Ch. 20 Notes
Voltaic Cells (Galvanic cells) – electron transfer takes place through an external pathway.
*a spontaneous redox reaction then releases energy that can be used to do work.
External circuit or electrodes connect two metals
*anode: oxidation ½ rxn takes place (so electrons become available)
*cathode: reduction ½ rxn takes place (where electrons are consumed)
Explanation of Voltaic cell Electrons become available at the anode while Zn is oxidized to Zn2+
 Electrons flow through the external circuit (and voltmeter) to cathode (Cu metal) where
they are consumed as Cu2+ is reduced to Cu solid.
Observations:
 Zn electrode loses mass
 [Zn2+] increases
Creates electrical (charge) imbalance which is restored
 Cu electrode gains mass
through use of a salt bridge or porous barrier.
2+
 [Cu ] decreases
Salt bridge- a gel-like electrolyte substance (NaNO3) that allows the migration of charged
particles to neutralize the building charge of the cell.
 Composition of ions must differ from those in the voltaic cell.
 Anions migrate toward the anode (NO3- and Zn2+)
 Cations migrate toward the cathode (Na+ and Cu)
Cell Electromotive force (EMF)Electrons flow from the anode  cathode due to a difference in P.E.
 Electrons at the anode have greater P.E. than in the cathode so spontaneous flow occurs
through the external circuit.
 The measured difference in P.E. = volts(V) or J/C
o 1J=1VK
o 1 C = 1 A x 1s
o 1 V is the P.E. difference needed to yield 1 J with an electric charge of 1C.
Cell EMF- (Ecell) the cell potential measured in volts or also called cell voltage.
 If the rxn is spontaneous Ecell will be positive.
 Depends on:
o Specific reactions at the anode and cathode
o Concentrations of the reactants and products
o Temperature – assumed to be 25oC
To begin, work w/standard emf or standard cell potentials (Eocell).
Ex., Zn(s) + Cu2+(aq, 1M)  Zn2+ (aq, 1M) + Cu (s) Eocell = 1.10 V
Standard Reduction Potentials – Eored, used to calculate Eocell.
Eocell = Eored (cathode) – Eored (anode)
E.g., Using Eored, calculate the standard emf for the following cell rxn:
Cr2O72- (aq) + 14 H+ (aq) + 6 I- (aq)  2 Cr3+ (aq) + 3 I2 (s) + 7 H2O (l)
Because every voltaic cell involves 2 half cells, it’s not possible to directly measure Eored of a ½
reaction. So, to measure the Eored of a half rxn one must use a reference half reaction.
Reference half-rxn: 2H+ (aq, 1M) + 2e-  H2 (g, 1 atm) Eored = 0 V
 Called a standard hydrogen electrode (SHE)
e.g.,
So, Zn2+ (aq, 1 M) + 2 e-  Zn (s) Eored = - 0.76 V
(always write the reaction as a reduction reaction although it may be an oxidation
reaction in the cell.
Eored
Of other half reactions have been calculated and are found in tables to allow easy calc. of Eocell.
Because Eocell measures P.E./electrical charge, Eored are intensive so changing the coefficient in a
half reaction does NOT affect the value of Eored . Therefore, Eored for 1 mol Zn = 2 mol Zn
*The more positive the Eored , the
greater the driving force for
reduction.
*To occur spontaneously, Eored
(cathode) must be more than the Eored
(anode)
*** Eocell is the “excess” driving force than can be used to push e- through an external circuit.
Oxidizing and Reducing Agents –
The more positive the Eored for a half reaction, the greater the tendency for the reactant of the half
reaction to be reduced and therefore, to oxidize another species.
e.g., F2 (g) + 2e-  2 F- (aq) Eored = 2.87 V
* most easily reduced so strongest oxidizing agent
Most frequently used oxidizing agents are the halogens, O2, and oxyanions (MnO4-, Cr2O72-,
NO3-).
Li+ (aq) + e-  Li (s) Eored = 3.05 V
*most difficult species to reduce so it is the poorest oxidizing agent but one of the strongest
reducing agents.
Other good reducing agents include H2, active metals and alkaline earth metals. Reducing agents
are difficult to store because O2 is a good oxidizing agent.
Spontaneity of Redox ReactionsAny reaction that can occur in a voltaic cell to produce a positive emf must be spontaneous.
