Chem 1721
Brief Notes: Chapter 19
Chapter 19: Electrochemistry
Consider the same redox reaction set up 2 different ways:

Cu metal in a solution of AgNO3
electrically conducting wire
Cu
salt bridge
Cu
Ag
Ag+
NO3 -
Cu2+
Anode
Ag+
Cathode
What is the reaction?
In the experiment on the left:
over time the strip of copper metal decreases in size (and mass)
the solution starts clear and colorless, and over time becomes clear and light blue
over time a fine grayish-silver powder deposits on the bottom of the beaker
ox: Cu (s)  Cu2+ (aq) + 2 e
red: Ag+ (aq) + 1 e  Ag (s)
net: Cu + 2 Ag+  Cu2+ + 2 Ag
⎯
⎯
energy change? Any energy associated with the process goes into changing the temperature of the solution.
In the experiment on the right:
the Cu/Cu2+ redox couple and the Ag+/Ag redox couple are separated
Cu metal and Ag metal are connected by an electrically conducting wire
e transfer from Cu to Ag+ occurs through the wire generating a voltage
this is an electrochemical cell
⎯

there are two types of electochemical cells
1.
Galvanic cell: a spontaneous chemical reaction occurs that generates a voltage
2.
Eletrolytic cell: a nonspontaneous chemical reaction is driven by an applied current
Galvanic cells

2 compartments: anode and cathode
anode – where oxidation occurs; usually shown on the left
cathode – where reduction occurs; usually shown on the right
can write anode and cathode half reactions (parallel to oxidation and reduction half reactions respectively)
anode:
cathode:
net cell:
Cu  Cu2+ + 2 e
Ag+ + 1 e  Ag
Cu + 2 Ag+  Cu2+ + 2 Ag
⎯
⎯

each compartment must contain an electrically conducting solid; the electrode; where the wire connects
in the example above: copper is the anode, silver is the cathode

the direction of electron flow in a Galvanic cell is always anode  cathode
electrons are generated at the anode (where oxidation occurs)
electrons are consumed at the cathode (where reduction occurs)


battery designations:
anode designated ⎯ because electrons are produced
cathode designated + because electrons are consumed
compartments are connected by a salt bridge (or porous membrane – see the 2 figures below)
electrically conducting wire
electrically conducting wire
salt bridge
porous or semi-permeable membrane
Cathode
Anode
Anode
Cathode
salt bridge (or porous membrane) allows migration of spectator ions between compartments
maintain charge neutrality
anions flow toward the anode; e ’s leave the anode ∴ + charge builds up; i.e. Cu  Cu2+
cations flow toward the cathode; e ’s arrive at the cathode ∴ positive charge is decreasing; i.e. Ag+  Ag
⎯
⎯

short-hand description of a Galvanic cell – line notation
basically: anode | anode compartment || cathode compartment | cathode
for the Cu/Ag+ cell: Cu (s) | Cu2+ (aq) || Ag+ (aq) | Ag (s)

ex. Consider a Galvanic cell with the following cell reaction: Fe (s) + 2 Fe3+ (aq)  2 Fe2+ (aq). Write the anode
and cathode half reactions, sketch the set-up of this cell, and write the corresponding line notation.
anode:
cathode:
net cell:
Fe  Fe2+ + 2 e
Fe3+ + e  Fe2+
Fe + 2 Fe3+  3 Fe2+
⎯
⎯
set up:
note: the cathode compartment requires a chemically inert conductor to facilitate the e transfer
common chemically inert conductors include Pt (s), C (s) i.e a graphite rod
⎯
line notation: Fe (s) | Fe2+ (aq) || Fe2+ (aq), Fe3+ (aq) | Pt (s) or other conductor
Cell voltage or cell potential; E and E°

a Galvanic cell produces a particular voltage depending on the exact combination of half reactions (half cells),
temperature, pressures of gases, and molar concentration of solutions
cell potential, E, measured in volts (V)
1 volt = difference in potential between 2 points
1 J of work is done when 1 Coulomb of charge moves between 2 points (of higher potential to lower potential)
differing by 1V
1 V = 1 J/C


