Chemistry 100 – Chapter 20

Electrochemistry

Voltaic Cells

A Schematic Galvanic Cell

e-

Reducing Agent

e-

e-

Oxidizing Agent

Anode Cathode

Porous Disk

The Galvanic Cell Defined Galvanic cells – an electrochemical cell

that drives electrons through an external circuit as a result of the spontaneous redox reaction occurring inside.

The Zn/Cu Galvanic Cell

Voltaic Cells We expect the Zn electrode to lose mass

and the Cu electrode to gain mass. “Rules” of voltaic cells:

At the anode electrons are products. (Oxidation)

At the cathode electrons are reactants (Reduction)

Electrons flow from the anode to the cathode.

The Anode and Cathode Galvanic cells - the anode is

negative and the cathode is positive.

Electrons are made to flow through an external circuit. (Rule 3.)

Cell Potentials (Electromotive Force or EMF Values) Electromotive force (emf) - aka

the cell potential the force required to push electrons

through the external circuit. Ecell is the emf of a cell (old

notation). Now talk about the cell potential!

Cell Reactions The difference in the RHS and the LHS

reactionCu2+ (aq) + Zn (s) Cu (s) + Zn2+ (aq)

For each half reaction, we can write the reaction quotient (see Chapter 15) as followsCu2+ (aq) + 2 e- Cu (s) Q = 1/ [Cu2+] Zn2+ (aq) + 2 e- Zn (s) Q = 1/ [Zn2+]

Overall Qcell = [Zn2+] / [Cu2+]

The Cell Potential and G From the reaction Gibbs energy

cellcello

rxnrxn E F QlnRTGG

cellcell

orxnrxn E

F

QlnRT

F

G

F

G

The Nernst Equation

E - standard cell potential Cell potential under

standard conditions. [Solutes] = 1.000

mole/L T = 298.15 K P = 1.00 atm pressure

cellcell QlnF

RTEE

F

GE

orxn

cell

Cell Potentials

Standard Reduction Potentials We cannot measure the potential

of an individual half-cell! We assign a particular cell as being

our reference cell and then assign values to other electrodes on that basis.

Cell Potentials are Intensive Properties In the previous example, the cell

potential was simply the difference between the standard potential for the Sn4+/Sn2+ reduction and the Fe3+/Fe2+ reduction.

Reason: standard cell potentials are intensive quantities.

F

GE rxn

cell

F

GE

orxn

cell

[H+] = 1.00

H2 (g)

e-

Pt gauze

The Standard Hydrogen electrode Eo (H+/H2) half-cell = 0.000 V

p{H2(g)} = 1.00 atm

A Galvanic Cell With Zinc and the Standard Hydrogen Electrode.

Note - [Zn2+]= [H+] = 1.000 M

The Cell Equation for the Zinc-Standard Hydrogen Electrode.

The cell reaction 2 H+ (aq) + Zn (s) H2 (g) + Zn2+ (aq)

When we measure the potential of this cell

Ecell = ERHS - ELHS but ERHS = E(H+/H2) = 0.000 V Ecell = E(Zn2+/Zn) = -0.763 V

The Spontaneous Direction of a Cell reaction Examine the magnitude the of the

standard cell potential!

F

GE

orxn

cell

If Eo is positive, the rG is negative! Under standard conditions, the cell will proceed spontaneously in the direction written for the cell reaction.

The Composition Dependence of the Cell Potential

Nernst equation the nonstandard cell potential (Ecell) will

be a function of the concentrations of the species in the cell reaction.

cellcell QlnF

RTEE

To calculate Ecell, we must know the cell reaction and the value of Qcell.

Electrochemical Series Look at the following series of reactions Cu2+ (aq) + 2 e- Cu (s) E(Cu2+/Cu) = 0.337 VZn2+ (aq) + 2 e- Zn (s) E(Zn2+/Zn) = -0.763 V Zn has a thermodynamic tendency to

reduce Cu2+ (aq) Pb2+ (aq) + 2 e- Pb (s) E(Pb2+/Pb) = -0.13 VFe2+ (aq) + 2 e- Fe (s) E(Fe2+/Fe) = -0.44 V

Fe has a thermodynamic tendency to reduce Pb2+ (aq)

Differences in Reduction Potentials

• The larger the difference between Ered values, the larger Ecell.

• In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode).

Oxidizing and Reducing Agents The more positive Ered the stronger

the oxidizing agent on the left. The more negative Ered the

stronger the reducing agent on the right.

Spontaneous Oxidation Processes A species on the higher to the left of

the table of standard reduction potentials will spontaneously oxidize a species that is lower to the right in the table.

Any species on the right will spontaneously reduce anything that is higher to the left in the series.

Oxidizing and Reducing Agents

Concentration Cells Two identical half-cells.

RHS AgCl (s) + e- Ag (s) + Cl- (aq, 0.10 M)

LHS AgCl (s) + e- Ag (s) + Cl- (aq, 0.50 M)

Electrolyte concentration cell – the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells.

The Nernst equation for the cell

LHS

RHS

cellcell

]Cl[

]Cl[ln

F

RT

QlnF

RTE

Cells at Equilibrium When the electrochemical cell has

reached equilibrium

cellcellcell KQV 0E

Kcell = the equilibrium constant for the cell reaction.

RT

FE KlnKln

F

RTE cellcell

Knowing the E° value for the cell, we can estimate the equilibrium constant for the cell reaction.

Equilibrium Constants from Cell Potentials Examine the following cell.

