Basic concepts of thermochemistry

Thermochemistry - the study of the energy changes that accompany physical or chemical changes in matter.

Work - the amount of energy transferred by a force over a distance; SI units - joules (J).

Energy - the ability to do work; SI units - joules (J).

Potential energy
the energy of a body or system due to its position or composition
Kinetic energy
the energy of an object due to its motion


Thermal energy - the total quantity of kinetic and potential energy in a substance.

Heat - the transfer of thermal energy from a warm object to a cooler object.

Temperature - a measure of the average kinetic energy of entities in a substance.

Chemical system - a group of reactants and products being studied .

Surroundings - all the matter that is not part of the system

Law of Conservation of Energy

Energy cannot be created or destroyed but can be converted from one form to another.

3 Types of thermodynamic systems
Open system
a system in which both matter and energy are free to enter and leave the system
Closed system
a system in which energy can enter and leave the system, but matter cannot
Isolated system
neither matter nor energy can move into or out of the system.


  1. Endothermic and Exothermic Reactions
  2. Nuclear Energy
  3. Heat Capacity
  4. Calorimetry and Thermal Energy Transfer
  5. Enthalpy Change
  6. Molar Enthalpy Change
  7. Representing Molar Enthalpy Changes
  8. Bond Energies
  9. Enthalpy and Bond Energies
  10. Enthalpy Change and Hess’s Law
  11. Standard Enthalpy of Formation

Endothermic and Exothermic Reactions

Exothermic - releasing energy to the surroundings

The combustion of methane releases a quantity of energy with the magnitude ∆Ep, which flows to the surroundings, mainly as thermal energy and light energy. This is an exothermic reaction.

Energy Changes during an Exothermic Reaction in an Open System

Energy Changes during an Exothermic Reaction in an Open System

Endothermic - absorbing energy from the surroundings

The synthesis of nitric oxide from its elements absorbs a quantity of thermal energy with the magnitude ∆Ep from the surroundings. This is an endothermic reaction.

Energy Changes during an Endothermic Reaction in an Open System

Energy Changes during an Endothermic Reaction in an Open System

Nuclear Energy

All nuclear reactions are exothermic.

Fusion - the process of combining two or more nuclei of low atomic mass to form a heavier, more stable nucleus.


An element with the chemical symbol X is represented as AZX, where

  • A is the mass number (number of protons plus number of neutrons) and
  • Z is the atomic number (number of protons).

A neutron is a subatomic entity, and the symbol 10n is used to show that a neutron has a mass number of 1 and an atomic number of 0.

Fission - the process of using a neutron to split a nucleus of high atomic mass into two nuclei with smaller masses.


Nuclear power plants use the fission of uranium-235 to produce electricity.

When a neutron collides with uranium-235 nucleus, the uranium nucleus splits into smaller nuclei and releases energy and additional neutrons.

Uranium fission

Heat Capacity

Specific heat capacity (c) - the quantity of thermal energy required to raise the temperature of 1 g of a substance by 1 °C; SI units – J/(g∙°C).

Substances with a high specific heat capacity take longer to heat or to cool.

Specific Heat Capacities of Some Common Substances

Specific Heat Capacities of Some Common Substances

Calorimetry and Thermal Energy Transfer

Calorimetry - the experimental process of measuring the thermal energy change in a chemical or physical change.

Calorimeter - a device that is used to measure thermal energy changes in a chemical or physical change.

In general, a calorimeter consists of:

  • a well-insulated reaction chamber,
  • a tight-fitting cover with insulated holes for a thermometer, and
  • some mechanism to stir the calorimeter contents.

Using an insulated chamber minimizes energy losses to the surroundings. A tight lid prevents matter from leaving or entering the calorimeter.

In calorimetry, the total amount of thermal energy absorbed or released by a chemical system is given the symbol q.

The magnitude of q depends on three factors:

  1. the mass of the substance
  2. the specific heat capacity of the substance
  3. the temperature change experienced by the substance as it warms or cools

The value of q is calculated using the equation

q = mc∆T


  • m is the mass of the substance,
  • c is the specifi c heat capacity of the substance,
  • ∆T is the change in temperature of that substance.

The value of q has two parts:

  1. the magnitude of q tells you how much energy is involved,
  2. the sign tells you the direction of energy transfer.

If q has a negative value, the system transfers thermal energy to its surroundings and the change is exothermic.

Enthalpy Change

Enthalpy (H) - the total amount of thermal energy in a substance.

Enthalpy change (ΔH) - the energy released to or absorbed from the surroundings during a chemical or physical change.

As long as pressure remains constant, the enthalpy change of the chemical system is equal to the flow of thermal energy in and out of the system.