standard cell potential, E°
cell potential (in V) when solids and liquids are in their pure form; gasses are at 1 atm pressure; solutions at 1 M
concentration
E°cell = E°ox + E°red = E°anode + E°cathode
one standard half-cell is designated as a reference and assigned E° = 0.00 V
standard hydrogen electrode: 2 H+ (aq) + 2 e  H2 (g); E° = 0.00 V
specifically with: [H+] = 1.00 M, PH2 = 1.00 atm; temp = 298 K
⎯

consider the following Galvanic cell:
anode:
Zn  Zn2+ + 2 e
cathode: 2 H+ + 2 e  H2
net cell: Zn + 2 H+  Zn2+ + H2; E°cell = 0.76 V
⎯
⎯
if E°cathode = 0.00 V then E°anode = 0.76 V
now consider:
anode:
Zn  Zn2+ + 2 e ; E° = 0.76 V
cathode: Cu2+ + 2 e  Cu; E° = ????
net cell: Zn + Cu2+  Zn2+ + Cu; E°cell = 1.10 V
⎯
⎯
if E°cell = E°anode + E°cathode then E°cathode = E°cell ⎯ E°anode = 0.34 V
Tabulated Standard Reduction Potentials (Appendix I and Table 19.1)

all half-reactions written as reductions
i.e. oxidizing agent + e  reducing agent
⎯

tabulated in order of increasing E°
increasing reactivity in direction written
increasing tendency for reduction to occur
decreasing tendency for oxidation (reverse reaction) to occur

stronger reducing agents at top of table; stronger oxidizing agents at bottom of table

+ value of E° indicates a spontaneous half reaction; ⎯ value of E° indicates a nonspontaneous half reaction

an electrochemical cell needs a reduction and an oxidation ∴ 2 half cells
the equation corresponding to the oxidation (anode) half cell must be written in reverse of table
when the half-reaction is reversed the sign of E° is also reversed
when a half reaction is multiplied by some factor to change the stoichiometry you do NOTHING to E°
Spontaneous cell reactions and E°cell
 a spontaneous reaction has a + E°cell
 Galvanic cells involve spontaneous chemical change ∴ E°cell must be +
ex.
consider Galvanic cells based on:
Li+ + e  Li
E° = ⎯3.04 V
Cl2 + 2 e  2 Cl
E° = 1.36 V
AND
Pb2+ + 2 e  Pb
E° = ⎯.13 V
PbO2 + SO42 + 4H+ + 2 e  PbSO4 + 2 H2O
⎯
⎯
⎯
⎯
⎯

⎯
E° = 1.69 V
the position of the half-reaction in the Table of Standard Reduction Potentials allows you to predict the anode and
cathode of any pair of half cells
the half reaction that is lower in the table (larger E°) will be the cathode (reduction) half reaction
the half reaction that is higher in the table (smaller E°) will be the anode (oxidation) half reaction
remember: must end up with a + E°cell for a Galvanic cell

relative strengths of oxidizing and reducing agents
the lower in the table the stronger the oxidizing agent
can cause any half reaction above it to proceed in reverse
ex.

F2 is the strongest oxidizing agent listed
Li is the strongest reducing agent listed
a related concept . . . the Activity Series of Metals
ranks metals as reducing agents
metals that like to be oxidized are good reducing agents
good reducing agents have small (and frequently negative) E°’s
Activity Series: Li > K > Ba > Ca > Mg > Be > Al > Zn > Fe > Cu > Ag > Au
most active metal

least active metal
best reducing agent

worst reducing agent on the list
a metal in the Activity Series can reduce any Mn+ to the right of it

lots of potential questions – several examples
ex.
NO3 + 4 H+ + 3 e  NO + 2 H2O
Fe3+ + e  Fe2+
If combined in a Galvanic cell, what will be the net cell reaction and E°cell?
Consider these half reactions:
⎯
⎯
⎯
ex.
Can MnO4 in acidic solution (i.e H+ present) oxidize Ni? Ag?
ex.
Which is a stronger reducing agent, Cr or Mn?
ex.
Based on the Acitivity Series, can aluminum metal reduce Ca2+? Ag+?
E° = 0.96 V
E° = 0.77 V
⎯
Complete description of a Galvanic cell
 at this point you should be able to
write/identify the anode and cathode half reaction of a cell
write the net cell reaction
determine E°cell, or E°anode or E°cathode if E°cell is given
sketch a diagram for a Galvanic cell including: identification of the anode and cathode compartments and their
components (i.e electrode and ions in solution); direction of electron flow; salt bridge or porous barrier and
ion migration
write the line notation that describes a Galvanic cell
Galvanic cells, work and free energy
 the amount of work that a Galvanic cell can do can be quantified
wmax = ⎯nFE°
wact = ⎯nFE
cell efficiency = (wact/wmax)*100
efficiency = (E/E°)*100
n = mol e ;
⎯
F = Faraday constant, 96485 C/mol e
⎯
nF = quantity of charge transferred

free energy change for a reaction in a Galvanic cell is equal to the maximum work that can be achieved
ΔG° = ⎯nFE°

