AP Chemistry Unit9: Thermodynamics and Electrochemistry
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9.1 Introduction to Entropy
Spontaneous/Thermodynamically favored
A process that occurs without outside intervention
Entropy (∆S)
The measure of randomness, disorder, or degrees of freedom (chaos)
The Second Law of Thermodynamics
∆S > 0 = favored process
∆S < 0 = not favored
9.2 Absolute Entropy and Entropy Change
Standard Entropy
298K and 1 atm
∆Srxn° = ∑npS°(products) – ∑nrS°(reactants)
9.3 Gibbs Free Energy and Thermodynamic Favorability
Gibbs Free Energy
∆G < 0 = thermodynamically favored
∆G > 0 = not thermodyanmically favored
∆G° = ∑npGf°(products) – ∑nrGf°(reactants)
∆G° = ∆H° – T∆S°
9.4 Thermodynamics and Kinetic Control
Thermodynamically favorable ≠ Fast reaction speed
Reaction rate is determined by the activation energy.
A reaction that is thermodynamically favored and with a large activation energy will have a slow reaction rate.
Adding a catalyst has no effect on ∆G
Catalysts reduce the activation energy and increases the reaction rate, but because it doesn’t affect the values of ∆H and ∆S, it will not affect the thermodynamics of a reaction (∆G).
9.5 Free Energy and Equilibrium
Free Energy at Nonstandard Conditions
∆G = ∆G° + RTln(Q) This equation is not given on the equation sheet.
R = 8.314 J/mol•K
T = temperature (K)
Q = reaction quotient
∆G° = –RTln(K) This equation is given on the equation sheet.
This is at equilibrium
R = 8.314 J/mol•K
T = temperature (K)
K = equilibrium constant
Le Chatelier’s Principle and ∆G
Use Le Chatelier’s Principles to identify if the forward reaction is becoming more or less thermodynamically favorable.
Remember that changes in liquid or solids will not impact thermodynamically favorability.
Combining (couple) an unfavorable reaction with a favorable reaction to make the overall process favorable (–∆G). The reactions must have a common intermediate.
9.8 Galvanic (Voltaic) and Electrolytic Cells
Electrochemistry
Study of the interaction between chemical and electrical energy
Oxidation-reduction (redox) Reaction
Transfer of elections from one substance to another.
Helpful ways to remember
LEO
Lose Electrons Oxidation
GER
Gain Electrons Reduction
Half Reactions
The overall reaction is split into two half reactions: oxidation and reduction.
Electric Current = Opposite of Electron Flow
Electrons move towards the reduction side, which is the current's opposite direction.
Anode = Oxidation
Example half reaction (where X is the ion): X (s) → X(charge) + e–
Metals are likely to be oxidized
Cathode = Reduction
Example half reaction (where X is the ion): X(charge) + e– → X (s)
Halogens are likely to be reduced
Galvanic (Voltaic) Cell
Voltmeter
The voltmeter measures the current without impeding the flow of electrons.
Electron Flow
The electrons flow to the right. This means the left side is the oxidation reaction, while the right is the reduction reaction.
Electrons always flow from the anode to the cathode (applies for both Galvanic/Voltaic and Electrolytic cells)
The Salt Bridge
It is not involved in the overall reaction but is required for Galvanic cells to function and produce voltage.
Prevents charge buildup by maintaining electrical neutrality and avoiding the end of the flow of electrons. It essentially completes the circuit to allow the reaction to proceed.
Gain/Loss of Mass
Throughout this reaction, the Zinc bar loses mass as the reaction proceeds, while the Copper bar gains mass.
In Galvanic cells, the anode always loses mass as the reaction proceeds, while the cathode gains mass.
Active Metal
The metal that loses electrons more easily.
In Galvanic cells, the anode will contain the more active metal.
9.9 Cell Potential and Free Energy
Cell Potential (E°cell)
Force made on the electrons
Galvanic Cell Potiential
The half-reaction must be reversed when the overall cell potential is positive.
Free Energy and Cell Potential
∆G° = –nFE°
n = number of moles of electrons
F = Faraday Constant (96485 coulombs per mole of electrons)
E° = cell potential (V)
If E° < 0, then ∆G° > 0 (nonspontaneous)
If E° > 0, then ∆G° < 0 (spontaneous)
9.10 Cell Potential Under Non-Standard Conditions
Nernst Equation
Used if the electrochemical cell is NOT in standard state conditions.
Under standard state, Q = 1 ln(1) = 0 Thus, Ecell = E°cell
At equilibrium Q = K and Ecell is 0
When Q < K
Ecell > 0 (spontaneous)
When Q > K
Ecell < 0 (nonspontaneous)
When K > 1, E°cell is positive
When K < 1, E°cell is negative
Increasing the size of solid metals in a Galvanic cell will not affect the Q value. Thus, the voltage will not change.
Concentration Cell
A type of electrochemical cell where both electrodes are made of the same material, but the concentrations in the two half-cells are different.
E° = 0
Both half-reactions involve the same redox couple, so their standard electrode potentials will cancel each other out.
Differences in concentrations drive the reaction instead of the difference in metal type.
A running galvanic cell is NOT at equilibrium. Thus, Le Chatelier’s Principle can’t be used when comparing changes.
Galvanic cells “run” towards equilibrium.
An Ecell value with a nonzero number runs and produces a voltage. After it reaches equilibrium/0, it is a “dead” battery.
