1. AP Chemistry Chapter 4.9 and Chapter 17
Electrochemistry
AP Chemistry
Chapter 4.9 and Chapter 17

2. Vocabulary
Oxidation Reduction reactions are referred to as “Redox Reactions.”
Redox Reactions: electron transfer process.
Electrochemistry: study of the interchange between chemical change and electrical work.

3. Vocabulary continued Oxidation: loss of 1 or more e-
Reduction: gain of 1 or more e-
Oxidation Numbers: Imaginary charges associated with ions at different stages of a reaction.

4. Oxidation Numbers Pure Element: oxidation # is zero
Monatomic ion: oxidation number is the charge
Neutral compound: sum of oxidation #’s is zero
Polyatomic ion: sum of #’s is ion’s charge

5. Oxidation-Reduction Oxidation Involves the Loss of electrons… OIL
Oxidation is the loss of e- thus causing the oxidation number to increase (become more positive)
Reduction is the gain of e- thus causing the oxidation number to decrease (become more negative)
Oxidation Involves the Loss of electrons… OIL
Reduction Involves the Gain of electrons… RIG
Reduction reduces the oxidation number.

6. Redox Agents
When a species undergoes oxidation it is called the “reducing agent” as it causes another species to undergo reduction.
When a species undergoes reduction it is called the “oxidation agent” as it causes another species to undergo oxidation.

8. Example Find the oxidation states for each species below:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Reactants:
H: +1
C: -4
O: 0
Products:
C: +4
O:-2

9. Balancing Redox Reactions
Write separate equations (half-reactions) for oxidation and reduction
For each half reaction
Balance elements involved in e- transfer
Balance number e- lost and gained
To balance e- multiply each half-reaction by whole numbers

10. Balancing Redox Reactions: Acidic/Basic
Add half-reactions/cancel like terms (e-)
Acidic conditions:
Balance oxygen using H2O
Balance hydrogen using H+
Basic conditions:
Balance oxygen using OH-
Balance hydrogen using H2O
Check that all atoms and charges balance

11. Homework Pg. 174 Review Redox: https://youtu.be/lQ6FBA1HM3s
# 67, 71, 73, 75
Review Redox:

12. Electrochemistry
Electrochemistry encompasses the study of redox reactions that occur within electrochemical cells.
The reactions either generate electrical current in galvanic cells, or are driven by an externally applied electrical potential in electrolytic cells.

13. Types of Cells Voltaic (galvanic) cells: Electrolytic cells:
A spontaneous reaction generates electrical energy.
Electrolytic cells:
Absorb free energy from an electrical source to drive a nonspontaneous reaction.

14. Common Parts of a Cell Electrodes: Electrolyte: Salt bridge:
Conduct electricity between cell and surroundings.
Electrolyte:
Mixture of ions involved in reaction or carrying charge.
Salt bridge:
Completes circuit (provides charge balance)
A porous disk could be used instead

15. Electrodes
Oxidation occurs at the anode and reduction occurs at the cathode for all electrochemical cells.
Oxidation at the Anode (vowels)
Reduction at the Cathode (consonants)

20. Cell Potential, E◦cell
E◦cell = cell potential under standard conditions
Elements in their standard states (298K)
Solutions are 1M
Gases are at 1atm
E◦cell = E◦cathode + E◦anode

23. Homework
Pg. 831
# 25, 29, 33,
# 35, 43, 47,

24. Electrical Work
In a galvanic cell, the electrons are transferred as a result of the electrodes’ abilities to “push” the electrons away in an attempt to reach a sense of equilibrium.
This “push” of electrons through the cell results in electrical work.

25. Cell Potential
Cell Potential/Electromotive Force (EMF): the “push/pull” or driving force on electrons which is measured in voltage (potential difference).
Work is measured in Joules and charge is measured in Coulombs

26. Charge
Charge is symbolized by “q” and is related to the number of moles of electrons.
The charge of one mole of electrons is a constant called the faraday constant (F).
F = 96,485 C/mol e-
q = nF
Thus… w = -qEcell = -nFEcell

27. Free Energy Previously we learned that …
ΔG = energy available to do work
Thus ΔG = -nFEcell
For standard conditions: ΔG◦ = -nFE◦cell
Why does this matter?
Because now we have a way to experimentally determine the change in free energy of a reaction.

28. Cell Potential and Free Energy
E◦cell > 0 ΔG◦ < 0 Spontaneous
E◦cell < 0 ΔG◦ > 0 Not Spontaneous
E◦cell = 0 ΔG◦ = 0 At Equilibrium

29. Cell Potential and Equilibrium
Under standard conditions and at equilibrium:
ΔG◦ = 0, K = Q, Temp=298K

30. Nernst Equation Under nonstandard conditions:
As the products increase, Q increases thus Ecell decreases
As the products decrease, Q decreases thus Ecell increases

31. Concentration Cells
Concentration cells have the same species present, but at different concentrations.
As soon as you “hook up” a galvanic cell it becomes a concentration cell that runs until ΔG = 0.

32. Concentration Cell
The cell will run until the solutions on both sides are in equilibrium.

33. Homework Pg
# 37, 41, 51, 52
If you are up to a challenge: 38

34. Electrolytic Cell
Electrolysis: forcing a current through a cell to produce a chemical change for which the cell potential is negative.
Nonspontaneous based on the electrical work.
Plating: depositing neutral metal on the electrode by reducing the metal ions in solution.

35. Stoichiometry of Electrolytic Cells
We can convert between the following:
Current and time
Ampheres (C/s)
Quantity of charge in coulombs
Moles of electrons
Moles of species
Mass of species.

37. Homework Pg. 834 Review Electrochemistry: https://youtu.be/IV4IUsholjg
# 79, 87, 88
Review Electrochemistry: