3.1 - Heat Change in Reactions

Exothermic and Endothermic Reactions

Chemical reactions normally involve heat energy being released or absorbed.

Exothermic reaction	Endothermic reaction
Releases heat to the surroundings. The surrounding temperature rises and the container becomes hot.	Absorbs heat from the surroundings. The surrounding temperature falls and the container becomes cold.
Examples: respiration, oxidation of metals, neutralisation, combustion of fuel, production of ammonia and dissolving sodium hydroxide in water.	Examples: photosynthesis, thermal decomposition of metal carbonates or metal nitrates, and dissolving ammonium salts in water.

Alcohol evaporation in hand sanitiser absorbs heat, making the hand feel cold. Detergent powder in contact with water releases heat, making the hand feel warm.

Heat Change When Substances Dissolve

Heat changes during dissolution can be identified from the initial temperature and the highest or lowest temperature of the mixture.
The comparison uses $20 c m^{3}$ of distilled water and one spatula of solid sodium hydroxide, $NaOH$ , ammonium nitrate, $N H_{4} N O_{3}$ , sodium thiosulphate, $N a_{2} S_{2} O_{3}$ , or anhydrous calcium chloride, $CaC l_{2}$ .
A lidded polystyrene cup, thermometer, measuring cylinder and spatula are used; the water is left with the thermometer for two minutes before its initial temperature is taken, and the mixture is stirred carefully.
A temperature rise indicates heat release; a temperature fall indicates heat absorption.
Precaution: Sodium hydroxide is corrosive and requires careful handling.

Heat of Reaction and Energy Level Diagrams

Heat of reaction, $Δ H$ , is the heat change when one mole of reactant reacts or one mole of product forms. Its unit is $kJ mo l^{- 1}$ .
$Δ H$ is negative when heat is released; the reaction is exothermic.
$Δ H$ is positive when heat is absorbed; the reaction is endothermic.
An energy level diagram shows the difference between the heat-energy content of the reactants and products:

Δ H = H_{products} - H_{reactants}

Feature	Exothermic	Endothermic
Relative energy level	Products lower than reactants	Products higher than reactants
Direction of $Δ H$ arrow	Downwards, from reactants to products	Upwards, from reactants to products
Sign of $Δ H$	Negative	Positive
Temperature of mixture	Increases	Decreases

Exothermic example:

Mg (s) + H_{2} S O_{4} (aq) \to MgS O_{4} (aq) + H_{2} (g) Δ H = - 467 kJ mo l^{- 1}

The products have lower total energy than the reactants; $467 kJ$ is released when the equation reacts in the stated molar amounts.

Energy axis with magnesium and sulphuric acid above magnesium sulphate and hydrogen, and a downward arrow labelled ΔH = −467 kJ mol⁻¹. — Energy level diagram for an exothermic reaction

Endothermic example:

N_{2} (g) + O_{2} (g) \to 2NO (g) Δ H = + 180 kJ mo l^{- 1}

The product has higher total energy than the reactants; $180 kJ$ is absorbed when the equation reacts in the stated molar amounts.

Energy axis with nitrogen and oxygen below nitrogen monoxide, and an upward arrow labelled ΔH = +180 kJ mol⁻¹. — Energy level diagram for an endothermic reaction

To construct an energy level diagram, label the vertical axis Energy, draw and label the reactant and product levels, draw the $Δ H$ arrow from the reactant level to the product level, and write its sign, value and unit.

Energy Change During Bond Breaking and Bond Formation

Breaking bonds in reactants absorbs heat energy.
Forming new bonds in products releases heat energy.
An exothermic reaction releases more energy during bond formation than it absorbs during bond breaking.
An endothermic reaction absorbs more energy during bond breaking than it releases during bond formation.

For $H_{2} (g) + C l_{2} (g) \to 2HCl (g)$ :

Bond process	Energy change
Break one H–H bond and one Cl–Cl bond	Absorb $436 + 243 = 679 kJ mo l^{- 1}$
Form two H–Cl bonds	Release $2 (432) = 864 kJ mo l^{- 1}$

The net heat change is $679 - 864 = - 185 kJ mo l^{- 1}$ , so the reaction is exothermic.