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CIE iGCSE Co-ordinated Sciences-C5.1 Exothermic and endothermic reactions- Study Notes- New Syllabus

CIE iGCSE Co-ordinated Sciences-C5.1 Exothermic and endothermic reactions – Study Notes

CIE iGCSE Co-ordinated Sciences-C5.1 Exothermic and endothermic reactions – Study Notes -CIE iGCSE Co-ordinated Sciences – per latest Syllabus.

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CIE iGCSE Co-Ordinated Sciences-Concise Summary Notes- All Topics

Exothermic Reactions

An exothermic reaction is a chemical reaction that releases thermal energy to its surroundings. This transfer of energy usually results in an increase in the temperature of the surroundings.

  • Energy is released because the total energy of the products is less than the total energy of the reactants.
  • The excess energy is given off as heat, light, or both.
  • Common examples include combustion reactions, neutralisation reactions, and many oxidation reactions.

Example 

When 50 g of magnesium reacts with oxygen to form magnesium oxide, the temperature of the surroundings increases. Explain why this is an exothermic reaction.

▶️ Answer/Explanation
  • Reaction: \( 2\text{Mg} + \text{O2} \rightarrow 2\text{MgO} \)
  • Magnesium oxide has lower chemical potential energy than magnesium and oxygen.
  • The excess energy is released as heat and light to the surroundings.
  • Observation: Increase in temperature and bright white flame, indicating an exothermic reaction.

Endothermic Reactions

An endothermic reaction is a chemical reaction that absorbs thermal energy from its surroundings. This uptake of energy usually results in a decrease in the temperature of the surroundings.

  • Energy is absorbed because the total energy of the products is higher than the total energy of the reactants.
  • The required energy is taken from the surroundings as heat, leading to cooling.
  • Common examples include thermal decomposition reactions and photosynthesis.

Example 

When solid ammonium chloride dissolves in water, the temperature of the solution decreases. Explain why this is an endothermic reaction.

▶️ Answer/Explanation
  • Reaction: \( \text{NH4Cl (s)} \rightarrow \text{NH4}^+ (aq) + \text{Cl}^- (aq) \)
  • The dissolution process absorbs thermal energy from the water.
  • Observation: Temperature of the solution drops, indicating energy is taken from the surroundings.

Reaction Pathway Diagrams: Exothermic and Endothermic Reactions

Reaction pathway diagrams (also called energy profile diagrams) show the progress of a chemical reaction and the energy changes that occur from reactants to products.

Exothermic reactions:

  • The energy of the products is lower than the energy of the reactants.
  • Energy is released to the surroundings, often as heat or light.
  • The diagram shows a downward slope from reactants to products.
  • Activation energy is the energy peak that must be overcome for the reaction to proceed.

Endothermic reactions:

  • The energy of the products is higher than the energy of the reactants.
  • Energy is absorbed from the surroundings to allow the reaction to occur.
  • The diagram shows an upward slope from reactants to products.
  • Activation energy is still required to initiate the reaction.

Example 

Explain how you can identify whether a reaction is exothermic or endothermic by looking at its reaction pathway diagram.

▶️ Answer/Explanation
  • Compare the energy levels of reactants and products.
  • If the products are lower in energy than the reactants, the reaction is exothermic; energy is released to the surroundings.
  • If the products are higher in energy than the reactants, the reaction is endothermic; energy is absorbed from the surroundings.
  • The activation energy is the peak of the curve, representing the minimum energy required to start the reaction.

Enthalpy Change (\( \Delta H \))

The transfer of thermal energy during a chemical reaction is called the enthalpy change, \( \Delta H \), of the reaction. It represents the difference in energy between reactants and products. Understanding enthalpy change helps predict whether a reaction will release or absorb energy, and it is crucial in practical applications such as fuels, batteries, and industrial processes.

