CH353 - Physical Chemistry I
Spring 2015, Unique 51170

Lecture Summary, 5 February 2015

Thermochemistry:  We then went on and discussed the topic of thermochemistry, the science of the flow of heat due to chemical transformations:

   Reactant --> Product +/- heat

In chemical thermodynamics we almost always are interested in energies at constant pressure, therefore in reaction enthalpies of a system.  Because enthalpy is a state function, by defining a standard state of P = 1 bar and T = 298 K, we can tabulate enthalpies of formation for any substance under that standard, then use that information to calculate a reaction enthalpy under any set of circumstances, even those far away from the standard state. 
Second Law of Thermodynamics: We have spent quite a bit of time developing the first law of thermodynamics, which we have expressed as energy = heat + work.  We have found that we can find heat, work, internal energies, and enthalpies for a variety of different systems and different paths.  But clearly we are missing something.  For example, an an isothermal expansion of an ideal gas, deltaU = 0, but under some conditions, w < 0; i.e. the system does work on the surroundings.  How is this possible?  How can we do work without a change in internal energy?  Furthermore, we have seen that some transformations occur spontaneously only in one direction, even if deltaU = 0.  For example, if we open a bottle of perfume in a room, soon perfume molecules will fill the entire room, but the reverse process will never happen.  Finally, we know of simple examples in which a reaction occurs spontaneously even if it requires energy from the surroundings, for example sweating. 

We are clearly missing a kind of energy that will provide a driving force for spontaneous events to occur even with no change in internal energy.  We developed a state function, called entropy, S, which defines the direction of spontaneous change (deltaS > 0), and which will quantify this driving force:

      deltaS >= q/T.  
We have already seen that the second law of thermodynamics defines the direction of spontaneous change of a thermodynamic process.  However, in order to apply this, we have to know the total change in entropy of the universe, not just the change in entropy of our system:

    deltaS(tot) = deltaS(sys) + deltaS(surr) >= 0

The equality applies of the process is reversible, the inequality applies in all other cases.  This means that we have to determine not only the entropy change of the system but also of the surroundings.  Our rules for solving entropy problems are:

   1.  For any reversible process, deltaS(tot) = 0
   2.  To find deltaS(sys), use a reversible path from initial to final states.  You can do this because entropy is a state function and doesn't depend on the actual path.
   3.  Determine deltaS(surr) independently.  If the process is reversible, you already know deltaS(surr) = -deltaS(sys).  If the process is irreversible, deltaS(surr) = q(surr)/T, and you will have to figure out what q(surr) is.