1. 1
VII. ELECTROCHEMISTRY (Chapters 13-16)
FUNDAMENTALS (Chapter 13)
Basic Concepts:
Electric charge q is measured in Coulombs (C)
Charge of 1 electron = 1.602 x 10-19 C
Charge of 1 mole electrons = (6.022 x 1023 mol-1) x (1.602 x 10-19 C )
= 9.649 x 104 C = Faraday constant F
Charge of n mole electrons = q = nF
Oxidizing agent: accepts electrons from another substance which is oxidized,
while the oxidizing agent is reduced
(examples: O2, F2, acidified MnO4
-, acidified Cr2O7
2-, Ce4+)
Reducing agent: donates electrons to another substance which is reduced,
while the reducing agent is oxidized
(examples: alkali metals, SO3
2−, Fe2+)
2
ΔG = - E • q = - nFE
Relates the free-energy change of a reaction to
the voltage that can be generated by that reaction
Ohm’s law: E = R • I (where R is the resistance, in ohm Ω)
Power P is work done per unit time, in units of Js-1 or watt (W)
P = work / s = E • q / s = E • I
The maximal possible electrical work that can be done by a chemical reaction
on its surroundings (at ct. T and p) is the free-energy change (ΔG):
-ΔG = work done on surroundings
The difference in electric potential E between two points is the work needed
(or that can be done) to move an electric charge from one point to the other,
per unit charge. It is measured in Volts (V).
Work = E • q (in Joule, J)
The quantity of charge per second is called the current I (in ampere A = Cs-1)
C3100 - Winter 2011
Dr. G. Van Biesen
2. 3
Galvanic Cells
A galvanic cell (e.g. battery, fuel cell) uses a spontaneous chemical reaction to
generate electricity – one reagent is oxidized, the other is reduced. The
reagents are physically separated so that electrons have to flow through an
external circuit to go from one reactant to the other.
A potentiometer in the external circuit measures
the voltage generated by the cell.
The two half-cells are connected with a
salt bridge which contains a gel with a
high conc. of electrolyte that does not
affect the cell reaction – it allows ions to
diffuse, thereby maintaining
electroneutrality.
Oxidation occurs in one half-cell (anode),
while reduction takes place in the other
(cathode).
Cd(s) + 2Ag+ Cd2+ + 2Ag(s)
Ecell = E+ - E-
Cd(s) Cd(NO3)2 AgNO3 Ag(s)
Phase boundary Salt bridge
Change in ΔG
4
Standard reduction potentials (E°) for half cells are determined experimentally.
(‘Standard’ means at 25 °C and all activities 1)
LHS cell: Standard Hydrogen Electrode
(S.H.E.): Pt surface in contact with an
acidic solution (A(H+) = 1) through which
H2 gas is bubbled (A(H2) = 1)
H+ + e- ½ H2 (g)
By convention: E° (S.H.E.) = 0 V
RHS cell: Ag+ + e- → Ag
By convention: the left-hand electrode is
connected to the negative (reference)
terminal of the potentiometer, the right-
hand electrode to the positive terminal.
Idealized exp. to determine E° for: Ag+ + e- → Ag:
e- (lost by the species undergoing oxidation) flow
from the anode to the cathode, where they are
gained by the species undergoing reduction.
The potentiometer measures the difference in reduction potentials of both half
cells, and indicates a positive voltage when electrons flow into the negative
terminal.
V or Ecell = E+ - E-
C3100 - Winter 2011
Dr. G. Van Biesen
3. 5
Redox potentials always written as
reduction potentials. The more positive
the reduction potential, the stronger the
driving force for reduction to occur:
ΔG = - nFE
K+ is a very poor oxidizing agent –
does this imply it is a strong reducing
agent? What about K?
With the cell: S.H.E. Cd2+ Cd(s) we find Ecell = E°Cd2+ = -0.402 V and e-
flow in the other direction – the LH half cell is now the cathode; the RH cell is
the anode.
In a galvanic cell, reduction spontaneously occurs in the half-cell with the
more positive reduction potential (at the cathode). Oxidation spontaneously
occurs in the other half-cell (at the anode).
Here, Ecell = E°Ag+ = +0.799 V, which means that e- flow into the neg.
terminal; the LH cell is the anode; the RH cell is the cathode.
6
The Nernst Equation:
The sign and magnitude of Ecell can only be predicted from tabulated values if
both half cells operate under standard conditions; this will rarely be the case,
and half cell potentials under non-standard conditions have to be calculated
using the Nernst equation.
aA + ne- bB
E°= standard reduction potential (AA = AB = 1)
R = gas ct. = 8.314 J / (K•mol)
T = temperature in K
n = # e- in half reaction
F = Faraday ct. = 96485 C/mol
Reaction quotient Q = ; when all activities are unity, E = E°
(Note: pure solids, pure liquids, and solvents have activity = 1)
At 25 °C, we can simplify this equation to:
AB
b
AA
a
We also replace activities of
solutes by their concentrations;
for gases we use pressure (bar)
C3100 - Winter 2011
Dr. G. Van Biesen
4. 7
EXAMPLE: for the reduction of white phosphor to phosphine gas:
¼ P(s) + 3 H+ + 3 e- PH3(g) E° = -0.046 V
E = -0.046 - log
0.05916
3
PPH3
[H+]3
Note that if we multiply the half-reaction by any factor, this does not
change E…
P(s) + 12 H+ + 12 e- 4 PH3(g)
E = -0.046 - log
0.05916
12
P4
PH3
[H+]12
= -0.046 - 4 log
0.05916
12 3
PPH3
[H+]3
PPH3
[H+]3
4
= -0.046 - log
0.05916
12
8
…nor does it change E°: remember that work = E • q, or: E = work / q
Potential difference is work done per unit of charge carried through that
potential difference (i.e. regardless of the amount of charge; q = nF can be
any number).
