Exothermic and Endothermic Reactions Revision Notes for Leaving Cert Chemistry

Exothermic and Endothermic Reactions

Heat Changes in Chemical Reactions

Chemical reactions often result in energy changes, typically observed as changes in temperature. These changes can be either exothermic or endothermic depending on whether heat is released or absorbed during the reaction.

Exothermic Reactions

Definition: In an exothermic reaction, energy is released to the surroundings, usually in the form of heat. This causes the temperature of the surroundings to increase.
Sign of $∆H$ : Negative ( $\Delta H < 0$ ) because heat is lost by the system.

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Examples: Combustion of hydrocarbons:

The combustion of alkanes, such as the burning of methane, is highly exothermic.

CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O \, (\Delta H = -890 \, \text{kJ/mol})

Neutralization reactions: The reaction between acids and bases releases heat.

Demonstration:

Burning of Magnesium: When magnesium burns in oxygen, it produces an intense exothermic reaction that releases heat and light.

Endothermic Reactions

Definition: In an endothermic reaction, energy is absorbed from the surroundings, causing the temperature of the surroundings to decrease.
Sign of $∆H$ : Positive ( $\Delta H > 0$ ) because heat is taken in by the system.

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Examples: Thermal decomposition:

The decomposition of calcium carbonate absorbs heat:

CaCO_3 \rightarrow CaO + CO_2 \, (\Delta H > 0)

Photosynthesis: Plants absorb sunlight to convert carbon dioxide and water into glucose.

Demonstration:

Dissolution of Ammonium Nitrate: When ammonium nitrate dissolves in water, the solution feels cold, indicating an endothermic process.

Heat of Reaction

The heat of reaction refers to the amount of heat energy absorbed or released during a chemical reaction. It depends on the number of moles of reactants reacting as indicated by the balanced chemical equation.

Heat of Combustion

Definition: The heat of combustion is the heat released when one mole of a substance is completely burned in excess oxygen.
Measurement: The bomb calorimeter is an instrument used to measure the heat of combustion accurately.

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Example of Heat of Combustion For methane:

CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O \, (\Delta H = -890 \, \text{kJ/mol})

This value represents the amount of heat released when one mole of methane combusts completely.

Bond Energy

Bond energy is the amount of energy required to break one mole of a particular covalent bond in a gaseous molecule.

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Example: $C–H$ Bond in Methane The energy required to break a single $C–H$ bond in methane is approximately 412 kJ/mol.

Breaking all four C–H bonds in one mole of methane requires $4 × 412 = 1,648 kJ$ .

Law of Conservation of Energy

This law states that energy cannot be created or destroyed, only transferred or transformed from one form to another.

Hess's Law

Hess's law states that the total enthalpy change of a reaction is the same, regardless of the pathway taken. This means that the energy change in a reaction can be calculated by adding the energy changes of intermediate steps.

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Example: The enthalpy change for the formation of carbon dioxide from carbon can be calculated using intermediate steps such as the formation of carbon monoxide.

Simple Calculations Using Heats of Formation

The heat of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states.

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Example If the heat of formation of water is $\Delta H_f = -286 \, \text{kJ/mol}$

Then the enthalpy change for the formation of 2 moles of water is:

2 \times (-286) = -572 \, \text{kJ}

Heat of Combustion and Bomb Calorimeter

The bomb calorimeter is used to measure the calorific value of fuels or food. It determines the energy released during combustion in a controlled, insulated environment.

Kilogram calorific value: This refers to the amount of energy released when 1 kg of fuel is burned.
Applications: Determining the calorific value of fuels (such as gasoline) and foods (like carbohydrates) is essential in industries like energy production and nutrition.

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Exam Tip:

Know how to differentiate between exothermic and endothermic reactions.
Be familiar with the concept of bond energy and how it relates to heat changes in chemical reactions.
Understand the practical uses of a bomb calorimeter and how it measures the energy content of fuels and foods.
Use Hess's law to simplify the calculation of enthalpy changes.

Exothermic and Endothermic Reactions (Leaving Cert Chemistry): Revision Notes