ΔH system = | qsystem |

For a chemical reaction, the enthalpy change, ΔH, is given by the equation -

ΔH = Hproducts - Hreactants

When the products of a reaction have a greater enthalpy than the reactants, ΔH will be positive. The system absorbs thermal energy from its surroundings and the reaction is endothermic.

On the other hand, if the enthalpy of the products is less than that of the reactants, ΔH will be negative. In this case, the system releases thermal energy to its surroundings and the reaction is exothermic.

Molar Enthalpy Change

Molar enthalpy change (ΔHr) - the enthalpy change associated with a physical, chemical, or nuclear change involving 1 mol of a substance; SI units – J/mol.

To write the balanced equation for the molar enthalpy change of formation of a product, the coefficient of that product must always be 1. Other substances in the equation may have fractional coefficients as a result.

Molar enthalpy change

The quantity of energy involved in a change (the enthalpy change, ∆H, expressed in kJ) depends on the quantity of matter that undergoes the change.

To calculate an enthalpy change, ∆H, for some amount of substance other than 1 mol, you need to obtain the molar enthalpy value, ∆Hr, from a reference source, and then use the formula

∆H = n∆Hr,

where n is the amount and ∆Hr is the molar enthalpy change of the reaction.

Representing Molar Enthalpy Changes

Potential energy diagram - a graphical representation of the energy transferred during a physical or a chemical change.

Potential energy diagram

(a) The condensation reaction of 1 mol of water vapour is exothermic. The reactant has a higher potential energy than the product.

(b) The vaporization reaction of liquid water to water vapour is endothermic. The reactant has a lower potential energy than the product.

Bond Energies

Bond dissociation energy - the energy  required to break a given chemical bond.

Average Bond Energies (kJ/mol)

Average Bond Energies

Multiple bonds have larger bond energies than single bonds.

A relationship exists between the number of bonds between atoms (that is, number of electrons shared) and the length of a covalent bond (that is, distance between nuclei):

as the number of bonds increases, the bond length shortens.

Bond length

The bond dissociation energy of a given bond depends on the types of atoms and bonds in the same molecule. Therefore, the bond energy of a C–H bond is affected by the number of atoms and bonds around it.

Bond dissociation energy

The use of an average bond energy is convenient for predicting enthalpy changes in chemical reactions.

Enthalpy and Bond Energies

Bond energy values can be used to calculate approximate enthalpy changes, ΔH, for reactions.

During a chemical reaction, the bonds in the reactants must first break.

  • For bonds to be broken, energy must be added—an endothermic process.
  • Making new bonds in the products releases energy—an exothermic process.

The enthalpy change for a reaction 

The enthalpy change for a reaction

  • n is the amount (in moles) of a particular bond type, and
  • D is the bond energy per mole of bonds.

The first step in using bond energies to predict ∆H for a reaction is to determine how many of each type of bond must be broken in the reactants.

Next, determine the number of bonds of each type that form in the products.

Finally, use bond energy data to calculate the total energy required to break the reactant bonds, followed by the total energy released by the formation of product bonds.

The energy change, ∆H, of the reaction is the difference between these two sums. The bonds between specific atoms and the bond energies associated with these bonds are very similar, even when the bonds are located in different entities.

Enthalpy Change and Hess’s Law

From experimental evidence, chemists have found that the change in enthalpy in a chemical process is independent of the path taken.

This means that in going from an initial set of reactants to a final set of products, the change in enthalpy is the same regardless of whether the conversion happens in one step or in a series of steps.

Hess's law

The enthalpy change for the conversion of reactants to products is

the same whether the conversion occurs in one step or several steps.


Hess’s Law

Hess’s law is very useful for studying energy changes in chemical reactions that cannot be analyzed using calorimetry.

To use Hess’s law to calculate enthalpy changes for chemical reactions, it is necessary to apply the following two rules:

  1. If you reverse a chemical reaction, you must also reverse the sign of ΔH.
  2. The magnitude of ΔH is directly proportional to the number of moles of reactants and products in a reaction. If the coefficients in a balanced equation are multiplied by a factor, the value of ΔH is multiplied by the same factor.

Standard Enthalpy of Formation

Standard enthalpy of formation (ΔHf °) - the change in enthalpy that accompanies  the formation of 1 mol of a compound  from its elements in their standard states.

A substance is in a standard state when it is in its most stable  form at SATP (standard conditions for temperature and pressure - 25 °C, 100 kPa). The standard state for a substance in solution is at a concentration  of 1 mol/L.

A schematic diagram of the energy changes for the reaction represented by the equation

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

The energy changes for the reaction

The enthalpy change for any reaction can be calculated by subtracting the sum of the enthalpies of formation of the reactants from the sum of the enthalpies of formation of the products.

The enthalpy change

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Reaction rates

Chemical equilibrium