Stepwise procedure for finding the potential of a galvanic cell:
Identify in each half-cell the element that occurs in two oxidation states
Write reduction half-reactions for each half-cell – find E° for each in a table
(e.g. Appendix H in Harris 7th Ed.). Multiply half-reactions so that they
contain the same # of e- (do not multiply E°!).
Write Nernst equations for E+ (RH cell; pos. terminal) and E- (LH cell; neg.
terminal).
Find: Ecell = E+ - E- If Ecell > 0, then net cell reaction is spontaneous, if
Ecell < 0 then net cell reaction is spontaneous in the reverse direction.
For a balanced net cell reaction, subtract left half-reaction from right half-
reaction (i.e. reverse left half-reaction and add it to the right half-reaction).
C3100 - Winter 2011
Dr. G. Van Biesen
5. 9
A more intuitive way of looking at cell potentials:
Electrons always flow towards
the more pos. potential
Half-cell with higher E ‘wants to
be reduced’, the one with the
lower E ‘wants to be oxidized’
Half-cell reactions can sometimes be written in
different ways, but the half-cell potential should be
the same:
e.g.: AgCl(s) + e- Ag(s) + Cl-
E°+ = 0.222 V
E+ = 0.222 – 0.05916 log[Cl-]
= 0.222 – 0.05916 log(0.0334)
= 0.309 V
10
We could also have written the RH half-cell reaction as:
Ag+ + e- Ag(s) E°+ = 0.799 V
(In both reactions, Ag+ is reduced to Ag)
E+ = 0.799 – 0.05916 log
1
[Ag+]
[Ag+] =
Ksp
[Cl-]
=
1.8 x 10-10
0.0334
= 5.4 x 10-9 M
= 0.310 V
E° and the Equilibrium Constant:
The potentiometer that measures the
voltage of the cell has a very high
resistance (1013 Ω), and thus allows
only a very small current (e.g. if Ecell =
1 V, then I = 10-13 A). This does not
noticeably change the concentrations
of the reactants in the cell.
A galvanic cell produces electricity because
the cell reaction is not at equilibrium.
Replacing the potentiometer with a wire,
there would be an appreciable current flowing
through it, and the concentrations of the
reactants would change until equilibrium is
reached (Ecell = 0)
C3100 - Winter 2011
Dr. G. Van Biesen
6. 11
aA + ne- cC E°+
dD + ne- bB E°-
True at any time
Special case: cell is at equilibrium: Ecell = 0; Q = K
12
EXAMPLE: Find K for the reaction:
Cu(s) + 2Fe3+ 2Fe2+ + Cu2+
The two half-reactions are:
2Fe3+ + 2e- 2Fe2+ E°+ = 0.771 V
- (Cu2+ + 2e- 2Cu(s)) E°- = 0.339 V
E° = E°+ - E°- = 0.432 V K = 10(2)(0.432)/(0.05916) = 4 x 1014
Note: the procedure can also be used to find K for reactions that are not redox
reactions! (e.g. FeCO3(s) Fe2+ + CO3
2- K = Ksp).
It is just a matter of finding half-reactions with known E° values that, when
combined, add up to that particular reaction. A redox potential is simply another
way of expressing the free-energy change of a reaction. See Box 13-3 on
Latimer diagrams in Harris 8th Ed.
C3100 - Winter 2011
Dr. G. Van Biesen
7. 13
Biochemists use E°':
Standard reduction potentials are defined with all participating species at
activity = 1. Often, H+ is involved in reactions, and this means that all E°
values are for pH = 0. This is not very appropriate for biological systems,
where pH ≈ 7.
E°’ is the formal potential at pH 7
(‘formal’ means that it applies under a specific set of conditions e.g. at certain
ionic strength, conc. of complexing agents, or at a certain pH)
aA + ne- bB + mH+ E°
E = E° - log
0.05916
n
[B]b[H+]m
[A]a
Conversion from [A] to FA : see
fractional composition equations.
IV. Acid/Base Chemistry, slides
35-39
A and B could be acids/bases
14
ELECTRODES AND POTENTIOMETRY (Chapter 14)
Potentiometry = use of electrodes to measure voltages that provide chemical
information
A solution containing an electroactive analyte (analyte that can donate or accept
electrons at an electrode) is turned into a half-cell by inserting an indicator
electrode (e.g. Pt wire), and is connected via a salt bridge to another half-cell
with a fixed composition (fixed potential), the reference electrode. The cell
voltage is then measured, and depends on the concentration of the analyte.
Reference Electrodes:
Consider a galvanic cell with Fe2+/Fe3+ and a Pt wire as one half-cell, and a
constant potential half-cell:
RH cell: Fe3+ + e- Fe2+ E°+ = 0.771 V
LH cell: AgCl(s) + e- Ag(s) + Cl- E°- = 0.222 V
E = 0.771 – 0.05916 log - 0.222 - 0.05916 log[Cl-]
[Fe2+]
[Fe3+]
= ct
Cell voltage only depends
on [Fe2+] / [Fe3+] !
