot of energy'. In fact, strong bonds are stable and have low potential energy; it takes a lot of energy to break them, and a lot of energy is released when they are formed.

Worked example 13 marks

Explain, in terms of bond breaking and bond making, why the combustion of hydrogen (2H₂ + O₂ → 2H₂O) is a strongly exothermic reaction.

  1. 1
    1. State the principle for reactants: Energy must be supplied to break the H-H bonds in the hydrogen molecules and the O=O double bond in the oxygen molecules.
  2. 2
    1. State the principle for products: Energy is released when new O-H bonds are formed in the water molecules.
  3. 3
    1. Compare the energy changes: In this reaction, the amount of energy released when forming the strong O-H bonds in two moles of water is much greater than the amount of energy absorbed to break the bonds in the hydrogen and oxygen molecules.
  4. 4
    1. Conclude: Because more energy is released than absorbed, the reaction is exothermic overall.

Recap

  • Energy is ALWAYS required to break chemical bonds (endothermic).
  • Energy is ALWAYS released when chemical bonds are formed (exothermic).
  • An exothermic reaction releases more energy making bonds than it absorbs breaking them.
  • An endothermic reaction absorbs more energy breaking bonds than it releases making them.

Quick check

  1. Is the process of forming a chemical bond endothermic or exothermic?1 mark

4. Calculating Reaction Energy

For the Extended syllabus, you can calculate the overall enthalpy change (ΔH) for a reaction using average bond energies. You will be given a table of bond energy values. The calculation is a two-step energy balance:

  1. Calculate the total energy absorbed to break all the bonds in the reactant molecules. This is the 'energy in'. Be sure to multiply the bond energy by the number of that specific bond in the balanced equation.
  2. Calculate the total energy released when forming all the bonds in the product molecules. This is the 'energy out'.
  3. The overall enthalpy change is the difference: ΔH = Energy In - Energy Out.

ΔH = Σ (bond energies of bonds broken) - Σ (bond energies of bonds formed)

Examiner insight

Show your working clearly by listing the bonds being broken and formed separately. Marks are often given for correctly identifying the bonds and for the 'energy in' and 'energy out' totals, even if the final subtraction is incorrect.

Common pitfall

Forgetting to multiply bond energies by the stoichiometric coefficients in the balanced equation. For example, in 2H₂O, there are four O-H bonds to be formed, not two.

Worked example 13 marks

Calculate the enthalpy change for the reaction between hydrogen and chlorine: H₂(g) + Cl₂(g) → 2HCl(g). Use the following bond energies (in kJ/mol): H-H = 436, Cl-Cl = 242, H-Cl = 431.

  1. 1

    Step 1: Calculate energy IN to break bonds in reactants.

  2. 2

    Bonds broken: 1 × (H-H) and 1 × (Cl-Cl)

  3. 3

    Energy in = (1 × 436) + (1 × 242) = 678 kJ

  4. 4

    Step 2: Calculate energy OUT from forming bonds in products.

  5. 5

    Bonds formed: 2 × (H-Cl) because the product is 2HCl.

  6. 6

    Energy out = 2 × 431 = 862 kJ

  7. 7

    Step 3: Calculate the overall enthalpy change (ΔH).

  8. 8

    ΔH = Energy in - Energy out = 678 - 862 = -184 kJ/mol

  9. 9

    The reaction is exothermic as the sign is negative.

Recap

  • Sum the energies of all bonds broken in the reactants (energy in).
  • Sum the energies of all bonds formed in the products (energy out).
  • Calculate ΔH by subtracting the energy out from the energy in.
  • Always pay close attention to the balancing numbers (stoichiometry) in the equation.
  • A negative result for ΔH means the reaction is exothermic; a positive result means it is endothermic.

Quick check

  1. For the reaction N₂ + 3H₂ → 2NH₃, how many N-H bonds are formed in total?1 mark

5. Fuels and Combustion

A fuel is a substance that stores chemical energy, which can be released easily, usually as heat, through a process called combustion. Combustion is a rapid reaction with oxygen that is always exothermic. The energy released can be harnessed for heating, cooking, generating electricity, or powering engines. The most common fuels are fossil fuels: coal (solid), petroleum/crude oil (liquid, from which we get petrol, diesel, and kerosene), and natural gas (mainly methane). These were formed from the remains of ancient organisms over millions of years and are non-renewable. Another vital exothermic process is aerobic respiration in our bodies, which is like a slow, controlled combustion of glucose to release the energy we need to live: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy.

Hydrocarbon + Oxygen → Carbon Dioxide + Water

Key term

Combustion: A high-temperature exothermic reaction between a fuel and an oxidant, usually oxygen, to produce heat and light.

Fun fact

The energy released from burning just 1 kg of natural gas is enough to power an average UK home for about 2 hours.

Worked example 13 marks

Propane (C₃H₈) is a fuel used in patio heaters and for cooking. Write a balanced chemical equation for the complete combustion of propane.

