ELECTROCHEMISTRY
Review

Electrochemistry – the study of the interchange of chemical
and electrical energy

Oxidation–reduction (redox) reaction – involves a transfer of
electrons from the reducing agent to the oxidizing agent

Oxidation – loss of electrons

Reduction – gain of electrons

Reducing agent – electron donor

Oxidizing agent – electron acceptor

Balance redox reactions
Applications
 Moving electrons is electric current.
 8H++MnO4-+ 5Fe+2 +5e



→ Mn+2 + 5Fe+3 +4H2O
Helps to break the reactions into half
reactions.
8H++MnO4-+5e- → Mn+2 +4H2O
5(Fe+2 → Fe+3 + e- )
In the same mixture it happens without doing
useful work, but if separate
Schematic of a method to separate the oxidizing and reducing agents of
aredox reaction. (The solutions also contain counter ions to balance the
charge.)
Galvanic Cell
Galvanic cells can contain a salt bridge or a porous-disk connection .
A salt bridge contains a strong electrolyte held in a Jello-like matrix.
A porous disk contains tiny passages that allow hindered flow of ions.
Galvanic Cell
 Device in which chemical energy is changed to electrical
energy.
 Uses a spontaneous redox reaction to produce a current
that can be used to do work.
 Oxidation occurs at the anode.
 Reduction occurs at the cathode.
Galvanic Cell
Digital voltmeters draw only a negligible current and are convenient to use.
Cell Potential
 Oxidizing agent pushes the electron.
 Reducing agent pulls the electron.
 The push or pull (“driving force”) is called the
cell potential Ecell
 Also called the electromotive force (emf)
 Unit is the volt(V)
 = 1 joule of work/coulomb of charge
 Measured with a voltmeter
Standard Reduction Potentials
Standard Hydrogen Electrode
 This is the reference all
other oxidations are
compared to
H2 in
 Eº = 0
 º indicates standard
+
H
states of 25ºC,
1
Cl
atm, 1 M solutions.
1 M HCl
Standard Reduction Potentials




All half-reactions are given as reduction processes in standard
tables. [1 M, 1atm, 25°C]
When a half-reaction is reversed, the sign of E° is reversed.
When a half-reaction is multiplied by an integer, E° remains the
same.
A galvanic cell runs spontaneously in the direction that gives a
positive value for E°cell.
•
E0 is for the reaction as written
•
The more positive E0 the greater the
tendency for the substance to be
reduced
•
The more negative E0 the greater the
tendency for the substance to be
oxidized
•
Under standard-state conditions, any
species on the left of a given halfreaction will react spontaneously with
a species that appears on the right of
any half-reaction located below it in the
table
(the diagonal rule)
•
The half-cell reactions are reversible
•
The sign of E0 changes when the
reaction is reversed
•
Changing the stoichiometric
coefficients of a half-cell reaction does
not change the value of E0
Example: Fe3+(aq) + Cu(s) → Cu2+(aq) + Fe2+(aq)



Half-Reactions:
 Fe3+ + e– → Fe2+
E° = 0.77 V
 Cu2+ + 2e– → Cu
E° = 0.34 V
To balance the cell reaction and calculate the cell potential, we
must reverse reaction 2.
 Cu → Cu2+ + 2e–
– E° = – 0.34 V
Each Cu atom produces two electrons but each Fe3+ ion accepts
only one electron, therefore reaction 1 must be multiplied by 2.
 2Fe3+ + 2e– → 2Fe2+ E° = 0.77 V
Overall Balanced Cell Reaction
2Fe3+ + 2e– → 2Fe2+
E° = 0.77 V (cathode)
Cu → Cu2+ + 2e–
– E° = – 0.34 V (anode)
 Balanced Cell Reaction:
Cu + 2Fe3+ → Cu2+ + 2Fe2+
 Cell Potential:
E°cell = E°(cathode) – E°(anode)
E°cell = 0.77 V – 0.34 V = 0.43 V
A galvanic cell involving
the half reactions
Zn → Zn2+ + 2e- (anode)
and Cu2+ + 2e- →Cu (cathode)
with E ocell = 1.10 V.
Line Notation







solidAqueousAqueoussolid
Anode on the leftCathode on the right
Single line different phases.
Double line porous disk or salt bridge.
If all the substances on one side are aqueous,
a platinum electrode is indicated.
For the last reaction
Cu(s)Cu+2(aq)Fe+2(aq),Fe+3(aq)Pt(s)
Galvanic Cell


1)
2)
3)
4)
The reaction always runs spontaneously in
the direction that produced a positive cell
potential.
Four things for a complete description.
Cell Potential
Direction of flow
Designation of anode and cathode
Nature of all the components- electrodes
and ions
Practice
What is the standard emf of an electrochemical cell
made of a Cd electrode in a 1.0 M Cd(NO3)2 solution
and a Cr electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e- →
Cr3+ (aq) + 3e-
Cd (s) E0 = -0.40 V
→
Anode (oxidation):
Cathode (reduction):
Cr (s)
Cr (s)
→
E0 = -0.74 V
Cr3+ (1 M) + 3e-
2e- + Cd2+ (1 M)
→ Cd (s)
0
0
E0 = Ecathode
- Eanode
cell
E0 = -0.40 – (-0.74)
Cd is the stronger
oxidizer
Cd will oxidize Cr
x2
x3
Potential, Work and ∆G





emf = potential (V) = work (J) / Charge(C)
E = work done by system / charge
E = -w/q
Charge is measured in coulombs.
-w = qE
 Faraday = 96,485 C/mol e q = nF = moles of e- x charge/mole e w = -qE = -nFE = ∆G
Potential, Work and ∆G