C3100 - Winter 2011
Dr. G. Van Biesen
8. 15
The type of reference electrode in the RH figure is a silver-silver chloride
reference electrode:
Its standard reduction potential is 0.222 V, but this is with A(Cl-) = 1.
The activity of Cl- in a saturated solution ≠ 1, and Esat. KCl = 0.197 V
AgCl(s) + e- Ag(s) + Cl- E°- = 0.222 V
Reference electrode
connected to reference
terminal (-) of potentiomenter
16
Problem with these type of electrodes: can
get clogged, not compatible with some types
of solutions (e.g. solutions containing Ag and
Ag-complexing compounds such as Tris).
Double-junction may be necessary
Another commonly used reference electrode is
the saturated calomel electrode (S.C.E.):
½ Hg2Cl2(s) + e- Hg(l) + Cl- E° = 0.268 V
Esat. KCl= 0.241 V
C3100 - Winter 2011
Dr. G. Van Biesen
9. 17
The S.H.E. (slide 4) is not commonly used as a reference electrode because
it is not practical to work with.
The following diagram allows conversion from voltages obtained using one
reference electrode, to another:
Point A: –0.220 V from the S.H.E.
-0.461 V from the S.C.E.
-0.417 V from the Ag/AgCl electrode
Point B: +0.230 V from the S.H.E.
-0.011 V from the S.C.E.
+0.033 V from the Ag/AgCl electrode
18
Limitations to the Use of Electrode Potentials:
Voltages measured in lab are not always the same as calculated ones, mainly
because of two reasons:
M(s)
Mn+(aq)
Note: four types of electrodes used in electrochemistry:
M(s), MXn(s)
X-(aq) Mn+(aq),
Mm+(aq)
Pt(s)
Pt(s)
Metal / metal ion
e.g. Cd/Cd2+
Metal / 'insoluble' salt
e.g. Ag/AgCl/Cl-
Redox electrode
e.g. Pt/Sn2+/Sn4+
Gas electrode
e.g. S.H.E
C3100 - Winter 2011
Dr. G. Van Biesen
10. 19
Liquid junction potentials:
Most galvanic cells have a salt bridge → electrolyte solutions of different
composition in contact with each other (with a porous glass disk: allows ions to
go through, but minimizes mixing) = liquid junction
Liquid junctions always develop a potential at their interface, usually a few mV.
This potentials contributes to the cell voltage (because the voltage of a galvanic
cell is due to the potential difference developed at each of the electrode-solution
and liquid-liquid interfaces), but is generally not known and puts a fundamental
limitation on the accuracy of potentiometric measurements
Development of a liquid junction potential:
Na+ and Cl- diffuse into the water (from high to low conc.)
Cl- has a higher mobility than Na+ → a region rich in Cl-
(with excess neg. charge) develops at the front;
immediately behind it is a pos. charged region,
with more Na+ than Cl-. This creates a potential difference at the junction of the NaCl and H2O phases.
The steady-state potential opposes the movement of Cl-,and accelerates the movement of Na+.
Since K+ and Cl- have similar mobilities, KCl is often used in salt bridges
because this minimizes the junction potential. Where Cl- cannot be used (e.g.
with half-cells containing Ag+), NO3
- is a good substitute.
20
Indicator Electrodes:
Metal electrodes: electric potential develops in response to a redox reaction at
the metal surface
Ion-selective electrodes: selective binding of analyte generates electric potential
no oxidation-reduction takes place!
Reversibility of the electrode reaction:
For the potential of an electrode to be properly described by the Nernst
equation, the electrode reaction must be reversible, and rapid, but that is not
always the case!
e.g.: 2 CO2(g) + 2 H+ + 2 e- H2C2O4
The rate at which CO2 reacts to produce oxalic acid is extremely
slow. Changes in partial pressure of CO2 therefore have little effect
on E; the Nernst equation does not apply.
E = E° - log
0.05916
2
[H2C2O4]
[H+]2
(pCO2)2
C3100 - Winter 2011
Dr. G. Van Biesen
11. 21
Metal electrodes can be made from the same metal as the cation to be
analyzed (Ag/Ag+, Cu/Cu2+,…) or from an inert metal (Pt, Au), or from C.
Ag+ + e- Ag(s) E°+ = 0.799 V
½ Hg2Cl2(s) + e- Hg(l) + Cl- Esat. KCl= 0.241 V
E = 0.799 – 0.05916 log - 0.241 V = (0.558 + 0.05916 log [Ag+]) V
1
[Ag+]
Ideally, E changes by 59.16 mV for every
factor of 10 change in [Ag+]
This type of cell can also be used to
measure halide concentrations (e.g. Cl-)
if AgX is present at the surface of the
Ag electrode, since [Ag+] [X-] = Ksp
22
Ion-selective electrodes respond selectively to one ion (ideally) – they do not
involve redox processes. There are various types, but most have a thin
membrane which binds only a specific analyte ion (solid-state ion-selective
electrodes are based upon doped inorganic crystals). Analyte ions equilibrate
with ion-exchange sites at the outer surface of the membrane. This creates a
charge imbalance (=potential) across the phase boundary between the solution
and the membrane. This potential difference can be related to concentration
using calibration curves.
