Fig. 22-1a (p.629) A galvanic electrochemical cell at open circuit

1.1 components• Conducting electrodes (metal, carbon)• External wires (electrons carry current)• Ion electrolyte solution (ions carry current) • Interfaces or junctions• Complete electrical circuit

conduction:

electrons moving from one electrode to another thru. External wire

within solutions, migration of ions carry current at electrode surface, oxidation or reduction reactions occur

coupling the electron conduction of electrodes w/ ion conduction of

solutions

1.2 Galvanic Cells – cell develops spontaneous potential differenceOverall: 2Ag+(aq) + Cu(s) 2Ag(s) + Cu2+(aq)

Half reactions: Cu(s) Cu2+ (aq) + 2e- OxidationAg+ + e- Ag (s) Reduction

Convention: Cathode where reduction occursAnode where oxidation occurs

Galvanic cell: Cu anode (negative)Ag cathode (positive)

Cu2+ = 0.02 M, Ag+ = 0.02 M,E = 0.412 V 0 Vreaction equilibriumPotential difference (voltage) is measure of tendency to move to equilibrium

Fig. 22-1b (p.629) A galvanic cell doing work

1.3 electrolytic cells – require potential difference greater than galvanic potential difference ( to drive away from equilibrium)

Overall: 2Ag(s) + Cu2+(aq) 2Ag+(aq) + Cu(s)

[chemically reversible cell]

Half reactions: Ag(s) Ag+ (aq) + 2e- Oxidation

Cu2+(aq) + 2e- Cu(s) Reduction

electrolytic cell: Ag anode (negative)

Cu cathode (positive)

Fig. 22-1c (p.629) An electrolytic cell

1.4Cell w/o liquid junctions

AgCl(s) Ag+ (aq) + Cl- (aq)

H2 (g) H2 (aq)

Cathode: AgCl(s) Ag+ + Cl-

Ag+ + e- Ag (s)-

Anode: H2 (aq) 2H+ (aq) + 2e-

Overall: 2AgCl (s) + H2(g) 2Ag(s) + 2H+ + 2Cl-

direct reaction of AgCl + H2 is very slow

Fig. 22-2 (p.631) A galvanic cell without a liquid junction

1.5 Schematic representation of cells

Short-hand cell notation

Cu|CuSO4(Cu2+ = 0.0200) ||AgNO3 (Ag+ = 0.0200)| Ag

| liquid-electrode interface

|| two phase boundaries, one at each end of the salt bridge

convention: anode on left

Galvanic cell as written

Electrolytic cell if reversed

Pt, H2 (p=1atm)|H+ (0.01M), Cl-(0.01M), AgCl (sat’d) | Ag

2.1 Cell potential (difference between anode and cathode potential)

Ecell = Ecathode – Eanode

when half-reactions written as reduction (electron on left)

Example:

2AgCl(s) + H2(g) 2Ag (s) + 2H+ + 2Cl-

Cathode: 2AgCl(s) + 2e- 2Ag (s) + 2Cl-

Anode: 2H+ + 2e- H2(g)

Galvanic cell Ecell = Ecathode – Eanode = EAgCl/Ag – EH+/H2 = +0.46 V

Can’t measure potential on each electrode independently – only differences

2.2 Standard reference electrode

Standard hydrogen electrode (SHE)

Pt, H2(p=1.00atm) | H+ (H+ = 1.00M)||…

- SHE assigned 0.000V- can be cathode or anode {depending on the half cell which it couples with}

- Pt does not take part in reaction, coated with a finely divided layer of platinum to provide large surface area

- controlled activity of reactants- rarely used for routine measurement

Fig. 22-5 (p.637) measurement of the electrode potential for M electrode

Alternative references electrodes {which are simple to prepare}

a. Calomel electrode

Hg|Hg2Cl2(sat’d), KCl(xM)||

Hg2Cl2 (s) + 2e- 2Cl- + 2Hg (l)

Ereference depends on Cl-1

Ereference = +0.24V vs. SHE for saturated calomel electrode (SCE)

b. Silver-silver chloride electrode

Ag|AgCl(sat’d), KCl(xM)||

AgCl(s) + e- Cl- + Ag (s)

- Similar construction to calomel electrode

Ag wire coated with AgCl

Solution of KCl sat’d with AgCl

- Again, Ereference depends on Cl-, and = + 0.22 V vs. SHE

- Can be used for uncontrolled temperature (lower temperature coefficient, see Table 23-1)

Can be used for temp > 60C- But Ag reacts with more ions (e.g, proteins), while Hg reacts with few

sample components,

Fig. 23-2 (p.661) Typical commerical reference electrode a) A saturated calomel electrode, and b) a silver-silver chloride electrode

2.3 electrode and standard electrode potential (E and E0)- Definition: potential of electrodes vs. SHE

- Electrode potential varies with activity of ion

Activity x = x[X]

x: activity coefficient, varies with presence of other ions (ionic strength)

(see Appendix 2)

Note: activity of pure liquid or solid in excess = 1.00

Use pressure (atm) for gases

If = 1.00M, the electrode potential E, becomes standard electrode potential E0

Appendix 3

Cu2+ + 2e- Cu(s) E0 = +0.337 V

2H2+ + 2e- H2(g) E0 = +0.000 V

Cd2+ + 2e- Cd(s) E0 = - 0.403 V

Zn2+ + 2e- Zn(s) E0 = -0.763 V

2.4 Nernst equation In general, E and Ecell can be calculated for any activity using Nernst equation: pP + qQ + ne- rR + sS

E = E0 when log quotient is unityE0 is relative to SHEE0 is measure of driving force for half-cell reduction

qQ

pP

sSR

qQ

pP

sSR

qQ

pP

sSR

aa

ara

nEE

aa

araIn

nEE

F

RT

aa

araIn

nF

RTEE

)()(

)()(log

0592.0

)()(

)()(10568.2

10568.2

)()(

)()(

0

20

2

0

3. Electrical double layer*

electrons transferred at electrode surface by redox reactions, occurring at solution/electrode interface

Electrical double layer formed

1. tightly bound inner layer

2. loosely bound outer layer

Fig. 22-3 (p.632) Electrical double layer formed at electrode surface as a result of an applied potentials