# 12.2: Thermodynamics: Basic Terms

- Page ID
- 8242

## Internal Energy, U

The internal energy of a system is the sum of the potential energy and the kinetic energy. For many applications it is necessary to consider a small change in the internal energy, dU, of a system.

dU = dq + dw = CdT - PdV

dq = the heat supplied to a system

dw = the work performed on the system

C = heat capacity

dT = change in temperature

P = pressure

dV = change in volume

At constant volume,

dU = C_{V}dT

## Enthalpy, H

Enthalpy is the constant pressure version of the internal energy. Enthalpy,

H = U + PV.

Therefore, for small changes in enthalpy,

dH = dU + PdV + VdP.

At constant pressure,

dH = C_{P}dT.

## Entropy, S

Entropy is a measure of the disorder of a system. In terms of molecular disorder, the entropy consists of the configurational disorder (the arrangement of different atoms over identical sites) and the thermal vibrations of the atoms about their mean positions. A change in entropy is defined as,

\[dS \geq \frac{dq}{T}\]

For reversible changes, i.e. changes under equilibrium conditions,

dq = TdS

For natural changes, i.e. under non-equilibrium conditions,

dq < TdS

## Gibbs free energy, G

The Gibbs free energy can be used to define the equilibrium state of a system. It considers only the properties of the system and not the properties of its surroundings. It can be thought of as the energy which is available in the system to do useful work.

Free energy, G, is defined as,

G = H - TS = U + PV - TS

For small changes,

dG = dH - TdS - SdT = VdP - SdT + (dq - TdS)

For changes occurring at constant pressure and temperature,

dG = dq - TdS

Therefore, dG = 0 for reversible (equilibrium) changes, and dG < 0 for non-reversible changes.

From this it is clear that G tends to a minimum at equilibrium.

The Helmholtz free energy, F, is sometimes used instead of G, and is the equivalent of G for changes at constant volume. It is defined as,

F = U - TS

## Thermodynamics of Solutions

Consider a mechanical mixture of two phases, A and B. If this is then transformed into a single solution phase with A and B atoms distributed randomly over the atomic sites, then there will be,

- An enthalpy change associated with interactions between the A and B atoms, ΔH
_{mix} - An entropy change, ΔS
_{mix}, associated with the random mixing of the atoms - A free energy of mixing, ΔG
_{mix}= ΔH_{mix}- TΔS_{mix}

Assume that the system consists of N atoms: x_{A}N of A and x_{B}N of B, where,

x_{A} = fraction of A atoms and x_{B} = (1 - x_{A}) = fraction of B atoms

## Enthalpy of mixing

In calculating ΔH_{mix} it is assumed that only the potential energy term undergoes any significant change during mixing. This change arises from the interactions between nearest-neighbour atoms. Consider an alloy consisting of atoms A and B. If the atoms prefer like neighbours, A atoms will tend to cluster and likewise B atoms, so a greater number of A-A and B-B bonds will form. If the atoms prefer unlike neighbours a greater number of A-B bonds will form. If there is no preference A and B atoms will be randomly distributed.

Let w_{AA} be the interaction energy between A - A nearest neighbours, w_{BB }that for B - B nearest neighbours and w_{AB} that for A - B nearest neighbours.

All of these energies are negative, as the zero in potential energy is for infinite separation between atoms.

Let each atom of A and B have co-ordination number z.

Therefore, the total number of nearest-neighbour pairs is Nz/2.

Probability of A - A neighbours = x_{A}^{2}

Probability of B - B neighbours = x_{B}^{2}

Probability of A - B neighbours = 2x_{A}x_{B}

For a solid solution the total interaction energy is,

H_{s} - U_{s} = Nz/2 (x_{A}^{2} w_{AA} + x_{B}^{2} w_{BB} + 2x_{A}x_{B} w_{AB})

For pure A, H_{A} = (Nz/2)w_{AA}

For pure B, H_{B} = (Nz/2)w_{BB}

Hence the enthalpy of mixing is given by,

ΔH_{mix} = H_{s} - (x_{A}H_{A} + x_{B}H_{B}) = (Nz/2)x_{A}x_{B} (2w_{AB} - w_{AA} - w_{BB})

We can define an interaction parameter

W = (Nz/2)(2w_{AB} - w_{AA} - w_{BB})

Therefore,

ΔH_{mix} = Wx_{A}x_{B}

If A-A and B-B interactions are energetically more favourable than A-B interactions then W > 0. So, ΔH_{mix} > 0 and there is a tendency for the solution to form A-rich and B-rich regions.

If A-B interactions are energetically more favourable than A-A and B-B interactions, W < 0, ΔH_{mix} < 0, and there is a tendency to form ordered structures or intermediate compounds.

Finally if the solution is ideal and all interactions are energetically equivalent, then W = 0 and ΔH_{mix} = 0.

## Entropy of mixing

Per mole of sites, this is

ΔS_{mix} = kN (- x_{A}lnx_{A} - x_{B}lnx_{B})

(the derivation of this result makes use of Stirling's approximation)

where N = Avogadro's number, and kN = R, the gas constant.

Hence,

ΔS_{mix} = R (- x_{A}lnx_{A} - x_{B}lnx_{B})

A graph of ΔS_{mix} versus x_{A} has a different form from ΔH_{mix}. The curve has an infinite gradient at x_{A} = 0 and x_{A} = 1.

The free energy of mixing is now given by,

ΔG_{mix} = ΔH_{mix} - TΔS_{mix} = x_{A}x_{B}W + RT (x_{A} lnx_{A} + x_{B}lnx_{B})

For W < 0, ΔG_{mix} is negative at all temperatures, and mixing is exothermic. For W > 0, ΔH_{mix} is positive and mixing is endothermic.