chemistry 232
DESCRIPTION
Chemistry 232. Electrochemistry. A Schematic Galvanic Cell. Galvanic cells – an electrochemical cell that drives electrons through an external circuit spontaneous redox reaction occurring inside cell. . e -. Porous Disk. e -. e -. Oxidizing Agent. Reducing Agent. Anode. Cathode. - PowerPoint PPT PresentationTRANSCRIPT
Chemistry 232
Electrochemistry
A Schematic Galvanic Cell
Galvanic cells – • an electrochemical
cell that drives electrons through an external circuit
• spontaneous redox reaction occurring inside cell.
Anode Cathode
e-
Reducing Agent
e-
e-
Oxidizing Agent
Porous Disk
The Zinc/Copper galvanic cell.
e-
1.10 V
a(Zn2+) = 1.00e-
e-
Anode Cathode
Porous Disk or Salt BridgeZn(s) Cu(s)
e-
a(Cu2+) = 1.00
The Zn/Cu Galvanic Cell
Cu2+ (aq) + 2 e- Cu (s) (cathode, RHS)
Zn2+ (aq) + 2 e- Zn (s) (anode, LHS)
Cell Reactions The difference in the RHS and the LHS reaction
Cu2+ (aq) + Zn (s) Cu (s) + Zn2+ (aq) For each half reaction, we can write the reaction
quotient as followsCu2+ (aq) + 2 e- Cu (s) Q = 1/ a(Cu2+) Zn2+ (aq) + 2 e- Zn (s) Q = 1/ a(Zn2+)
Overall Qcell = a(Zn2+) / a(Cu2+)
Cell Diagrams A shorthand way of expressing what
takes place in an electrochemical cell. For the above electrochemical cell.
Pt Cu (s) Cu2+ (aq) Zn2+ (aq) Zn (s) Pt
Note phase boundary liquid junction
salt bridges
Another Example The cell reaction
H2 (g) + Cu2+ (aq) 2 H+ (aq) + Cu (s)
Pt H2 (g) H+ (aq) Cu2+ (aq) Cu (s) Pt
Electrochemical cells a cell that has not reached equilibrium can
do electrical work by driving electrons through an external wire.
Reversible Electrochemical Cells In order for us to make measurements on
an electrochemical cell, it must be operating reversibly. • Place an opposing source of potential in the
external circuit • Cell operates reversibly and at a constant
composition.
we,max = G
The Measurement of Cell Potentials Measure the potential of an electrochemical cell
when the cell is at equilibrium, i.e., the state between the galvanic and the electrolytic cell.
e-
Reducing Agente-
e-
Oxidizing Agent
Anode Cathode
Porous Disk
Counter potential (load)
Derivation of the Nernst Equation Consider an electrochemical cell that
approaches the equilibrium state by an infinitesimal amount d
d GddG rxnJJ
J Reminder
PTJ
JJr d
dG G,
The Work in Transporting Charge The maximum work
d Gdw rxne max,
F = Faraday’s constant = e NA = 96485 C/mole
For the passage d electrons from the anode (LHS) to the cathode (RHS)
d Fd eN A
The Cell Potential The work to transport charge
celle E d F dw max,
Gd
dwrxn
e
max,
cellrxn E- F G
Standard Cell Potentials From the reaction Gibbs energy
cellcello
rr E F QRTGG ln
cellcell
orr E
F QRT
F G
F G
ln
We define
F GE
or
cell
The Nernst Equation
E represents the standard cell potential, the potential of the cell when all cell components are under standard conditions. • f (all gases) = 1 • a (solutes) = 1• T = 298.15 K• P = 1.00 bar pressure
cellcell QF
RTEE ln
Cells at Equilibrium When the electrochemical cell has
reached equilibrium
cellcellcell KQV 0E
Kcell = the equilibrium constant for the cell reaction.
RTFE KK
F RTE cellcell
lnln
Knowing the E° value for the cell, we can estimate the equilibrium constant for the cell reaction.
Equilibrium Constant Calculations from Cell Potentials Examine the following cell.
Pt Sn2+ (aq), Sn4+ (aq) Fe3+ (aq) Fe2+ (aq) Pt
Half-cell reactions.Sn4+ (aq) + 2 e- Sn2+ (aq) E(Sn4+/Sn2+) = 0.15 VFe3+ (aq) + e- Fe2+ (aq) E (Fe3+/Fe2+) = 0.771 V
Cell ReactionSn2+ (aq) + 2 Fe3+ (aq) Sn4+ (aq) + 2 Fe2+ (aq)
Ecell = (0.771 - 0.15 V) = 0.62 V
Standard Reduction Potentials Standard reduction potentials are
intensive properties. We cannot measure the potential of an
individual half-cell! We assign a particular cell as being our
reference cell Assign values to other electrodes on that
basis.
a (H+) = 1.00
H2 (g)
e-
Pt gauze
The Standard Hydrogen Electrode Eo (H+/H2) half-cell = 0.000 V
f{H2(g)} = 1.00
A Galvanic Cell With Zinc and the Standard Hydrogen Electrode.
e-
Zn2+, SO42-
a (H+) = 1.00
Anode Cathode
Porous Disk or Salt Bridge
Source of H+ (e.g.,HCl (aq), H2SO4 (aq))
a(Zn2+) = 1.00
H2 (g)
0.763 Ve-
Zn(s)
Pt gauze
The Cell Equation for the Zinc-Standard Hydrogen Electrode. The cell reaction 2 H+ (aq) + Zn (s) H2 (g) + Zn2+ (aq)
Pt Zn (s) Zn2+ (aq),a=1 H+ (aq), a=1 H2 (g), f=1 Pt
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 the standard cell potential is positive, the rG is negative!
