Download CH 20

CHAPTER 20 – Electrochemistry (The chemistry of Oxidation-Reduction Reactions)
I.
A Review of Oxidation-Reduction (Redox) Reactions
A. Examples
H2 + 1/2 O2 → H2O
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) a molecular equation
Mg + 2 H+ → Mg2+ + H2 a net ionic equation
Is this net ionic equation balanced?
Why yes or no?
List the spectator ion(s) that were eliminated.
B. Terminology
1. In the equations above, what is oxidized?
What is reduced?
2. What is the oxidizing agent?
What is the reducing agent?
3. What is the reductant?
What is the oxidant?
II.
Balancing more complicated Redox reactions
A. Balancing by the Oxidation Number Method
Ag+(aq) + Mg(s) → Ag(s) + Mg+(aq)
Is the reaction above balanced?
Why?
Write the balanced equation:
In the space below, write the balanced equation for: Fe + Fe3+ → Fe2+
In the space below, write the balanced equation for the following reaction.
Use the oxidation number method to balance this equation.
I2O5 + CO → I2 + CO2
1
B. Balancing by the Half Reaction Method
1. Divide the unbalanced equation into two half reactions, an oxidation half
reaction and a reduction half reaction.
Fe + Fe3+ → Fe2+
oxid rxn:
Fe → Fe2+
balanced oxid ½ reaction =
red rxn: Fe3+ → Fe2+
balanced red ½ reaction =
2. Add the two half reactions together eliminating the electrons.
3. Balancing Redox Equations in Acidic Solutions – Sample Problem
In acidic solutions, use H+ ions whenever necessary to mass balance and
use H+ to eliminate OH− ions from the final balanced equation.
Balance the following equation using the half reaction method.
Br− + BrO− → Br2 (acidic solution) = Br− + BrO− + H+ → Br2 + H2O
4. Balancing Redox Equations in Basic Solutions - Sample Problem –
In basic solutions, use OH− ions whenever necessary to mass balance and
use OH− ions to eliminate H+ ions from the final balanced equation
Balance the following equation using the half reaction method.
Cl2 + S2O32−→ SO22− + Cl− (basic solution) =
Cl2 + S2O32− + OH−→ SO22− + Cl− + H2O
2
III.
Voltaic (Galvanic) Cells – Devices that use spontaneous Redox reactions to
generate current, a flow of electrons through an external circuit, or harness the
free energy change of a spontaneous redox reaction.
A. Zn + Cu2+ → Zn2+ + Cu
Is this reaction spontaneous?
Observation = chemical demonstration Yes!
Based upon the “Activity Series”? Chem 200 Yes!
Therefore there is a spontaneous transfer of electrons between Zn and Cu2+.
B. Voltaic (Galvanic) Cell Construction and Features (see packet page
1. Oxidation half cell: Zn → Zn2+ + 2 e− (This is the Anode)
Oxidation always occurs at the anode.
)
2. Reduction half cell: Cu2+ + 2 e− → Cu (This is the Cathode)
Reduction always occurs at the cathode.
3. The salt bridge maintains electrical neutrality in the cell. It allows for the
flow of ions to replace ions used in the Redox reaction.
4. The standard (25 oC, all ions at 1 M) cell voltage for this reaction is 1.10 V.
C. Eo = The Standard (conditions) Cell Voltage (Cell EMF)
A cell voltage (reaction) must simultaneously contain an oxidation and a
reduction reaction, these cannot be measured independent of the other.
Eocell = Eoox + Eored
For the reaction: Zn + Cu2+ → Zn2+ + Cu, Eo = 1.10 v. For this
reaction what are the values of Eoox and Eored?
D. Standard Reduction (Half-Cell) Potentials
By definition, Eored (H+ + e− → ½ H2) = 0.00 volts = Eoox (½ H2 → H+ + e−) .
All other potentials (voltages) are measured relative to this half reaction as the
arbitrary reference standard.
Zn + 2 H+ → Zn2+ + H2, Eo = 1.10 v
Eo = 0.76 v = Eoox + Eored = Eoox + 0.00 v
Eoox = 0.76 v − 0.00 v = 0.76 v (Zn → Zn2+ + 2 e−)
Eored = − Eoox
Eored (Zn2+ + 2 e− → Zn) = −0.76 v
3
E. Calculation of standard cell potentials using standard reduction potentials
from Brown’s Chemistry the Central Science.