A glass electrode for pH measurements is the most common ion-selective
electrode. It is usually a combination electrode, which incorporates a Ag/AgCl
reference electrode.
Other ion-selective electrodes: K+, Ca2+, Mg2+, F-, Cl-, S2-
Extensively used in clinical diagnosis
Respond to activity of uncomplexed analyte ion only!
Often used with standard addition
C3100 - Winter 2011
Dr. G. Van Biesen
12. 23
pH Measurement with a Glass Electrode
Glass electrode = ion-selective electrode
Sensitive to H+, to some extent also to Na+ (especially at pH > 12)
Usually as a combination electrode = both glass and reference in one unit
Ag(s) AgCl(s) Cl-
(aq) H+
(aq, outside) glass membrane H+
(aq, inside), Cl-
(aq) AgCl(s) Ag(s)
The two reference electrodes measure the electric potential difference across
the glass membrane. Except for Eb, which depends on [H+] on either side of
the membrane, all other potentials are constant. Since [H+]intside is fixed, Eb in
fact only depends on [H+]outside (i.e. pH of analyte solution).
Ext. ref. electrode Int. ref. electrode
Glass electrodeEref1 Ej Eb Eref2Easy
Ecell = Eref2 + Eb + Easy – Eref1 + Ej
24
C3100 - Winter 2011
Dr. G. Van Biesen
13. 25
O
Si
cation
The glass membrane consists for ~70% SiO2,
with the remainder oxides of Ca, Ba, Li, and Na.
In the silicate lattice of the glass membrane, those
oxygen atoms that are not shared by two Si
atoms have a neg. charge and are coordinated to
cations. Singly charged cations (Na+, Li+) are
mobile in this lattice and are responsible for
electrical conduction within the membrane.
(H+ cannot cross the membrane)
The surfaces of the glass membrane in
contact with water become hydrated (they
swell), and act as cation-exchangers. Metal
cations in these regions become mobile,
diffuse out of the membrane, and can be
replaced by H+ from the solution– H+ is the
only ion that binds significantly to the
hydrated gel layer.
26
H+ + Na+Gl- Na+ + HGl
soln glass soln glass
K is large, so ordinarily surface of
hydrated glass membrane is entirely
HGl (except at high pH)
Each side of the membrane acquires a negative charge, the magnitude of
which depends on the [H+] of the solution with which it is in contact. The
potential that develops across the membrane is the result of an unequal
charge buildup at opposing surfaces of the membrane
The glass in the hydrated layer can be
considered as a weak acid:
HGl H+ + Gl-
glass soln glass
E = k + 0.05916 log Ao
(Ao is the activity of H+ in the outer solution)
The response of the glass electrode is given by:
In practice: E = k - β (0.05916) pH
Or: E = k - 0.05916 pH
C3100 - Winter 2011
Dr. G. Van Biesen
14. 27
Calibration of a glass electrode, and pH measurement
Use two (or more) buffer solutions:
First buffer (known pHS1) gives a voltage ES1
Second buffer (known pHS2) gives a voltage ES2
Equation of the line
through both points:
For a solution of unknown
pH, we measure Eunknown
and pHunknown is
automatically calculated
and displayed via the
previously determined
relationship
With the second buffer β is determined; should be
close to 1
28
Using glass electrodes:
•Electrode has to be stored (ideally) in same solution as that in the reference
compartment. If the membrane dehydrates, it has to be rehydrated for several
hours before use.
•Remove cap at upper end of electrode which covers the air inlet (prevents
evaporation of filling solution when not in use) – see slide 24
•Calibrate the electrode with appropriate buffers – equilibrate with stirring for at
least a minute – let reading stabilize.
•Rinse the electrode with deionized water when transferring to another solution –
the membrane can be gently blotted, but do not wipe it, because this may
produce a static charge on the glass.
•Most pH meters have a dial to compensate for temperature – see slide 6.
•Accuracy of pH measurements with glass electrodes are ± 0.02 pH units (this
corresponds to a uncertainty of ± 5% in AH+), although measurements of pH
differences between solutions can be 10x more accurate. Main sources of
errors are standards and liquid junction potentials.
C3100 - Winter 2011
Dr. G. Van Biesen
15. 29
REDOX TITRATIONS (Chapter 15)
Ce4+ + Fe2+ Ce3+ + Fe3+
Reactions at the Pt electrode:
Fe3+ + e- Fe2+ E° = 0.767 V
Ce4+ + e- Ce3+ E° = 1.70 V
E° for titration reaction = 1.70 – 0.767
= 0.933 V; K = 10(0.933/0.05916) ≈ 1016
(i.e. quantitative)
The Pt indicator electrode responds to
relative conc. (activities) of Ce4+ and
Ce3+, or Fe2+ and Fe3+
Before EP:
With the addition of each aliquot of Ce4+, Fe2+ is consumed and converted
into Fe3+. We can easily calculate [Fe2+], [Fe3+] and [Ce3+]; [Ce4+] is more
difficult to calculate. Therefore, we use:
30
E = E+ - E- = 0.767 – 0.05916 log - 0.241
[Fe2+]
[Fe3+]
E = 0.526 – 0.05916 log
[Fe2+]
[Fe3+]
Note: 1. When V = ½ VEP, E+ = E°+ for the Fe3+/Fe2+ couple
2. The voltage at V = 0 cannot be calculated (E = -∞?)
ES.C.E.