The Composition Dependence of the Cell Potential Nonstandard cell potential (Ecell) will be a
function of the activities of the species in the cell reaction.
cellcell QF
RTEE ln
To calculate Ecell, we must know the cell reaction and the value of Qcell.
Example For the following system
Pt H2 (g) H+ (aq) Cu2+ (aq) Cu (s) Pt
Calculate the value of the cell potential when the f (H2) = 0.50, a(Cu2+) = 0.20, and a(H+) = 0.40.
Concentration Cells Electrolyte concentration cell
• the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells.
Concentration Cells (II) Electrode concentration cells
• the electrodes themselves have different compositions. This may be due to.• Different fugacities of gases involved in
electrode reactions (e.g., The H+ (aq)/H2 (g) electrode).
• Different compositions of metal amalgams in electrode materials.
Applications of Electrochemistry
Measurement of activities and activity coefficients.
Electrochemical series.
Equilibrium constants and thermodynamic functions of cell reactions
Obtaining Standard Cell Potentials Look at the following cell
Pt H2 (g) HCl (aq) AgCl (s) Ag (s) Pt
)(ln2
cellcell HfCla Ha
FRTEE
Ecell = E(AgCl/Ag) - E (H+/H2) = E(AgCl/Ag)
Ecell Values and Activity Coefficients In dilute solution, using the DHLL
21
cell mF
RT3442AgAgClEmF
RT6064E .)(log.
Plot LHS vs. m1/2
Once Ecell is known, we can obtain experimental estimates of the mean activity coefficients.
The Calculation of Standard Cell Potentials
y = 0.0582x + 0.2222R2 = 0.9836
0.220
0.222
0.224
0.226
0.228
0.230
0.232
0.234
0.236
0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40
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)
Thermodynamic Information Note
GFEGr
PT
,
And GFE r
Entropy Changes To obtain the entropy change for the cell
reaction
PPTPTrxn T
EFGT
SS
,,
P
PTrxnrxn TEFG
TS
,
Enthalpy Changes To obtain the enthalpy change for the cell reaction
PTPTPTrxn
STGHH,,,
PTEFTFE
Prxn T
EFTFEH
The Liquid Junction Potential Examine the following electrochemical
cell Activity difference of the HCl between
compartment 1 and compartment 2 There should be a transport of matter
from one cell compartment to the other!
A Concentration Cell
e-
a (Cl -) = 0.0010
Left Right
Porous Disk or Salt Bridge
a(Cl-) = 0.010
0.0592Ve-
Ag(s) Ag(s)
The Development of Liquid Junction Potentials The cell
compartments are identical except for the activities of the electrolyte solutions.
HCl (a1)
HCl (a2)
Ag/AgCl electrode
Note that we now have the migration of both cations and anions through the liquid junction.
Cl-
Ag/AgCl electrode
H+
After a period of time
------------ - - - -
Ag/AgCl electrode
+ + + + +
Choose the lower compartment as our LHS electrode.
Ag AgCl Cl- (aq) a1 Cl- (aq), a2 AgCl (s) Ag (s)
Note: liquid junction
For the passage of one mole of charge through the cell
-F Ecell = GJ
The Cell Reactions For the LHS and RHS electrodes
AgCl (s) + e- Ag (s) + Cl- (a1) LHSAgCl (s) + e- Ag (s) + Cl- (a2) RHS
Net changeCl- (a1) Cl- (a2)
Note that the charge at the interface is transported by the anions and cations in the cell reaction!
The Transport Numbers How is the charge carried at the interface of the
cells?• t+ moles of charge carried by the H+ (cation).
• t- moles of charge carried by the Cl- (anion). Passage of one mole of “+” charge through the
interface • requires the passage of t+ moles of H+ (aq) from the LHS
RHS, and the passage of t- mole of Cl- charge from the RHS LHS.
At the boundary t+ H+(a1) + t- Cl-(a2) t+ H+(a2) + t- Cl-(a1)
For the entire cellCl- (a1) t+ H+(a1) + t- Cl-(a2) Cl- (a2) t+ H+(a2) +
t- Cl-(a1) The cell reaction involves the transport of t+
moles of HCl from the LHS to the RHs of the cell.
The Gibbs Energy Changes For the above cell reaction, we can write
the Gibbs energy expressions as follows
11
22
ClaRTClHaRTHtClaRTClHaRTHtG
)(ln)()(ln)()(ln)()(ln)(
1
2
ClaHaClaHaRTtG )()(
)()(ln
Cells With Transference Note a(H+) a (Cl-) = {a (HCl)}2
1
2
HClaHClaRTt2G )(
)(ln
1
2wt HCla
HClaF
RTt2E )()(ln
Note that the cell potential with transference, Ewt is determined as follows
Cells without Transference What if we were able to set up a cell so that the
transport at the interface did not contribute to the overall G?
The potential of this cell would be the cell potential without transference, Ewot.
Cl- (a1) Cl- (a2)
1
2
HClaHClaRTG )(
)(ln
1
2wot HCla
HClaF
RTE )()(ln
The Liquid Junction Potential The liquid junction potential is the
difference in the cell potentials with and without transference!
wotwtLJ EEE
1
2LJ HCla
HClaF
RTt21E )()(ln
L.J. Potentials Depend on Transport Numbers What is the following were true? t+ t- 0.5 ELJ would be very small
and would only make a small contribution to the overall cell potential !
L.J. Potentials Depend on Transport Numbers ELJ a potential problem any time we
measure the cell potential whose electrodes have different electrolytes
How does the salt bridge help?• e.g., for species with t+ t- 0.5, the ELJ
values are small and are readily established!