  1. 1
    1. Write the reactants and products: C₃H₈ + O₂ → CO₂ + H₂O
  2. 2
    1. Balance the Carbon atoms: There are 3 C atoms on the left, so we need 3 CO₂ on the right. C₃H₈ + O₂ → 3CO₂ + H₂O
  3. 3
    1. Balance the Hydrogen atoms: There are 8 H atoms on the left, so we need 4 H₂O on the right (4 × 2 = 8). C₃H₈ + O₂ → 3CO₂ + 4H₂O
  4. 4
    1. Balance the Oxygen atoms: On the right, there are (3 × 2) + (4 × 1) = 10 O atoms. So we need 5 O₂ on the left.
  5. 5
    1. Final balanced equation: C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l)

Recap

  • A fuel is a substance burned to release useful heat energy.
  • Combustion is the process of burning a fuel in oxygen, and it is always exothermic.
  • The three main fossil fuels are coal, petroleum (oil), and natural gas.
  • Complete combustion of hydrocarbons produces carbon dioxide and water.
  • Respiration is the body's exothermic process for releasing energy from glucose.

Quick check

  1. Name the two products of the complete combustion of any hydrocarbon fuel.2 marks

6. Reversible Reactions and Energy

Some chemical reactions are reversible, meaning the products can react to re-form the original reactants. We show this using a double arrow (⇌) in the chemical equation. Energy changes in reversible reactions are directly linked. If the forward reaction is exothermic, the reverse reaction must be endothermic, and it will absorb the exact same amount of energy that the forward reaction released. The classic example is hydrated copper(II) sulfate. When you heat blue hydrated copper(II) sulfate crystals, they turn into white anhydrous copper(II) sulfate powder and release water vapour. This is an endothermic reaction (it needs continuous heating). CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(g). If you then add water back to the white powder, it turns blue again and releases heat – the reverse reaction is exothermic.

CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)

Forward reaction (heating): Endothermic

Reverse reaction (adding water): Exothermic

Key term

Reversible Reaction: A reaction in which the products can react to re-form the reactants, so the reaction can proceed in both the forward and reverse directions.

Examiner insight

Students must know the colours of hydrated copper(II) sulfate (blue) and anhydrous copper(II) sulfate (white) and link the forward reaction (heating) to an endothermic change.

Worked example 14 marks

A student gently heats some blue hydrated copper(II) sulfate crystals in a test tube.(a) What two things would they observe?(b) How could they then show the reaction is reversible?

  1. 1

    (a) Observation 1: The blue solid crystals will turn into a white powder.

  2. 2

    Observation 2: Colourless liquid (water) or steam will be seen condensing at the cooler, upper part of the test tube.

  3. 3

    (b) To show it is reversible, the student should allow the test tube to cool and then carefully add a few drops of water to the white powder. They would observe the powder turning blue again and feel the test tube getting warm.

Recap

  • A reversible reaction can proceed in both forward and backward directions, shown by the ⇌ symbol.
  • If the forward reaction is exothermic (ΔH is negative), the reverse reaction is endothermic (ΔH is positive).
  • The magnitude of the enthalpy change is the same for both forward and reverse reactions, but the sign is opposite.
  • Heating blue hydrated copper(II) sulfate to form the white anhydrous salt is a key example of an endothermic reversible reaction.
  • Anhydrous copper(II) sulfate is used as a chemical test for water, turning from white to blue in its presence.

Quick check

  1. The Haber process, N₂ + 3H₂ ⇌ 2NH₃, is exothermic. What can you say about the energy change when ammonia decomposes back into nitrogen and hydrogen?1 mark

End-of-chapter exercise

Test yourself on the whole chapter. Work through these before moving on.

  1. Define an endothermic reaction and give one specific example.2 marks
  2. Draw a fully labelled energy level diagram for a reaction where the overall enthalpy change, ΔH, is +45 kJ/mol. Your diagram should include labels for axes, reactants, products, ΔH, and activation energy (Ea).4 marks
  3. When solid ammonium nitrate is dissolved in water, the temperature of the solution drops significantly. Explain this observation in terms of energy transfer.2 marks
  4. Explain, using the concepts of bond breaking and bond making, why the overall enthalpy change of some reactions is positive.3 marks
  5. The reaction to form hydrogen iodide from its elements is: H₂(g) + I₂(g) → 2HI(g). Use the bond energies provided to calculate the overall enthalpy change for this reaction. Bond energies (kJ/mol): H-H = 436, I-I = 151, H-I = 299.3 marks
  6. Write the balanced chemical equation for the complete combustion of ethanol (C₂H₅OH).2 marks
  7. The thermal decomposition of hydrated cobalt(II) chloride is a reversible reaction shown by the equation: CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g). The forward reaction is endothermic. State the expected colour of hydrated cobalt(II) chloride and anhydrous cobalt(II) chloride, and predict what would happen to the temperature if you added water to the anhydrous salt.3 marks
  8. A reaction has an activation energy of +70 kJ/mol. The overall enthalpy change for the reaction is -110 kJ/mol. What is the activation energy for the reverse reaction?2 marks
  9. Both respiration and the combustion of fossil fuels are exothermic processes. State one similarity and one difference between these two processes, other than their location.2 marks
  10. The combustion of hydrazine, N₂H₄, is highly exothermic, and it is used as a rocket fuel. The equation is: N₂H₄(l) + O₂(g) → N₂(g) + 2H₂O(g). Explain in detail, with reference to the relative strengths of the bonds being broken and formed, why this reaction releases a large amount of energy.4 marks

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