1. Electrochemistry
Dr. Susan Lagrone

2. ASSIGNING OXIDATION NUMBERS
1. F is always Elements in standard state = zero 3. Monatomic ions = charge 4. H = +1 with nonmetals (H= -1 with metals) 5. O = -2 usually, but in peroxide = -1, superoxides it can even be a fraction or with F it is positive! 6. Memorize polyatomic ions

3. OXIDATION – REDUCTION (REDOX)
DEFINITIONS: Oxidation is loss of eletrons oxidation number increases (more positive)
Reduction is gain of eletrons
oxidation number decreases (more negative)

4. OXIDATION – REDUCTION (REDOX)
Given the reaction below, identify
Species oxidized
Species reduced
2 H2O2 (aq) → O2 (g) + 2 H2O(l),
OIL
RIG
ASSIGN
OX #
H = +1
H = +1
O = 0
O = -1
O = -2

5. OXIDATION – REDUCTION (REDOX)
Given the reaction below, identify
Species oxidized
Species reduced
2 H2O2 (aq) → O2 (g) + 2 H2O(l),
H2O2 oxygen -1 to zero
H2O2 oxygen -1 to -2
H = +1
O = 0
H = +1
ASSIGN
OX #
O = -1
O = -2

6. OXIDATION – REDUCTION (REDOX)
FUN FACT: Disproportionation - the process by which the same substance is both OXIDIZED AND REDUCED 2 H2O2 (aq) → O2 (g) + 2 H2O(l),
O = 0
O = -1
O = -2

7. OXIDATION – REDUCTION (REDOX)
DEFINITIONS: Oxidizer is the substance that causes oxidation to occur, i.e. is reduced (sometimes called oxidizing agent)
Reducer is the substance that causes reduction to occur,
i.e. is oxidized (sometimes called reducing agent)

8. OXIDATION – REDUCTION (REDOX)
Given the reaction below, identify
Oxidizer
Reducer
Al(NO3)3 (aq) + Mg (s) → Al (s) + Mg(NO3)2(aq)
Reduced, Al3+ (aq)
Oxidized, Mg (s)
Al = +3
Al = 0
ASSIGN
OX #
O = -2
N = +5
Mg = 0
Mg = +2
O = -2
N = +5

9. BALANCING REDOX RXNS BY HALF RXNS
Assign oxidation numbers to split reaction into oxidation half-reaction and reduction half-reaction
Balance the atom oxidized /reduced and do not change these coefficients
Balance oxygen with water
Balance hydrogen with H+ ions
Add electrons to the more positive side to balance the charge
All reactions are balanced by atoms AND charge

10. BALANCING REDOX RXNS BY HALF RXNS
Assign oxidation numbers to split reaction into oxidation half- reaction and reduction half-reaction
Balance the atom oxidized /reduced and do not change these coefficients
Balance oxygen with water
Balance hydrogen with H+ ions
Add electrons to the more positive side to balance the charge
The result is balanced in ACID
CHECK:
All reactions are balanced by atoms AND charge

11. BALANCING in BASE – ADD THESE THREE STEPS
6. Add OH- for each H+ to BOTH SIDES OF EQUATION 7. Combine H+ and OH- appearing on same side to make WATER. 8. If possible SUBTRACT WATER FROM BOTH SIDES (cancel waters so it appears only on one side)
CHECK:
All reactions are balanced by atoms AND charge

12. BALANCING REDOX RXNS BY HALF RXNS
YOU TRY in acid:
Cr(s) + NO3- (aq)  Cr3+ (aq) + NO (g)
CH3OH(aq) + Ce4+ (aq)  CO2 (aq) + Ce3+ (aq)
Fe2+ (aq) + MnO4- (aq)  Mn2+ (aq) + Fe3+ (aq)
YOU TRY in base:
4. MnO4- (aq) + ClO4- (aq)  Mn2+ (aq) + ClO3- (aq)