At EP:
All Fe2+ and Ce4+ are converted to Fe3+ and Ce3+; which exist in equilibrium
with Fe2+ and Ce4+ as:
Fe3+ + Ce3+ Ce4+ + Fe2+ Since [Fe3+] = [Ce3+], [Fe2+] = [Ce4+]
Both half-reactions are in equilibrium at the Pt electrode at any time:
E+ = 0.767 – 0.05916 log
[Fe2+]
[Fe3+]
E+ = 1.70 – 0.05916 log
[Ce3+]
[Ce4+]
C3100 - Winter 2011
Dr. G. Van Biesen
16. 31
2 E+ = 0.767 – 0.05916 log
[Fe2+]
[Fe3+]
+ 1.70 – 0.05916 log
[Ce3+]
[Ce4+]
2 E+= 2.467 – log
[Fe2+]
[Fe3+]
[Ce3+]
[Ce4+]
2 E+ = 2.467 V ⇒ E+ = 1.23 V
The cell voltage E = E+ - Ecalomel = 1.23 – 0.241 = 0.99 V
[Fe3+] = [Ce3+] and [Fe2+] = [Ce4+] at EP
After EP:
The # moles of Ce3+ = the # moles of Fe3+, and there is a known excess of
(unreacted) Ce4+. We now use:
E = E+ - Ecalomel = 1.70 – 0.05916 log - 0.241
[Ce3+]
[Ce4+]
When V = 2 VEP, [Ce3+] = [Ce4+], and E+ = E°+ = 1.70 V
E = E+ - Ecalomel = 1.70 – 0.241 = 1.46 V
Independent of
conc., for this
particular titration
32
EXAMPLE: 100.0 mL 0.0500 M Fe2+ titrated with 0.100 M Ce4+
What are the cell potentials at 36.0, 50.0 (= VEP) and 63.0 mL?
•At 36.0 mL: A fraction of Fe of 36.0/50.0 is now in the form of Fe3+; a
fraction of 14.0/50.0 is still in the form of Fe2+
Ce4+ + Fe2+ Ce3+ + Fe3+
E = 0.526 – 0.05916 log
14/50
36/50
E = 0.550 V
•At 50.0 mL (=VEP), the cell voltage here is
independent of concentration, and V = 0.99 V
•At 63.0 mL: with the first 50.0 mL of Ce4+
converted into Ce3+, there is now a 13.0 mL
excess of Ce4+:
E = E+ - Ecalomel = 1.70 – 0.05916 log - 0.241
50
13
E = 1.424 V
The actual concentrations of Fe2+
and Fe3+ are 0.0103 and 0.0265
M, respectively
The actual concentrations of
Ce3+ and Ce4+ are 0.0307 and
0.00798 M, respectively
C3100 - Winter 2011
Dr. G. Van Biesen
17. 33
The titration curve is symmetrical around the EP only because the reaction
stoichiometry is 1:1. For a 2:1 ratio, the curve is not symmetrical. However,
with a steep rise in E near the EP, the error in VEP is minimal.
IO3
- + 2Tl+ + 2Cl- + 6H+
ICl2
- + 2Tl3+ + 3H2O
The more the E° values of the half-
reactions differ, the bigger the
voltage change at the EP – best
results are obtained with strong
oxidizing/reducing agents.
A difference in formal potentials of
> 0.200 V usually gives satisfactory
results.
Since half-cell potentials depend on the ratio of concentrations, they are
independent of the concentration! In practice, the lower limit is ~ 1 mM, since at
lower concentrations impurities will be present at much larger relative amounts,
and may consume a significant proportion of titrant.
34
EXAMPLE: potentiometric titration of Fe2+ (400.0 mL; 3.75 mM) with MnO4
-
(20.0 mM); calomel electrode as reference (Ecalomel = 0.241 V) in 1.00 M H2SO4
Fe3+ + e- Fe2+ E° = 0. 68 V (in 1 M H2SO4)
MnO4
- + 8H+ + 5e- Mn2+ + 4H2O E° = 1.507 V
MnO4
- + 8H+ + 5Fe2+ Mn2+ + + 5Fe3+ + 4H2O
VEP? # moles MnO4
- = # moles Fe2+
1 mol MnO4
-
5 mol Fe2+
VEP x 20 mM = 1/5 x 400 mL x 3.75 mM ⇒ VEP = 15.0 mL
•Before EP, e.g. 12.0 mL MnO4
- added:
E = E+ - E- = 0.68 – 0.05916 log - 0.241
[Fe2+]
[Fe3+]
E = 0.439 – 0.05916 log
3.0/15.0
12.0/15.0
E = 0.475 V
C3100 - Winter 2011
Dr. G. Van Biesen
18. 35
•At EP: we again add both half-reactions:
E+ = 0.68 – 0.05916 log
[Fe2+]
[Fe3+]
5 E+ = 1.507 – log0.05916
5
[Mn2+]
[MnO4
-] [H+]8
We multiply by 5 so we can sum
these equations more easily – this is
a purely algebraic operation and
now we have to multiply E+ as well!
