Enthalpy Change, ΔH

• Exothermic Reactions occur when the enthalpy of the products is lower than that of the reactants, resulting in a negative $\Delta H$ value.

• In exothermic processes, energy is released to the surroundings, typically causing an increase in the temperature of the environment.

• Endothermic Reactions occur when the products have higher enthalpy than the reactants, requiring an input of energy and resulting in a positive $\Delta H$ value.

• In endothermic processes, heat is absorbed from the surroundings, leading to a decrease in the temperature of the environment.

• Standard Enthalpy of Reaction ($\Delta H^{\theta}_{r}$): The enthalpy change when reactants react in stoichiometric amounts under standard conditions.

• Standard Enthalpy of Formation ($\Delta H^{\theta}_{f}$): The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states.

• Standard Enthalpy of Combustion ($\Delta H^{\theta}_{c}$): The enthalpy change when one mole of a substance is completely burned in excess oxygen; this is always an exothermic process.

• Standard Enthalpy of Neutralisation ($\Delta H^{\theta}_{neut}$): The enthalpy change when an acid and alkali react to form one mole of water.

Calorimetry

• This experimental technique measures the heat transfer ($q$) to a known mass of surroundings (usually water) using the formula:

$q = m \cdot c \cdot \Delta T$

• Where $m$ is the mass of the liquid, $c$ is the specific heat capacity (e.g., $4.18\text{ J g}^{-1}\text{ K}^{-1}$ for water), and $\Delta T$ is the temperature change.

Bond Energies

• Enthalpy change can be estimated by calculating the energy required to break bonds and the energy released when new bonds form.
• Bond breaking is always endothermic (requires energy), while bond forming is always exothermic (releases energy).
• Formula: $\Delta H = \sum(\text{bond energies of reactants}) - \sum(\text{bond energies of products})$.

• It is critical to distinguish between the system (the chemical reaction) and the surroundings (the water in a calorimeter or the air).

• In an exothermic reaction, the system loses energy to the surroundings, so the system's enthalpy decreases while the surroundings' temperature rises.

• Check the Units: Calorimetry calculations ($q = mc\Delta T$) yield results in Joules (J). You must divide by $1000$ to convert to kilojoules (kJ) before calculating $\Delta H$ in $\text{kJ mol}^{-1}$.

• Sign Consistency: Always include the $+$ or $-$ sign in your final answer for $\Delta H$. Omitting the sign is a frequent cause of lost marks.

• Definition Precision: When defining $\Delta H^{\theta}_{f}$, ensure you specify 'one mole of compound' and 'from its elements in their standard states'.

• Stoichiometry: If a reaction produces $2$ moles of a product, the total energy change is $2 \times \Delta H^{\theta}_{f}$ for that substance.