1.6.2 Enthalpy Changes

What is Enthalpy Change ($ΔH$)?

Types of Reactions

• Endothermic reactions: Energy is absorbed from the surroundings, making $ΔH$ positive. The products have more energy than the reactants.
• Exothermic reactions: Energy is released to the surroundings, making $ΔH$ negative. The products have less energy than the reactants.

Formula for Enthalpy Change

$$ΔH = \text{Energy of products} - \text{Energy of reactants}$$

If the value of $ΔH$ is negative, the reaction is exothermic. If it is positive, the reaction is endothermic.

Standard Enthalpy Changes

Enthalpy changes are often measured under standard conditions:

• Pressure of 100 kPa (approximately 1 atmosphere).
• Temperature of 298 K (25°C).

• All reactants and products must be in their standard states (the most stable form of a substance under standard conditions). Standard enthalpy changes are represented with a superscript $θ$ (e.g $ΔH^\circ$)

Standard Enthalpy Change of Formation ($ΔfH°$)

The standard enthalpy change of formation is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.

The enthalpy change of formation for elements in their standard states is zero.

For example, the enthalpy of formation of $O_2(g)$ is zero.

Standard Enthalpy Change of Combustion ($ΔcH°$)

The standard enthalpy change of combustion is the enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions, with all substances in their standard states.

Importance of Enthalpy in Reactions

• State Symbols: When writing equations involving enthalpy changes, always include state symbols (e.g, $s$, $l$, $g$) to represent the physical states of the reactants and products.
• Exothermic vs Endothermic: Knowing the enthalpy change allows you to predict whether a reaction releases heat (exothermic) or absorbs heat (endothermic).
• Standard Conditions: These conditions ensure that enthalpy values are comparable across different reactions.