5 Steps to a 5 AP Chemistry

This is not molar (yet)

Thus:

Laws of Thermodynamics

The First Law of Thermodynamicsstates that the total energy of the universe is constant.

This is simply the Law of Conservation of Energy. This can be mathematically stated as:

ΔEuniverse=ΔEsystem+ΔEsurroundings= 0

The Second Law of Thermodynamics involves a term called entropy. Entropy(S) is related

to the disorder of a system. The Second Law of Thermodynamicsstates that all processes

that occur spontaneously move in the direction of an increase in entropy of the universe

(system +surroundings). Mathematically, this can be stated as:

ΔSuniverse=ΔSsystem+ΔSsurroundings> 0 for a spontaneous process

For a reversible process ΔSuniverse=0. The qualitative entropy change (increase or decrease

of entropy) for a system can sometimes be determined using a few simple rules:

1. Entropy increases when the number of molecules increases during a reaction.


2. Entropy increases with an increase in temperature.


3. Entropy increases when a gas is formed from a liquid or solid.


4. Entropy increases when a liquid is formed from a solid.
   Let us now look at some applications of these first two laws of thermodynamics.

Products Minus Reactants

Enthalpies

Many of the reactions that chemists study are reactions that occur at constant pressure.

During the discussion of the coffee-cup calorimeter, the heat change at constant tempera-

ture was defined as qp. Because this constant-pressure situation is so common in chemistry,

a special thermodynamic term is used to describe this energy: enthalpy. The enthalpy

change, ΔH, is equal to the heat gained or lost by the system under constant-pressure

conditions. The following sign conventions apply:

If ΔH>0 the reaction is endothermic

If ΔH<0 the reaction is exothermic

If a reaction is involved, ΔHis sometimes called ΔHreaction. ΔHis often given in association

with a particular reaction. For example, the enthalpy change associated with the formation

of water from hydrogen and oxygen gases can be shown in this fashion:

2 H 2 (g) +O 2 (g) →2 H 2 O(g) ΔH=−483.6 kJ

The negative sign indicates that this reaction is exothermic. This value of ΔHis for the

production of 2 mol of water. If 4 mol were produced, ΔHwould be twice −483.6 kJ. The

techniques developed in working reaction stoichiometry problems (see the Stoichiometry

chapter) also apply here.

−

×

− =−

28 52

3234

.kJ

8.8177 10 mol

3 kJ/mol

(.1 5886g) 8 8177 10.^3

(1mol)

(180.16 g)

=×−mol

126  Step 4. Review the Knowledge You Need to Score High

KEY IDEA