13. Voltaic Cell (Galvanic) – it is a battery Ecell > 0 (positive)
SPONTANEOUS CHEMICAL RXN
Thermodynamically favorable
∆G < 0 (negative)
Keq > 1 (rxn is product favored)

14. Voltaic Cell (Galvanic) – it is a battery Ecell is positive
Anode = oxidation ( an ox); lose eletrons
Cathode = reduction (redc at); gain electrons
e- flow from anode to cathode (FATCAT)
Salt bridge- maintains electrical neutrality
(keeps charge from building up in the cell)
Voltmeter – measures cell potential
Ammeter – measures current
NOTICE: No battery or power source used
Line notation – anode is written first
Zn(s)  Zn2+ (1.0 M)  Cu2+ (1.0M) Cu
Bar denotes change in phase of matter
Double bar denotes salt bridge (porous disk)
Standard conitions are 1.0 M , 1 atm , 25°C

15. Voltaic Cell (Galvanic) – it is a battery Ecell is positive
Question:
1. Which electrode will decrease in mass as the
cell operats?
2. Write the cell rxn which occurs at the cathode.
3. Calculate standard cell potential.
Reduction potentials:
Zn2+ Zn(s) = -0.76V
Cu2+ Cu = 0.34V
4. Which ions flow into the cell containing the
copper electroe?
NOTICE: No battery or power source used

16. The hydrogen electrode
2H+ (aq) + 2 e-  H2 (g) Defined: E0 = 0.00V
All reduction potentials are measured relative to this electrode.

17. Problem 5:
Sketch the galvanic cell given the half cell reactions below. Mn e-  Mn(s) E0 = -1.18V Fe e-  Fe(s) E0 = V

18. Indicate electron flow in cell Show ion movement through bridge
Problem 5: continued
Label the anode
Indicate electron flow in cell
Show ion movement through bridge
Write the line notation for this cell
Calculate ∆G0 for the cell

19. 2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s)
Problem 6:
2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s)
Respond to the following statements and questions that relate to the species and the reaction represented above.
Write the complete electron configuration (e.g., 1s2 2s ) for Zn2+.
Which species, Zn or Zn2+, has the greater ionization energy? Justify your answer.
Identify the species that is oxidized in the reaction

20. 2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s)
Problem 7:
2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s)
Respond to the following statements and questions that relate to the species and the reaction represented above.
Indicate the movement of ions in the salt bridge.
Label the cathode.
Indicate electron flow in the cell.
Calulate standard cell potential, E0.
Which electrode gains mass during cell operation
If the concentration of zinc nitrate changes to 0.50M
will the cell potential increase, decrease, or remain the same? Justify your answer.

21. Problem 8:
Place the following in order of increasing oxidizing strength. Use table on next slide. (all under standard conditions) Cadmium(II), iodate ion, potassium ion, water, iodine, and AuCl4-

23. Problem 9: Answer the following questions using the table on previous slide.
Is H+ (aq) capable of oxidizing Cu(s) to Cu2+ (aq)?
Is Fe+3 (aq) capable of oxidizing I- (aq)?
Is H2 (g) capable of reducing Ag+ (aq)?
Is Fe+2 (aq) capable of reducing Cr3+ (aq) to Cr2+ (aq)?

24. Problem 10: Consider the species at standard conditions
Problem 10: Consider the species at standard conditions. Na+, Cl-, Ag+, Ag, Zn2+, Zn, Pb
Which is the strongest oxidizing agent? (oxidizer)
Which is th strongrest reducing agent?
Which species can be oxidized by sulfate ion, SO42-, in acid?
Which species can be reduced by aluminum solid?

25. Problem 11:
Use the half reactions below to explain why household bleach ( a highly alkaline solution of hypochlorite) should not be mixed with household ammonia or glass cleansers containing ammonia. ClO- + H2O + 2 e-  2 OH- + Cl- E0 = 0.90V N2H4 + 2 H2O + 2 e-  2 OH- + 2 NH3 E0 = -0.10V