In general = Eo = Eored (reduction process) – Eored (oxidation process)
 Positive E = spontaneous
 Negative E = nonspontaneous
Ex., Using a list of Eored , determine whether the following reactions are spontaneous under
standard conditions.
a. Cu(s) + 2 H+ (aq)  Cu2+ (aq) + H2 (g)
b. Cl2 (g0 + 2 I- (aq)  2 Cl- (aq) + I2 (s)
EMF and Free-Energy ChangeG is a measure of the spontaneity of a process that occurs at constant T & P. Because emf
indicates whether the reaction is spontaneous the relationship is:
n= number of moles of e- transferred
F= Faraday’s constant, 1 F = 96,500 C/mole e1 F = 96,500 J/V mole eA positive value of E and a negative G = spontaneous
G = - nFE or

Go = - nFEo
E.g., Use Eored to calculate the
Go for:
2 Br- (aq) + F2 (g)  Br2 (l) +2 F- (aq)
The equation,
G = - nFE , can also be used to calculate how long a cell will be active
because 1 C = 1 A-s.
E.g., The reduction of TiO2+ to Ti3+ can be carried out electrochemically. What is the minimum
time, in seconds, required to reduce 50.000 mL sample of 0.115 M TiO2+ with a direct current of
1.06 amperes?
Effect of Concentration on cell EMFAs a cell is discharged, emf falls until E = 0, so the cell is dead.
 Emf depends on the concentrations of the R and P.
o When [R] increases, emf increases
o When [P] increases, emf decreases
The emf generated under nonstandard conditions is calculated using the Nernst Equation.
E.g., Calculate the emf generated by a cell when [Cr2O72-] = 2.0 M, [H+] = 1.0 M, [I-] = 1.0 M
and [Cr3+] = 1.0 x 10-5M.
Cr2O72- (aq) + 14 H+ (aq) + 6 I- (aq)  2 Cr3+ (aq) + 3 I2 (s) + 7 H2O (l)
1. Calculate Eo and Q
2. Determine the number of moles in the reaction
3. Plug values in and calculate
The Nernst equation can also be used to calculate the [R] or [P] (e.g., [H+] and also pH)
E.g,. If the voltage of a Zn/H+ cell is 0.45 V at 25oC when [Zn2+] = 1.0 M and PH2 = 1.0 atm,
what is the [H+]?
Equilibrium Constants for Redox ReactionsAt equilibrium, when G = 0 and E = 0 (Q=K)
Inserting E = 0 and Q = K into the Nernst equation:
0 = Eo – 0.592 log K
n
OR
Log K = nEo
0.0592
So, K can be obtained from the Eo for a reaction.
E.g., Using a list of Eored, calculate K for the oxidation of Fe2+ by O2 in acidic solution, according
to the following reaction:
O2 (g) + 4 H+ (aq) + 4 Fe2+ (aq)  4 Fe3+ (aq) + 2 H2O (l)
Electrolysis – forcing a current through a cell to produce a chemical change for which the cell
potential (E) is negative.
 Consists of 2 electrodes in a molten salt (used in the production of Na or AL) or solution
 Use battery of alternative electrical current to push the e-. The negative terminal connects
to the cathode (receives the e-) and the positive terminal connects to the anode (pulls eoff)
 Use stoichiometry to determine how much chemical change occurs with the flow of a
given current for a specified time.
o Current and time (A) and (s)
o Quantity of charge (C)
o Mole of e- = Faraday’s constant 96,500 C/mole eo Mole of substance = ox or red
o Grams of substance = ox or red
E.g., What mass of copper is plated out when a current of 10.0 amps is passed for 30.0 min?
Electrolysis of waterAnode –
Cathode –
However, so few ions exist in pure water that only a small current actually flows. The addition of
a soluble salt is needed to facilitate the process.
Electrolysis of Mixtures of Ions=
If a solution contains multi0ple ion is solution, then use Eo values to determine the order in
which the metals plate out.
Cu2+, Ag+, Zn2+
0.34 V 0.80 V -0.76V
Reactions want to run spontaneously so Eored that is most positive goes first.
Ag+ > Cu2+ > Zn2+
e.g.,
Calc. Eocell
When adding two half reactions, Eored the one that is less positive is reversed b/c the Eocell is less
positive (therefore spontaneous) for a voltaic cell.
e.g.,