Exothermic reactions:

  • Energy is released to the surroundings, usually as heat or light.
  • \( \Delta H \) is negative (\( \Delta H < 0 \)).
  • The products have lower energy than the reactants.
  • Common examples: combustion of fuels, neutralisation reactions, respiration.
  • Observation: Temperature of the surroundings increases.
  • Applications: Heating systems, hand warmers, combustion engines.

Endothermic reactions:

  • Energy is absorbed from the surroundings, usually as heat.
  • \( \Delta H \) is positive (\( \Delta H > 0 \)).
  • The products have higher energy than the reactants.
  • Common examples: thermal decomposition, photosynthesis, dissolving ammonium chloride in water.
  • Observation: Temperature of the surroundings decreases.
  • Applications: Instant cold packs, photosynthesis in plants, certain industrial processes.

Additional Points:

  • Enthalpy change depends on the type and number of bonds broken and formed.
  • Energy profile diagrams can show the relative energy levels of reactants and products and the activation energy.
  • Measuring \( \Delta H \) helps chemists design reactions efficiently and safely.

Example 

A reaction releases 250 kJ of energy to the surroundings. State the sign of \( \Delta H \), identify whether the reaction is exothermic or endothermic, and explain what happens to the temperature of the surroundings.

▶️ Answer/Explanation
  • Energy is released → heat flows to the surroundings.
  • \( \Delta H \) is negative: \( \Delta H = -250\,\text{kJ} \).
  • The reaction is exothermic.
  • The temperature of the surroundings increases due to the released energy.

Activation Energy (\( E_a \))

Activation energy, \( E_a \), is the minimum energy that colliding particles must have for a chemical reaction to occur. Without this energy, even if particles collide, the reaction will not proceed. It is an essential concept in understanding reaction rates and energy profile diagrams.

  • Represents the energy barrier between reactants and products.
  • Determines how quickly a reaction occurs; higher \( E_a \) → slower reaction at a given temperature.
  • Can be lowered by a catalyst, which provides an alternative pathway with a lower activation energy.
  • Shown as the peak of a curve in a reaction pathway or energy profile diagram.

Example 

In a reaction, the reactants have an energy of 50 kJ/mol, and the peak of the reaction pathway diagram is at 120 kJ/mol. Calculate the activation energy and explain its significance.

▶️ Answer/Explanation
  • Activation energy: \( E_a = 120 – 50 = 70\,\text{kJ/mol} \)
  • This is the minimum energy the particles must have to overcome the energy barrier and react.
  • Particles with energy less than \( E_a \) will collide but will not form products.

Reaction Pathway Diagrams: (a) Reactants

In a reaction pathway diagram, the reactants are the substances present at the start of a chemical reaction. They are shown on the left-hand side of the diagram and their position on the y-axis represents their relative energy.

  • For both exothermic and endothermic reactions, reactants appear at the beginning of the energy curve.
  • Their energy level serves as a reference point to compare with the products.

Example 

On a reaction pathway diagram for the combustion of methane, label the reactants. The reaction is: \( \text{CH}_{4} + 2\text{O}_{2} \;\rightarrow\; \text{CO}_{2} + 2\text{H}_{2}\text{O} \).

▶️ Answer/Explanation
  • Reactants: \( \text{CH}_{4} \) and \( \text{O}_{2} \)
  • They are positioned at the start of the energy curve (left-hand side).
  • This shows the initial energy of the system before the reaction begins.

Reaction Pathway Diagrams: (b) Products

In a reaction pathway diagram, the products are the substances formed at the end of the reaction. They are shown on the right-hand side of the diagram, and their position on the y-axis represents their relative energy compared to the reactants.

  • For an exothermic reaction, the products have lower energy than the reactants because energy is released to the surroundings.
  • For an endothermic reaction, the products have higher energy than the reactants because energy is absorbed from the surroundings.

Example 

On a reaction pathway diagram for the combustion of methane, label the products. The reaction is: \( \text{CH}_{4} + 2\text{O}_{2} \;\rightarrow\; \text{CO}_{2} + 2\text{H}_{2}\text{O} \).