6 E+ = 8.215 – 0.05916 log
[Mn2+]
[MnO4
-] [H+]8
[Fe2+]
[Fe3+]
From the net reaction, we
see that at the EP [Fe2+] = 5
[MnO4
-], and 5 [Mn2+] = [Fe3+]
6 E+ = 8.215 – 0.05916 log 1
[H+]8
With [H+] = 1.00 M = 0.964 M
400
415
E+ = 1.368 V; the cell voltage E = 1.368 – 0.241 = 1.127 V
•After EP: e.g. 17.0 mL MnO4
- added: excess of 17.0 – 15.0 = 2.0 mL MnO4
-
We now use the other half-reaction:
36
[Mn2+]? All the Mn2+comes from oxidation of Fe2+ by MnO4
-, and only as much
Mn2+ can be formed as the amount of MnO4
- added at VEP:
[Mn2+] = = 0.719 mM
MnO4
- + 8H+ + 5e- Mn2+ + 4H2O E° = 1.507 V
E+ = 1.507 – log
0.05916
5
[Mn2+]
[MnO4
-] [H+]8
15.0 mL x 20.0 mM
(400.0 mL + 17.0 mL) Or using [Fe2+]:
[Mn2+] =
400.0 mL x 3.75 mM x 1/5
400.0 mL + 17.0 mL
= 0.719 mM
[MnO4
-]? All MnO4
- that is added after the
EP is in excess:
[MnO4
-] = = 0.0959 mM
(17.0 mL – 15.0 mL) x 20.0 mM
(400.0 mL + 17.0 mL)
[H+]? The original [H+] is simply diluted by the addition of the titrant:
[H+] = 1.00 M = 0.959 M
400.0 mL
400.0 mL + 17.0 mL
The small amount of H+
consumed during the
reaction can be ignored.
C3100 - Winter 2011
Dr. G. Van Biesen
19. 37
E+ = 1.507 – log = 1.495 V
0.05916
5
0.719 x 10-3
(0.0959 x 10-3) x (0.959)8
The cell voltage E = 1.495 – 0.241 = 1.254 V
Finding the End Point of a Redox Titration:
Can use 2nd derivative method from a plot of V vs. volume, analogous to
acid-base titrations (Section V, slide 57), or can use a redox indicator.
Sometimes, the color of the titrant itself indicates the end point (e.g. the
previously described titration of Fe2+ with MnO4
-).
A redox indicator changes color going from the oxidized to the reduced state.
Prediction of the potential range over which an indicator changes:
In(ox) + ne- In(red)
E = E° – log0.05916
n
In(red)
In(ox)
38
As with acid-base indicators, the color
of In(red) will be observed when:
In(red)
In(ox)
≥
10
1
The color of In(ox) will be observed when:
In(red)
In(ox)
≤
1
10
The color change thus occurs over the range: E = E° ± V
0.05916
n
Ferroin (E° = 1.147 V) will change color in the range of 1.088 – 1.206 V with
respect to the S.H.E.; this is the range of 0.847 – 0.965 V vs. S.C.E. This
would be a good indicator for the titration on slide 29.
C3100 - Winter 2011
Dr. G. Van Biesen
20. 39
Adjustment of Analyte Oxidation State:
Sometimes an analyte has to be oxidized or reduced before it can be titrated.
This adjustment of the analyte oxidation state has to be quantitative, and the
reagent used has to be eliminated/removed completely before analysis.
Preoxidation:
Peroxydisulfate (or persulfate): S2O8
2- (requires Ag+ as a catalyst)
S2O8
2- + Ag+ SO4
2- + SO4
- + Ag2+
Excess reagent is destroyed by boiling after oxidation is complete:
2 S2O8
2- + 2 H2O 4 SO4
2- + O2 + 4 H+
Silver(II)oxide: AgO
Dissolves in mineral acids to give Ag2+; similar oxidizing power as persulfate
Excess AgO is removed by boiling: 4 Ag2+ + 2 H2O 4 Ag+ + O2 + 4 H+
Solid sodium bismuthate: NaBiO3
Similar oxidizing strength as above oxidants, excess reagent removed by
filtration.
2 powerful oxidants (E°(Ag2+) ≈ 2.0 V)
40
Hydrogen peroxide: H2O2 (in basic solution); medium oxidizing power
H2O2 + 2e- 2 OH- E° = 0.88 V
Can be used for oxidation of Co2+ → Co3+; Fe2+ → Fe3+; Mn2+ → MnO2
Excess removed by boiling: 2 H2O2 2 H2O + O2 (disproportionation)
In acidic solution, it can reduce Cr2O7
2- to Cr3+, and MnO4
- to Mn2+
Analytes that can be pre-oxidized by the reagents listed above:
Mn2+ to MnO4
-; Ce3+ to Ce4+; Cr3+ to Cr2O7
2-
The oxidized analytes can then be titrated with e.g. a standardized Fe2+ solution
Prereduction:
Stannous chloride (SnCl2): reduces Fe3+ to Fe2+ in hot HCl
Excess SnCl2 is destroyed by adding HgCl2:
SnCl2 + 2 HgCl2 Sn4+ + Hg2Cl2 + 4 Cl-
C3100 - Winter 2011
Dr. G. Van Biesen
21. Ox + Zn(Hg)(s) Red + Zn2+ + Hg(l)
(analyte)
Zn(Hg)(s)
(analyte)
41
Jones reductor: uses a column filled with Zn / Zn(Hg):
Zn(s) + HgCl2 Zn(Hg)(s) + ZnCl2
Zn(Hg) is preferred over Zn because it reduces H+ much more
slowly than Zn, and thus can be used for acidic solutions
Non-selective (E0 = -0.80 V); reduces many ions e.g. Fe3+, Cr3+, UO2
2+
Most reduced analytes are oxidized again by air, so the eluent has to be
titrated immediately.