ΔG = ⎯nFE
consider again the Cu/Ag Galvanic cell:
anode:
cathode:
net cell rxn:
Cu  Cu2+ + 2 e ;
2*(Ag+ + 1 e  Ag);
Cu + 2 Ag+  2 Ag + Cu2+;
⎯
⎯
E° = ⎯0.34 V
E° = 0.80 V
E°cell = 0.46 V
Calculate the maximum work that can be done by this cell, ΔG°, and K at 25°C. Recall, 1 V = 1 J/C.
wmax = ΔG° = ⎯nFE°
wmax = ΔG° = ⎯(2mol e )(96485 C/mol e )(0.46 J/C)
wmax = ΔG° = ⎯8.9 x 104 J, or ⎯89 kJ
ΔG° = ⎯RTlnK
ln K = (⎯8.9 x 104 J)/(⎯8.314 J•K 1•mol 1*298 K)
ln K = 36
K = 4.3 x 1015
⎯
⎯
⎯
⎯
Now consider that this cell operates at only 76% efficiency. Calculate E?
efficiency = (E/E°)*100
E = (effiency/100)*E°
E = (.76)(.46V) = 0.35V
and the actual work done by the cell . . .
wact = ⎯nFE
wact = ⎯(2mol e )(96485 C/mol e )(0.35 J/C)
wact = ⎯6.8 x 104 J or ⎯68 kJ
⎯
⎯
The Nernst Equation
 relates E to E°
 for cells operating at nonstandard conditions; [sol’n] ≠ 1.0 M and Pgas ≠ 1.0 atm

E = Eo

nF
E = Eo
lnQ
.02569 lnQ; at 25o C
n
for a system at equilibrium, E = 0; Q = K
Eo =

RT
RT
nF
*note – a “dead” battery is at equilibrium
Eo = .02569 lnK; at 25o C
n
lnK
Calculate E for a Galvanic cell at 25°C with a Ni cathode in 3.3 M NiCl2 (aq) and a Tl anode in 0.25 M TlNO3 (aq).
anode:
2(Tl  Tl+ + 1 e )
E° = 0.34 V
cathode:
Ni2+ + 2 e  Ni
E° = ⎯0.24V
as this cell runs, [Ni2+] will decrease (reactant)
2+
+
net cell:
2 Tl + Ni  Ni + 2 Tl
E°cell = 0.10 V
and [Tl+] will increase (product) until Q = K
⎯
⎯
E = 0.10 J/C
(8.314 J/Kmol)(298 K)
(2 mol e)(96485 C/mol e)
E = 0.15 V
ln *
0.252
3.3
Electrolytic cells

reactions in electrolytic cells are nonspontaneous; ΔG° is +; E°cell is ⎯

E°cell is the minimum voltage required for reaction to occur

you need to know:
balanced chemical equation to identify mole electrons transferred (n = mol e )
Faraday constant defining quantity of charge transferred per mole of electrons; F = 96485 C/mol e
current used or required during the electrolysis in amperes; 1 A = 1 C/s
electric power in Watts; 1 W = 1 J/s
⎯

⎯
some examples:
ex.
Determine the time required to plate 85.5 g Zn if 23.0 A passed through a solution of ZnSO4 (aq).
cathode rxn: Zn2+ + 2 e  Zn
⎯
logical pathway: g Zn  mol Zn  mol e  C  t
⎯
ex.
answer = 1.10 x 104 s or 183 min or 3.05 h
Determine the current required to plate 2.86 g chromium metal from CrCl3 (aq) in 2.5 min.
cathode rxn: Cr3+ + 3 e  Cr
⎯
logical pathway: g Cr  mol Cr  mol e  C; C÷s = A
⎯
answer = 106 A