Chemistry AP Study Notes 5 (Thermodynamics)

Thermodynamics
Enthalpy

• Enthalpy is the heat gained or lost by a system under constant pressure conditions, ΔH. ΔH > 0 if the reaction is endothermic. ΔH < 0 if the reaction is exothermic.
• Exothermic means that the temperature goes up during the reaction. Endothermic means the temperature goes down during the reaction.
• Calorimetry is a technique to measure the heat released or absorbed during a change. That quantity is known as q.
• Heat capacity is the amount of heat needed to change the temperature 1K.

C_p = heat capacity = q/ΔT (units of joules/kelvin)

• Specific heat capacity is the amount of heat needed to raise one gram of a substance 1K.

c = q/(mΔT) (units of joules per gram-kelvin)

• Molar heat capacity is the amount of heat needed to raise one mole of a substance 1K.
• Hess' Law states that if a reaction occurs in steps then the total enthalpy change will equal the totals of the individual steps. 
• You do not have to know the actual steps because heat reactions are a state function. That means that the total only depends on the beginning and end states, not on the pathway used to get there. 
• The heat of formation of a product is symbolized by ΔH_f
• A degree to the right of the ΔH implies a standard state (1 atm, 1 M, etc). ΔH°
• So ΔH°_f gives the total heat of formation when 1 mole of a substance is formed from elements and all the substances are in their standard state. This is the standard enthalpy of formation.
• The ΔH°_f of an element in its standard state is zero.
• The ΔH°_{f rxn} for a reaction is the sum of all the ΔH°_f for the products minus those for the reactants.

Entropy

• The First Law of Thermodynamics is that the total energy of the universe is constant. It amounts to the Law of Conservation of Energy.
• The Second Law of Thermodynamics is the famed entropy law. Entropy is the inevitable overall tendency of a system toward disorder. ΔS_universe = ΔS_system + ΔS_surroundings > 0.
• The entropy increases: 1) when the number of molecules increases during a reaction, 2) with an increase in temperature, 3) when a gas is formed from a liquid or solid, and 4) when a liquid is formed from a solid.
• The standard molar entropy (S°) can be summed up like the standard enthalpy (ΔH°). You take the sum of the entropies of the products and subtract from them the sum of the entropies of the reactants.

Gibbs Free Energy

• Some guidelines for predicting a spontaneous reaction are a negative enthalpy and a positive entropy. These are put together in the Gibbs free energy equation.

ΔG = ΔH - TΔS

• ΔG is the best indicator as to whether a spontaneous reaction will occur.
• If ΔG > 0, the reaction will not be spontaneous. More energy is needed.
• If ΔG < 0, the reaction will be spontaneous.
• If ΔG = 0, the reaction is in equilibrium.
• The standard Gibbs free energy change ΔG° is again the sum of the products minus the sum of the reactants.
• If the concentrations or pressures are not 1, then we need the Gibbs free energy equation for non-standard conditions:

ΔG = ΔG° + ln RT Q or ΔG° + 2.303 log Q

where Q is the ratio of the sum of products over the sum of reactants, the activity quotient.