Thermochemistry
Heat, Enthalpy & Energy

Thermochemistry is the branch of physical chemistry that studies heat changes accompanying chemical reactions and physical transformations. When substances react, their chemical bonds rearrange and their energy changes. Depending on the process, energy may be transferred as heat between the system and its surroundings. Understanding thermochemistry helps explain combustion, neutralisation, dissolution, melting, freezing and many industrial processes.

In thermochemistry, the system is the substance or reaction being studied, while the surroundings include everything outside that system. Heat flows because of a temperature difference. The direction of heat transfer and the sign of the enthalpy change help classify a process as endothermic or exothermic.

The first law of thermodynamics expresses conservation of energy. Using the chemistry sign convention, where heat added to the system and work done on the system are positive, the relationship is ΔU = q + w. Internal energy, U, is a state function. Its change depends on the initial and final states rather than the route taken between them.

Enthalpy, represented by H, is defined as H = U + PV. At constant pressure, when only pressure–volume work is involved, the heat transferred to the system equals its enthalpy change: ΔH = q_p. This makes enthalpy particularly useful for studying reactions that take place in open containers under approximately constant atmospheric pressure.

Calorimetry is an experimental method for measuring heat transferred during a chemical reaction or physical change. A calorimeter helps measure the temperature change of a known quantity of material. By combining that temperature change with the mass and specific heat capacity, the heat absorbed by the material can be calculated.

For a substance that experiences a moderate temperature change without a phase change, the basic calorimetry equation is q = mcΔT. Here, q is heat transferred, m is mass, c is specific heat capacity and ΔT is the final temperature minus the initial temperature. If mass is measured in grams, specific heat capacity in J g⁻¹ K⁻¹ and temperature in kelvin or degrees Celsius difference, heat is obtained in joules.

In a coffee-cup calorimeter, a reaction takes place in an insulated container at approximately constant atmospheric pressure. The measured heat change of the solution can be used to estimate the reaction enthalpy. If the solution absorbs heat, its q value is positive; the reaction’s heat change is approximately equal in magnitude and opposite in sign, provided heat loss and calorimeter effects are negligible.

A bomb calorimeter operates differently. It holds a sample in a sealed, rigid vessel and is commonly used to measure combustion energy at constant volume. Under suitable conditions, the measured heat is related to the change in internal energy, ΔU, rather than directly to ΔH. This distinction is important when interpreting experimental thermochemistry data.

Choose a reaction type to see an illustrative heat-transfer display. The temperature change shown is a conceptual example rather than an experimental prediction for a specific quantity of reactants. Actual temperature changes depend on reaction extent, heat capacity, concentration and heat loss.

The standard enthalpy of formation, ΔH°_f, is the enthalpy change when one mole of a substance is formed from its constituent elements in their standard states. Standard formation enthalpies are useful because they allow reaction enthalpies to be calculated from tabulated values, without measuring every reaction directly.

These approximate values are commonly tabulated at 298.15 K and a standard pressure of 1 bar. The physical state matters: liquid water and water vapour have different standard enthalpies of formation. Elements in their standard states have a standard enthalpy of formation of zero by convention.

Endothermic and exothermic sign conventions

The sign of ΔH describes the enthalpy change of the system. A negative value means that the system loses enthalpy in the process as written. A positive value means the system gains enthalpy. The surroundings usually respond in the opposite direction when heat exchange dominates and other energy transfers are negligible.

Hess’s Law

Hess’s Law states that the overall enthalpy change is independent of the route between the same initial and final states. This works because enthalpy is a state function. If a reaction can be represented as a series of steps, the enthalpy changes of those steps can be added to obtain the overall change. When an equation is reversed, its ΔH changes sign; when the equation is multiplied by a number, its ΔH is multiplied by the same number.

Bond enthalpy method

Breaking chemical bonds requires energy, whereas forming bonds releases energy. The approximate reaction enthalpy can therefore be estimated by subtracting the total bond enthalpies of bonds formed from those of bonds broken. This method generally provides an estimate because tabulated bond enthalpies are average values and can depend on the molecular environment.

Combustion and energy production

Thermochemistry helps quantify the energy released by fuels such as methane, petrol and hydrogen. Engineers use reaction enthalpies to compare fuels, estimate heat output and plan combustion processes.

Food energy and metabolism

Calorimetric methods help measure energy released by food during combustion. Biological energy conversion is more complex than direct burning, but energy accounting remains essential for understanding metabolic processes.

Industrial chemical manufacturing

Many industrial reactions release or absorb substantial heat. Thermochemical calculations help determine cooling or heating requirements, improve process safety and maintain suitable operating conditions.

Cold packs and heat packs

Some dissolution processes absorb heat and can be used in instant cold packs. Other chemical processes release heat and are used in warming products. The observed temperature change depends on the amount of material, the enthalpy change and the heat capacity of the surroundings.

Environmental and materials chemistry

Enthalpy data support the study of fuel combustion, cement production, material synthesis and other processes with energy demands or environmental consequences. Thermochemistry can help compare process options, although a full environmental assessment requires more than reaction enthalpy alone.

• Reversing the sign of ΔH: Exothermic reactions have ΔH < 0, while endothermic reactions have ΔH > 0 for the process as written.
• Confusing system and surroundings: Heat gained by the system is heat lost by the surroundings when the exchange is isolated and other transfers are negligible.
• Using Celsius in a temperature ratio: Temperature differences have the same numerical value in kelvin and Celsius, but thermodynamic equations and absolute-temperature ratios require kelvin.
• Forgetting stoichiometric coefficients: Multiply each standard enthalpy of formation by its coefficient in the balanced equation.
• Ignoring physical states: Formation enthalpies for liquid water and water vapour are different.
• Changing a reaction equation without changing ΔH: Reversing a reaction reverses the sign of ΔH; multiplying the equation scales ΔH by the same factor.
• Assuming every calorimeter is constant-pressure: A coffee-cup calorimeter is commonly used at constant pressure, whereas a bomb calorimeter is a constant-volume device.

Review the video alongside the explanations, equations, numerical examples and practice quiz on this page.

Thermochemistry explains how energy is transferred as heat during reactions and physical changes. The first law expresses energy conservation, enthalpy is especially useful for constant-pressure processes, and calorimetry connects measured temperature changes with heat transfer.

Remember the central relationships: ΔU = q + w, H = U + PV, q = mcΔT, and the formation-enthalpy method for calculating reaction enthalpy. Hess’s Law allows enthalpy changes from multiple steps to be combined. Always balance equations, use the correct physical states and keep track of the sign convention.

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