Half-cell reactions. Sn4+ (aq) + 2 e- Sn2+ (aq) E(Sn4+/Sn2+) =

0.15 V Fe3+ (aq) + e- Fe2+ (aq) E (Fe3+/Fe2+) =

0.771 V Cell Reaction

Sn4+ (aq) + 2 Fe3+ (aq) Sn2+ (aq) + 2 Fe2+ (aq)

Ecell = (0.771 - 0.15 V) = 0.62 V

Lead-Acid Battery A 12 V car battery - 6 cathode/anode

pairs each producing 2 V.Cathode: PbO2 on a metal grid in sulfuric

acid:PbO2(s) + SO4

2-(aq) + 4H+(aq) + 2e- PbSO4(s) + 2H2O(l).

Anode: Pb:Pb(s) + SO4

2-(aq) PbSO4(s) + 2e-

Lead-Acid Battery The overall electrochemical reaction is PbO2(s) + Pb(s) + 2SO4

2-(aq) + 4H+(aq) 2PbSO4(s) + 2H2O(l)

for whichEcell = ERHS - ELHS

= (+1.685 V) - (-0.356 V)= +2.041 V.

Wood or glass-fiber spacers are used to prevent the electrodes form touching.

A Picture of a Car Battery

An Alkaline Battery Anode: Zn cap:

Zn(s) Zn2+(aq) + 2e- Cathode: MnO2, NH4Cl and carbon paste:

2 NH4+(aq) + 2 MnO2(s) + 2e- Mn2O3(s) +

2NH3(aq) + 2H2O(l) Graphite rod in the center - inert cathode. Alkaline battery, NH4Cl is replaced with KOH. Anode: Zn powder mixed in a gel:

The Alkaline Battery

Fuel Cells Direct production of electricity from

fuels occurs in a fuel cell. H2-O2 fuel cell was the primary source

of electricity on Apollo moon flights. Cathode: reduction of oxygen:

2 H2O(l) + O2(g) + 4e- 4OH-(aq) Anode:

2H2(g) + 4OH-(aq) 4H2O(l) + 4e-

Fuel Cells

Corrosion of Iron Since E(Fe2+/Fe) < E(O2/H2O) iron can

be oxidized by oxygen. Cathode

O2(g) + 4H+(aq) + 4e- 2H2O(l).

Anode Fe(s) Fe2+(aq) + 2e-.

Fe2+ initially formed – further oxidized to Fe3+ which forms rust, Fe2O3• xH2O(s).

Rusting (Corrosion) of Iron

Preventing the Corrosion of Iron Corrosion can be prevented by coating

the iron with paint or another metal. Galvanized iron - Fe is coated with Zn. Zn protects the iron (Zn - anode and Fe -

the cathode)

Zn2+(aq) +2e- Zn(s), E(Zn2+/Zn) = -0.76 VFe2+(aq) + 2e- Fe(s), E(Fe2+/Fe) = -0.44 V

Preventing the Corrosion of Iron

Preventing the Corrosion of Iron To protect underground pipelines,

a sacrificial anode is added. The water pipe - turned into the

cathode and an active metal is used as the sacrificial anode.

Mg is used as the sacrificial anode:

Mg2+(aq) +2e- Mg(s), E(Mg2+/Mg) = -2.37 VFe2+(aq) + 2e- Fe(s), E(Fe2+/Fe) = -0.44 V

Corrosion Prevention

Electrolysis of Aqueous Solutions Nonspontaneous reactions require

an external current in order to force the reaction to proceed.

Electrolysis reactions are non-spontaneous.

In voltaic and electrolytic cells: reduction occurs at the cathode, and oxidation occurs at the anode.

Voltaic vs.Electrolytic Cells Electrolytic cells – electrons are

forced to flow from the anode to cathode.

In electrolytic cells the anode is positive and the cathode is negative. (In galvanic cells the anode is negative and the cathode is positive.)

Electrolysis of Aqueous Solutions

Electrolysis of Molten Salts Decomposition of molten NaCl.

Cathode: 2Na+(l) + 2e- 2Na(l) Anode: 2Cl-(l) Cl2(g) + 2e-.

Industrially, electrolysis is used to produce metals like Al.

Electrolysis With Active Electrodes

Active electrodes: electrodes that take part in electrolysis.

Example: electrolytic plating.

Electrolysis With Active Electrodes (cont’d) Consider an active Ni electrode and

another metallic electrode placed in an aqueous solution of NiSO4: Anode: Ni(s) Ni2+(aq) + 2e-

Cathode: Ni2+(aq) + 2e- Ni(s). Ni plates on the inert electrode. Electroplating is important in

protecting objects from corrosion.

Quantitative Aspects of Electrolysis Consider the reduction of Cu2+ to Cu.

Cu2+(aq) + 2e- Cu(s). 2 mol of electrons 1 mol of Cu. How

much material is obtained?Q = I t

current (I) time (t) of the plating process.

Gibbs energy – the maximum amount of useful work that can be obtained from a system. nFEw

nFEG

wG

max

max

Gibbs Energy and Work

Note – if wmax is negative, then work is performed by the system and E is positive.

Electrical Work Eelectrolytic cell – external source

of energy is required to force the reaction to proceed.

External emf must be greater than Ecell.

From physics: work has units watts.1 W = 1 J/s.

Units of Electrical Work Electric utilities use units of

kilowatt-hours:

J. 106.3

W1

J/s 1

h 1

s 3600h 1 W1000kWh 1

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