Walden reductor: column is filled with Ag and 1 M HCl: more selective
(higher E°: 0.222 V), Cr3+ and TiO2
+ are not reduced and thus do not
interfere with the analysis of e.g. Fe3+
The eluent can also be collected in an acidic Fe3+ solution (except when Fe
is analyzed), which is reduced by the analyte(s) to Fe2+ (stable) and can then
be titrated with an oxidant such as MnO4
- (indirect determination)
H+
42
Oxidizing Agents for Redox Titrations:
Potassium Permanganate: KMnO4
MnO4
- + 8H+ + 5e- Mn2+ + 4H2O E° = 1.507 V acid
MnO4
- + 4H+ + 3e- MnO2 + 2H2O E° = 1.692 V neutral
MnO4
- + e- MnO4
2- E° = 0.56 V basic
In strongly acidic solutions, MnO4
- is its own indicator (Mn2+ is colorless)
KMnO4 is not a primary standard (contains traces of MnO2), and dissolved
KMnO4 reacts with organic impurities to produce some MnO2. Freshly
prepared KMnO4 solutions are boiled for an hour (speeds up the reaction)
and MnO2 is filtered (not with filter paper!). In addition, it oxidizes H2O:
4 MnO4
- + 2 H2O 4 MnO2(s) + 4 OH- + 3 O2
Frequent standardization is necessary – can use Na2C2O4 or pure iron wire:
C3100 - Winter 2011
Dr. G. Van Biesen
22. 43
[MnO4
− + 8 H+ + 5 e− Mn2+ + 4 H2O] x 2
[H2C2O4 2 CO2 + 2 H+ + 2 e−] x 5
2 MnO4
− + 5 H2C2O4 + 6 H+ 2 Mn2+ + 10 CO2 + 8 H2O
This reaction is slow at room temperature; when most of the titrant (90-95%)
is added, the reaction mixture is heated up to ~60 °C to drive off CO2 and
shift the equilibrium to the right. A blank is performed to account for the
amount of titrant necessary to impart a pink color to the solution.
A pure iron wire can be dissolved in warm 1.5 M H2SO4 (under N2); the cooled
solution can be titrated directly – 5 mL of H3PO4 / 100 mL solution can be
added to mask the yellow color of Fe3+ to make the endpoint more distinct.
If great accuracy is not needed, Fe(NH4)2(SO4)2.6H2O can also be used
(sufficiently pure for most purposes).
An indirect application of KMnO4 as a titrant is in the analysis of non-oxidizable
or difficult to oxidize cations such as Ca2+, Mg2+, Ce3+, Cu2+. They can be
precipitated with Na2C2O4, filtered, redissolved in acid and the H2C2O4 can then
be titrated as above.
44
Ammonium hexanitratocerate: (NH4)2Ce(NO3)6
Primary standard – usually dissolved in 1 M H2SO4 (slow decomposition in
other mineral acids such as HClO4and HNO3)
Ce(IV) binds anions such as ClO4
-, NO3
-, Cl- etc. strongly, as indicated by
variation of the E° in HClO4 (E° = 1.70 V), HCl (E° = 1.47 V) etc.
Need an indicator (Ce4+ is yellow, Ce3+ is colorless, but no sharp change),
e.g. ferroin.
Can be used instead of KMnO4 in most procedures.
Potassium dichromate: K2Cr2O7
Cr2O7
2- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O E° = 1.36 V
Primary standard, cheap, its solutions (orange) are stable. Cr3+ complexes
can be green to violet, so indicator necessary (or potentiometric).
Not as strong an oxidant as MnO4
- and Ce4+; mainly used for determination
of Fe2+ and for indirect determination of species that can oxidize Fe2+ to
Fe3+ (e.g. NO3
-, MnO4
-): a measured excess of Fe2+ is added to the
unknown, and excess Fe2+ is titrated with K2Cr2O7.
C3100 - Winter 2011
Dr. G. Van Biesen
23. 45
Methods Involving Iodine:
Iodimetry: titration using I2 (I3
-); I- is formed during the titration
Iodometry: titration of I2 (produced by a chemical reaction) with Na2S2O3
Molecular I2 is poorly soluble in water (1.3 x 10-3 M at 20 °C), but its solubility
increases by complexation with I- to form triiodide, and I2 solutions for titration
are prepared by dissolving I2 in water and adding excess I-:
I2 + I- I3
- K = 700 I3
- + 2e- 3 I- E° = 0.535 V
Sublimed I2 is a primary standard, but because of its high vapour pressure, it
is not practical to weigh it accurately. Instead, an approximate amount is
weighed, and dissolved with excess KI, and the solution is standardized with
Na2S2O3 or As4O6:
I3
- + 2 S2O3
2- S4O6
2- + 3 I-
As4O6 + 6 H2O 4 H3AsO3
H3AsO3 + I3
⎯ + H2O H3AsO4 + 3 I⎯ + 2 H+ Titration performed in
bicarbonate buffer to
keep [H+] low
46
IO3
- + 8 I- + 6 H+ 3 I3
- + 3 H2O
Alternatively, KIO3 can be added to a small excess of KI, then strong acid is
added to bring the solution to pH ≈ 1, and I3
- is produced by quantitative
reverse disproportionation:
Prepared solutions are stable at neutral pH (in the absence of heat, light…),
but not at acidic (6 I- + O2 + 4 H+ 2 I3
- + 2 H2O - slow) and basic pH
(disproportionation of I3
- to HOI, IO3
-, I-).