▶️ Answer/Explanation
  • Products: \( \text{CO}_{2} \) and \( \text{H}_{2}\text{O} \)
  • They are positioned at the end of the energy curve (right-hand side).
  • For this exothermic reaction, the products have lower energy than the reactants, indicating energy is released to the surroundings.

Reaction Pathway Diagrams: (c) Overall Energy Change (\( \Delta H \))

The overall energy change of a reaction, \( \Delta H \), is the difference in energy between the reactants and the products. It indicates whether a reaction is exothermic or endothermic.

  • For an exothermic reaction, \( \Delta H \) is negative (\( \Delta H < 0 \)) because the products have lower energy than the reactants. Energy is released to the surroundings.
  • For an endothermic reaction, \( \Delta H \) is positive (\( \Delta H > 0 \)) because the products have higher energy than the reactants. Energy is absorbed from the surroundings.
  • The magnitude of \( \Delta H \) is represented by the vertical difference between reactants and products on the energy axis of the reaction pathway diagram.

Example 

Using a reaction pathway diagram for the combustion of methane (\( \text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} \)), determine the overall energy change (\( \Delta H \)) and explain its sign.

▶️ Answer/Explanation
  • Reactants: \( \text{CH}_4 \) and \( \text{O}_2 \); Products: \( \text{CO}_2 \) and \( \text{H}_2\text{O} \)
  • Products have lower energy than reactants → energy released.
  • Overall energy change: \( \Delta H < 0 \), indicating an exothermic reaction.
  • The vertical difference between reactants and products on the diagram shows the magnitude of \( \Delta H \).

Reaction Pathway Diagrams: (d) Activation Energy (\( E_a \))

Activation energy, \( E_a \), is the minimum energy that reacting particles must have for a reaction to occur. On a reaction pathway diagram, it is represented as the energy difference between the reactants and the peak of the curve (the transition state).

  • For both exothermic and endothermic reactions, the curve rises from the reactants to a maximum point representing \( E_a \).
  • After reaching the peak, the curve falls to the products’ energy level.
  • The larger the activation energy, the slower the reaction occurs at a given temperature.
  • Catalysts work by providing an alternative pathway with a lower \( E_a \), increasing the reaction rate.

Example 

For the combustion of methane (\( \text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} \)), the reactants have an energy of 50 kJ/mol, and the peak of the reaction pathway is at 120 kJ/mol. Calculate the activation energy (\( E_a \)) and explain its significance.

▶️ Answer/Explanation
  • Activation energy: \( E_a = 120 – 50 = 70\,\text{kJ/mol} \)
  • This is the minimum energy required for methane and oxygen molecules to collide successfully and react.
  • Particles with energy less than \( E_a \) will collide but will not form products.

Bond Energy: Breaking and Making Bonds

In chemical reactions, energy is involved in breaking and forming bonds. This is an important concept for understanding why reactions are exothermic or endothermic.

Bond breaking: Energy is required to break chemical bonds between atoms. Therefore, bond breaking is an endothermic process and absorbs energy from the surroundings.

Bond making: Energy is released when new chemical bonds are formed. Therefore, bond making is an exothermic process and releases energy to the surroundings.

The overall energy change of a reaction (\( \Delta H \)) depends on the difference between the energy absorbed to break bonds and the energy released when new bonds are formed:

  • If more energy is released in bond formation than absorbed in bond breaking → reaction is exothermic (\( \Delta H < 0 \)).
  • If more energy is absorbed in bond breaking than released in bond formation → reaction is endothermic (\( \Delta H > 0 \)).

Example 

During the reaction \( \text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl} \), explain why bond breaking and bond making affect the overall energy change.

▶️ Answer/Explanation
  • Energy is required to break the H-H and Cl-Cl bonds → endothermic process.
  • Energy is released when two H-Cl bonds are formed → exothermic process.
  • If the energy released in forming H-Cl bonds is greater than the energy absorbed in breaking H-H and Cl-Cl bonds, the overall reaction is exothermic.
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