# Unit 6- Thermodynamics -Summary

Thermodynamics deals with energy changes in chemical or physical processes and enables us to study these changes quantitatively and to make useful predictions.

For these purposes, we divide the universe into the system and the surroundings. Chemical or physical processes lead to evolution or absorption of heat (q), part of which may be converted into work (w).

These quantities are related through the first law of thermodynamics via ∆U = q + w. ∆U, change in internal energy, depends on initial and final states only and is a state function, whereas q and w depend on the path and are not the state functions.

We follow sign conventions of q and w by giving the positive sign to these quantities when these are added to the system.

We can measure the transfer of heat from one system to another which causes the change in temperature. The magnitude of rise in temperature depends on the heat capacity (C) of a substance.

Therefore, heat absorbed or evolved is q = C ∆T. Work can be measured by w = –pex∆V, in case of expansion of gases.

Under reversible process, we can put pex = p for infinitesimal changes in the volume making wrev = – p ∆V. In this condition, we can use gas equation, pV = nRT.

At constant volume, w = 0, then ∆U = qV , heat transfer at constant volume. But in study of chemical reactions, we usually have constant pressure.

We define another state function enthalpy. Enthalpy change, ∆H = ∆U + ∆ngRT, can be found directly from the heat changes at constant pressure, ∆H = qp.

There are varieties of enthalpy changes. Changes of phase such as melting, vaporization and sublimation usually occur at constant temperature and can be characterized by enthalpy changes which are always positive.

Enthalpy of formation, combustion and other enthalpy changes can be calculated using Hess’s law. Enthalpy change for chemical reactions can be determined by

rH = å aifH(products) - å bifH(reactants)

and in gaseous state by

rH = å bond enthalpies of the reactants - å bond enthalpies of the products

First law of thermodynamics does not guide us about the direction of chemical reactions i.e., what is the driving force of a chemical reaction.

For isolated systems, ∆U = 0. We define another state function, S, entropy for this purpose. Entropy is a measure of disorder or randomness.

For a spontaneous change, total entropy change is positive. Therefore, for an isolated system, ∆U = 0, ∆S > 0, so entropy change distinguishes a spontaneous change, while energy change does not.

Entropy changes can be measured by the equation

∆S = qrev/T for a reversible process, qrev/T is independent of path

Chemical reactions are generally carried at constant pressure, so we define another state function Gibbs energy, G, which is related to entropy and enthalpy changes of the system by the equation:

rG  = ∆rH - T∆rS

For a spontaneous change, ∆Gsys < 0 and at equilibrium, ∆Gsys = 0

Standard Gibbs energy change is related to equilibrium constant by  ∆rG  = -RT ln K

K can be calculated from this equation, if we know ∆rG  which can be found from ∆rG  = ∆rH - T∆rS .

Temperature is an important factor in the equation.

Many reactions which are non-spontaneous at low temperature, are made spontaneous at high temperature for systems having positive entropy of reaction.