Energetics, or thermochemistry, is the study of energy changes that accompany chemical and physical processes. All chemical reactions involve energy changes, either absorbing energy from the surroundings (endothermic) or releasing energy to the surroundings (exothermic). This energy is typically exchanged as heat and is quantified as an enthalpy change (ΔH).

Enthalpy (H) is a thermodynamic property that represents the total heat content of a system at constant pressure. It is a state function, meaning its value depends only on the initial and final states of the system, not on the path taken.

Enthalpy change (ΔH) is the heat absorbed or released during a chemical reaction at constant pressure.

● Exothermic reactions: Release heat to the surroundings. The enthalpy of the products is lower than the enthalpy of the reactants, so ΔH is negative (ΔH < 0). Examples include combustion reactions and neutralization reactions.

● Endothermic reactions: Absorb heat from the surroundings. The enthalpy of the products is higher than the enthalpy of the reactants, so ΔH is positive (ΔH > 0). Examples include melting ice and photosynthesis.

Enthalpy changes are typically measured under standard conditions to allow for comparison. Standard conditions are defined as:
● Pressure: 100 kPa (1 atm)
● Temperature: 298 K (25 °C)
● Concentration: 1 mol dm⁻³ for solutions
A superscript circle (°) is used to denote standard conditions (e.g., ΔH°).

Let's examine some specific types of standard enthalpy changes:

Standard Enthalpy of Formation (ΔH_f°): The standard enthalpy of formation of a compound is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.

● By definition, the standard enthalpy of formation of an element in its most stable form under standard conditions is zero. For example, ΔH_f°(O₂(g)) = 0, ΔH_f°(C(graphite)) = 0.
● Example: C(graphite) + O₂(g) → CO₂(g) ΔH_f° = -393.5 kJ mol⁻¹ (for CO₂)

Standard Enthalpy of Combustion (ΔH_c°): The standard enthalpy of combustion of a substance is the enthalpy change when one mole of a substance undergoes complete combustion in excess oxygen under standard conditions.

● Combustion reactions are always exothermic, so ΔH_c° values are always negative.
● Example: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH_c° = -890.3 kJ mol⁻¹ (for CH₄)

Standard Enthalpy of Neutralization (ΔH_neut°): The standard enthalpy of neutralization is the enthalpy change when one mole of water is formed from the reaction of an acid and a base under standard conditions. For strong acid-strong base reactions, the enthalpy of neutralization is remarkably consistent, approximately -57.3 kJ mol⁻¹, because the net ionic equation is always the same: H⁺(aq) + OH⁻(aq) → H₂O(l)

● Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) ΔH_neut° ≈ -57.3 kJ mol⁻¹

Enthalpy changes are measured experimentally using a calorimeter. A simple calorimeter can be a polystyrene cup or a bomb calorimeter for more accurate measurements. The principle involves measuring the temperature change (ΔT) of a known mass of water (or solution) due to the heat released or absorbed by a reaction.
The heat exchanged (q) is calculated using the formula: q = mcΔT
Where:
● q = heat energy transferred (Joules, J)
● m = mass of the substance (water/solution) that changes temperature (grams, g)
● c = specific heat capacity of the substance (for water, 4.18 J g⁻¹ K⁻¹ or J g⁻¹ °C⁻¹)
● ΔT = change in temperature (K or °C)
Once 'q' is determined, the enthalpy change (ΔH) for the reaction can be calculated by dividing 'q' by the number of moles of reactant that caused that heat change. Remember to consider the sign convention: if the solution temperature increases (exothermic reaction), q is positive for the surroundings, but ΔH for the reaction is negative. If the solution temperature decreases (endothermic reaction), q is negative for the surroundings, and ΔH for the reaction is positive.