Eocell = Eored(Cathode) − Eored(Anode) OR Eocell = Eoox + Eored
Calculate the standard cell voltage for a cell that uses the reaction,
2 Al(s) + 3 I2 → 2 Al3+ + 6 I−.
2 Al(s) → 2 Al3+ Eoox (= −Eored) = + 1.66 V ≠ 2(1.66 V)
3 I2 → 6 I− Eored = 0.54 V ≠ 3(0.54 V)
Eocell = (1.66 + 0.54) V = 2.20 V
A voltaic cell is constructed as follows: a strip of Al is inserted into a 1.0 M
solution of Al(NO3)3 is connected by a salt bridge to a chlorine gas electrode
(Cl2(g) at 1.0 atm and KCl solution at 1.0 M) and these cells are connected to
by external circuit to a voltmeter.
a. What is the standard emf (cell potential) for this cell?
b. Which half cell is the anode and the cathode?
c. Write a balanced equation for the overall reaction
d. Will the Al strip gain or lose mass?
G. Relative Strength of Oxidizing and Reducing Agents (Relative ability to be
oxidized or reduced).
Is Cl2 an oxidizing or reducing agent?
Under standard conditions, which is a stronger oxidizing agent, F2 or Cl2?
Under standard conditions, which is a stronger reducing agent, Al or Mg?
4
IV.
Determination of the Spontaneity of Redox Reactions
A. ΔGo = − nFEo where n = mol of e− transferred in the balanced equation, F =
Faraday’s constant = 9.56 x 104 C/mol = 9.56 x 104 J/V·mol and Eo = the
standard cell voltage (potential)
B. Sign Convention: Spontaneous - ΔGo is negative and Eo is positive
Non-Spontaneous - ΔGo is positive and Eo is negative
Equilibrium - ΔGo = Eo = 0
Sample Problem: The Eo for the reaction: Zn(s) + 2 Ag+(aq) → Zn2+(aq) + 2 Ag(S), is
1.56 V. Calculate the ΔGo for this reaction.
C. The Nernst Equation – If the concentrations of reactants and/or products in a
chemical cell are not standard, the cell voltage is not equal to Eo. The Nernst
Equation allows us to calculate the actual cell voltage or emf (E) under
nonstandard concentration conditions.
ΔG = ΔGo + (RT)lnQ
−nFE = −nFEo + (RT)lnQ
E = Eo + (RT/−nF)lnQ =>
at 25 oC E = Eo − 0.0592 V/n logQ
Sample Problem: The Eo for the reaction: Zn(s) + 2 Ag+(aq) → Zn2+(aq) + 2 Ag(S), is
1.56 V. Calculate the ΔGo for this reaction if the concentration of Ag+ is 0.0100
M and the concentration of Zn2+ ions is 2.00 M.
5
D. Calculation of E for summed reactions.
Given the following reduction potentials:
Sn2+ + 2 e− → Sn(s)
Sn2+ → Sn4+ + 2 e−
Eo = −0.136 V
Eo = −0.150 V
Calculate Eo for the half reaction: Sn4+ + 4 e− → Sn(s)
V.
Electrolytic Cells (Electrolysis) – Electrochemical cells that use an external power
supply or battery to carry out non-spontaneous chemical reactions by supplying
the necessary free energy.
A. Typical Cell Construction
B. Types of Electrolytic Cells
1. Ore Refinement or Pure Metal Production
The production of Mg(S) and Cl2(g) from the electrolysis of molten MgCl2
a. Calculate the minimum (threshold) voltage required to electrolyze
molten MgCl2.
b. Calculate the number of kilowatt-hours required to produce 1.00 kg of
Mg metal from molten MgCl2 using an applied voltage (emf) of 5.00
V assuming 80.0% efficiency of the reaction. Use as necessary the
standard reduction potentials supplied by the Brown text.
6
c. A direct current of 10.0 amperes (amps) is passed through a molten
solution of MgCl2 for 1.00 hours. How many coulombs of charge pass
through the cell? How Faradays are passed through the cell? How
many grams of Mg (or Cl2) will be produced, assuming 100 %
efficiency of the reaction?
C(coul) = I(amps) * time
Charge(C) = flow (C/s) * time(s)
2. Electroplating
When 15.0 amps of current is passed through aqueous solution of
platinum salt for 2.67 minutes, it is found that 1.22 grams of platinum
metal is plated out. Determine the oxidation state of the platinum ion
in this solution.
VI.
Corrosion and Cathodic Protection
7