# Unit 9: Applications of Thermodynamics

## Entropy

• Entropy, S, is the amount of disorder or chaos in a system. More disorder, greater S value.

• Standard entropy is S° and measured at 25 celsius

• Standard entropy change ∆S° is measure at the end of a reaction

• ∆S° = (sum of ∆S° products) - (sum of ∆S° reactants)

• If a reaction goes from less moles to more moles (such as 2 moles on the reactant side to 3 moles on the product side) there is more disorder and a positive ∆S

• If a reaction goes from a gas to liquid, liquid to solid, or gas to solid, the reaction has a negative ∆S

• If bonds are broken and phase change becomes more disordered, the ∆S is positive

## Gibbs Free Energy

• ∆G is Gibbs Free Energy which determines if a process is thermodynamically favored or unfavored, also known as spontaneous or nonspontaneous

### Free Energy Change

• Standard free energy change, ∆G°, is calculated the same as ∆S°

• ∆G° = (sum of ∆G° products) - (sum of ∆G° reactants)

• For a reaction,

• If ∆G is negative, it is TFP (thermodynamically favored process)

• If ∆G is positive, it is not TFP

• If ∆G is 0, it is at equilibrium

### ∆G, ∆H, and ∆S

• TFP must result in decreasing enthalpy, increasing entropy, or both

• ∆G° = ∆H° - T∆S°

• T = temperature in Kelvin

• ∆S° is usually given in j/mol*K and must be converted to kj/mol*K

• Gibbs Free Energy is usually kj/mol*K

∆H

∆S

T

∆G

Favorability

-

+

LowHigh

--

Always TFP

+

-

LowHigh

++

Never TFP

+

+

Low High

+-

Not TFPTFP

-

-

Low High

-+

TFPNot TFP

### Standard Free Energy Change and the Equilibrium Constant

• Gibbs free energy can be calculated if equilibrium constant is known

• ∆G° = -RT(ln K)

• R = gas constant (8.31 j/mol*k)

• T = kelvin temperature

• K = equilibrium constant

• If ∆G° is negative, K is greater than 1, the products are favored at equilibrium

• If ∆G° is positive, K must be less than 1, the reactants are favored at equilibrium

### Reduction Potentials

• Every half reaction has electric potential. Potentials are given as reduction half-reactions. If the reaction is reversed, flip the sign to get the oxidation potential

## Galvanic Cells

• Galvanic cells (voltaic cell) use favored redox reactions to generate current

• Two half-reactions take place in separate chambers and the electrons from the oxidation pass to the reduction reaction which creates the current

• Current is defined as the flow of electrons from one place to another

• Oxidation takes place at the anode electrode and reduction takes place at the cathode electrode

• The salt bridge keeps electrical neutrality. Without the salt bridge the voltage would be zero. The potassium ion flows to the cathode and the chlorine flows to the anode.

• The cell voltage is equal to the total redox reaction voltage.

### Non-Standard Conditions

• Reduction potentials are give at standard conditions, 25 celsius, 1 atm, and 1 M

• Voltaic cells are very favored with equilibrium constant greater than 1. If the Q = K however, the voltage would drop to ero.

• If the reaction quotient increased it would become close to the equilibrium constant and the voltage would decrease.

## Electrolytic Cells

• Electrolytic cells use outside voltage sources to power unfavored redox reactions and mainly occur in aqueous solutions.

• The sign of total cell potential is always negative

### Electroplating

• Electrolytic cells are used for electroplating.

• I = (q/t)

• I = Current (amperes, A)

• q = charge (coulombs, C)

• t = time (second, s)

• Moles of electrons = (coulombs/ 96,500 coulombs per mol)

## Voltage and Favorability

• Redox is favored if the potetial has a positive value. reaction potential can be calculate gibb’s free energy

• ∆G° = -nFE°

• n = number of moles of electrons exchanged in the reaction

• F = Faraday’s constant. 96,500 coulombs/mol

• E° = standard reaction potential (V)

• If E° is positive, ∆G° is negative and is TFP

• If E° is negative, ∆G° is positive and not TFP