To calculate the entropy change for reactions, we simply look at the entropy of the final state minus the entropy of the initial state. The final state is the products in their standard state and the initial state is the reactants in their standard state. We define the entropy of a pure crystalline solid at T = 0 K as having an entropy of zero. This is the third law of thermodynamics. Since we now know the entropy at T = 0 K and we can calculate the entropy change for a temperature and phase change, we can then calculate the absolute entropy for any substance at any temperature. These numbers are thankfully tabulated (usually along with the enthalpies of formation). So for a reaction we simply sum the entropies of the products (times the number of moles in our thermochemical equation) minus the sum of the entropies of the reactants (times the number of moles in the thermochemical equation).
\[\Delta S_{\rm rxn}^\circ = \Sigma n S_{\rm products}^\circ - \Sigma n S_{\rm reactants}^\circ\]
Note: \(\Delta S_{rxn}\) is not equal to \(\Delta H_{\rm rxn} / T \). This is only true at a very particular condition of equilibrium that will be investigated extensively in CH302.
\(\Delta H_{rxn}\) is very important for finding the entropy change of the surroundings. Since at constant temperature \(q = \Delta H\), we can use the enthalpy of the reaction to find the entropy change of the surroundings.
\[\Delta S_{\rm surroundings} = {-q \over T_{surroundings}} = {-\Delta H_{sys} \over T_{surroundings}}\]