∆Gº = -nFE º
if E º < 0, then ∆Gº > 0 spontaneous
if E º > 0, then ∆Gº < 0 nonspontaneous
In fact, reverse is spontaneous.
Calculate ∆Gº for the following reaction:
Cu+2(aq)+ Fe(s) → Cu(s)+ Fe+2(aq)
 Fe+2(aq) + e-→ Fe(s)
 Cu+2(aq)+2e- → Cu(s)
Eº = 0.44 V
Eº = 0.34 V
Cell Potential and
Concentration
 Qualitatively - Can predict direction of change
in E from LeChâtelier.
 2Al(s) + 3Mn+2(aq) → 2Al+3(aq) + 3Mn(s)
 Predict if Ecell will be greater or less than Eºcell
if [Al+3] = 1.5 M and [Mn+2] = 1.0 M
 if [Al+3] = 1.0 M and [Mn+2] = 1.5M
 if [Al+3] = 1.5 M and [Mn+2] = 1.5 M
The Nernst Equation
It enables us to calculate E as a function of [reactants] and [products]
in a redox reaction.
 ∆G = ∆Gº +RTln(Q)
 -nFE = -nFEº + RTln(Q)
 E = Eo + (RT/nF) ln (Q)
 2Al(s) + 3Mn+2(aq) → 2Al+3(aq) + 3Mn(s)
Eº = 0.48 V
 Always have to figure out n by balancing.
 If concentration can gives voltage, then from
voltage we can tell concentration.
Spontaneity of Redox Reactions
∆G = -nFEcell
∆G0
n = number of moles of electrons in reaction
J
= -nFEcell0
F = 96,500
∆G0 = -RT ln K
RT
0 =
Ecell
nF
V • mol
= 96,500 C/mol
0
= -nFEcell
(8.314 J/K•mol)(298 K)
ln K =
n (96,500 J/V•mol)
0 =
Ecell
0.0257 V
n
ln K
0 =
Ecell
0.0592 V
n
log K
ln K
Spontaneity of Redox Reactions
If you know one, you can calculate
the other…
If you know K, you can calculate
∆Eº and ∆Gº
If you know ∆Eº, you can calculate
∆Gº
Concentration Cells
Batteries
A Common Dry Cell Battery
Leclanché cell
Anode:
Zn (s)
Cathode: 2NH4+(aq) + 2MnO2 (s) + 2eZn (s) + 2NH4 (aq) + 2MnO2 (s)
Zn2+ (aq) + 2eMn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s
Batteries
A Mercury Battery
Anode:
Cathode:
Zn(Hg) + 2OH- (aq)
HgO (s) + H2O (l) + 2eZn(Hg) + HgO (s)
ZnO (s) + H2O (l) + 2eHg (l) + 2OH- (aq)
ZnO (s) + Hg (l)
Batteries
Lead storage battery
Anode:
Cathode:
Pb (s) + SO2-4 (aq)
PbSO4 (s) + 2e-
PbO2 (s) + 4H+ (aq) + SO2-4 (aq) + 2e-
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq)
PbSO4 (s) + 2H2O (l)
2PbSO4 (s) + 2H2O (l)
Batteries
Solid State Lithium Battery
Corrosion
 Rusting - spontaneous oxidation.
 Most structural metals have reduction




potentials that are less positive than O2 .
Fe → Fe+2 +2eEº= 0.44 V
O2 + 2H2O + 4e- → 4OH- Eº= 0.40 V
Fe+2 + O2 + H2O → Fe2 O3 + H+
Reaction happens in two places.
Salt speeds up process by increasing
conductivity
Water
Rust
eIron Dissolves- Fe → Fe+2
Preventing Corrosion
 Coating to keep out air and water.
 Galvanizing - Putting on a zinc coat
 Has a lower reduction potential, so it is more.
easily oxidized.
 Alloying with metals that form oxide coats.
 Cathodic Protection - Attaching large pieces
of an active metal like magnesium that get
oxidized instead.
Electrolysis
 Running a galvanic cell backwards.
 Put a voltage bigger than the potential and
reverse the direction of the redox reaction.
 Used for electroplating.
1.10
e-
e-
Zn
Cu
1.0 M
Zn+2
Anode
1.0 M
Cu+2
Cathode
e-
A battery
>1.10V
Zn
e-
Cu
1.0 M
Zn+2
Cathode
1.0 M
Cu+2
Anode
Electrolysis and Mass Changes
charge (C) = current (A) x time (s)
1 mole e- = 96,500 C
So what is the charge on a single electron?
Other uses
 Electroysis of water.
 Seperating mixtures of ions.
 More positive reduction potential means the
reaction proceeds forward.
 We want the reverse.
 Most negative reduction potential is easiest
to plate out of solution.