Na2S2O3 is an almost universal titrant for triiodide. Titrations are carried
out in neutral or acidic solution (pH < 9) to prevent disproportionation of I3
-.
Na2S2O3 is not a primary standard, but is standardized with I3
- prepared
from KIO3, or from a I3
- solution standardized with As4O6.
The universal indicator for iodi- and iodometry is starch. The amylose
portion of starch (which is helical) binds iodine (as I6 chains) to form a dark
blue complex. Best at low temperature.
Has to be fresh, or preservative (HgCl2) has to be added – hydrolysis
product is glucose, which is a reducing agent (!).
C3100 - Winter 2011
Dr. G. Van Biesen
24. 47
Iodimetry: reducing agent + I3
- 3 I-
Can add starch at the beginning – first drop of excess I3
- turns solution blue
Iodometry: oxidizing agent + 3 I- I3
-
Here, an excess of I- is added to the solution containing the analyte; the I3
-
formed is then titrated with Na2S2O3. Starch is added immediately before EP
(as indicated by fading I3
- color), otherwise some I2 will remain bound to
starch after EP.
48
Applications of iodimetry: analysis of SO2, As3+, Sn2+, formaldehyde, glucose,
vitamin C…
Some of these analyses are back-titrations, whereby a known excess of I3
- is
added to the analyte solution, and the remaining I3
- is titrated with Na2S2O3
Applications of iodometry: analysis of Cl2, Br2, MnO4
-, Cr2O7
2-…
In all these applications, excess I- is initially added to the analyte solution to
reduce the analyte and generate a stoichiometric amount of I3
- which is
titrated with Na2S2O3
ELECTROANALYTICAL TECHNIQUES (Chapter 16)
(a brief, qualitative overview of Chapter 16)
Previous chapters dealt with potentiometry: measuring voltage in the absence
of any significant current. Following techniques all involve a significant current
to force chemical reactions at an electrode surface (i.e. electrolysis).
C3100 - Winter 2011
Dr. G. Van Biesen
25. 49
Electrogravimetric Analysis:
Analyte is quantitatively deposited on an electrode
(cathode) by electrolysis. The electrode is weighed
before and after deposition.
e.g.: Cu2+ + 2e- Cu(s)
Check for completion of the reaction by
disappearance of color, or by exposing fresh
surface of the cathode to the solution and
checking if deposit forms with continuing
electrolysis.
Coulometric Titration:
Controlled delivery of electrons to form a reagent (titrant) in situ.
e.g.: Br2 + + 2 Br-
Br
Br
Br2 is generated by oxidation of Br- (in large excess in solution) at a Pt anode
Analysis of cyclohexene
by titration with Br2
Fig. 16-5,
Harris 8th Ed.
50
When just enough Br2 has been formed to
react with all the cyclohexene, the moles of e-
liberated = 2 x moles of cyclohexene.
A constant current is applied (controlled via
hand-operated switch) between the two
generator electrodes
#moles of e- = I • t / F
End point: excess Br2 detected by measuring
current between two detector electrodes:
anode: 2 Br- Br2 + 2 e-
cathode: Br2 + 2 e- 2 Br-
(no current before any Br2 remains in solution)
Other types of titrations:
acids - via OH- generated from: 2 H2O + 2 e- 2 OH- + H2(g)
precipitation – via Ag+ generated at a Ag anode
End point detection can be via color change of an indicator, or potentiometry.
C3100 - Winter 2011
Dr. G. Van Biesen
26. 51
Amperometry:
Electric current I between pair of electrodes that drive an electrolysis reaction
is measured – analyte is one of the reactants and I ~ [analyte].
There are numerous variations on this technique.
Example: blood glucose monitor:
Working electrode 1 is coated with the enzyme glucose oxidase
(and a mediator):
52
Early glucose monitors measured
[H2O2] by oxidation at a single
working electrode held at +0.6 V vs.
Ag / AgCl:
H2O2 O2(g) + 2 H+ + 2e-
I ~ [H2O2] ~ [glucose]
Problem: depends on pO2 in enzyme layer, which is not well controlled.
Mediator substitutes for O2 (shuttles e- between glucose and electrode, is
regenerated).
Another problem: Vitamin C, acetaminophen (Tylenol), uric acid are also
oxidized at the applied potential. This is corrected for by having a 2nd working
electrode coated with mediator, but not with glucose oxidase. These species
are thus oxidized at this electrode as well, but not glucose. The current due
to glucose is Ielectrode1 – Ielectrode2.
Voltammetry:
Number of techniques where the relationship between current and voltage is
observed during electrochemical processes. These techniques include:
C3100 - Winter 2011
Dr. G. Van Biesen
27. 53
Polarography
Uses a dropping-Hg electrode
(fresh Hg surface = reproducible I – V behaviour)
Qualitative (identify analyte by half-wave potential) and quantitative
(diffusion current ~ [analyte])
Stripping Analysis
Analytes are first reduced from a solution onto an electrode (Hg)
Analytes are then oxidized by applying a positive potential – I ~ [analyte]
Most sensitive voltammetric technique.
C3100 - Winter 2011
Dr. G. Van Biesen
