Measuring energy changes

I. Introduction to Energy Changes

Exothermic reactions release energy, which is often evident as an increase in temperature.
Endothermic reactions absorb energy, typically causing a decrease in temperature.
The universe is composed of systems and surroundings; energy is transferred between these two entities.
The term ‘system’ refers to the reactants and products; ‘surroundings’ means everything external to the system.

II. Enthalpy (H)

Enthalpy is a measure of the total energy of a system.
It cannot be measured directly, but the change in enthalpy (∆H) can be.
For an exothermic reaction, ∆H is negative as energy is released to the surroundings.
For an endothermic reaction, ∆H is positive because energy is absorbed from the surroundings.

III. Enthalpy Changes

Enthalpy change of reaction (∆Hr) is the enthalpy change when the mole quantities of reactants as stated in a balanced equation react together.
Standard enthalpy change of formation (∆Hf) is the enthalpy change when one mole of a substance is formed from its elements in their standard states.
Standard enthalpy change of combustion (∆Hc) is the enthalpy change when one mole of a substance is completely burned in oxygen.
Enthalpy change due to bond breaking: bonds do not break freely; instead, energy must be inputted for this to happen, making bond breaking an endothermic process.
Enthalpy change due to bond formation: when bonds form, energy is released, making it an exothermic process.

IV. Measuring Enthalpy Changes

Energy changes can be measured using a thermometer in a calorimeter.
Using the heat capacity (specific heat) formula q = mc∆T, one can calculate energy change; where q is the heat energy, m is the mass, c is the specific heat capacity, and ∆T is the change in temperature.
Assumptions made for calorimetry: no heat loss to surroundings, all the heat is absorbed by the solution, and the specific heat capacity of all solutions is the same as water.

V. Hess’s Law

Hess’s law states that the total enthalpy change of a reaction is independent of the pathway, provided the starting and final conditions are the same.
It allows for the calculation of energy changes that are difficult to measure directly.
This law is used to create Hess’s cycles which use known enthalpy changes to determine unknown ones.

VI. Bond Enthalpies

Bond enthalpy is the energy needed to break one mole of a bond in a molecule in the gaseous state.
It is an average value, as the energy required to break a specific bond varies in different molecules.
Bond enthalpy values can be used to calculate the enthalpy change for reactions using the equation ∆H = bond breaking - bond making.