14-6 Cells as chemical probes
• It is important to distinguish two types of equilibrium in a galvanic
cell: equilibrium between the two half-cells and equilibrium within
each half cell.
A galvanic cell that can be
used to measure the formation
constant of Hg(EDTA)2Hg2+ + 2e- ↔ Hg(l)
Eo = 0.852 V
14-7 Biochemists use Eo’
• Formal potential: Potential of a half-reaction (relative to
a standard hydrogen electrode) when the formal
concentrations of reactants and products are unity. Any
other conditions (such as pH, ionic strength, and
concentrations of ligands) also must be specified.
Finding the formal potential
• Reduction potential of ascorbic
acid, showing its dependence
on pH.
(a) Graph of the function labeled
formal potential.
(b) Experimental polarographic
half-wave reduction potential
of ascorbic acid in a medium of
ionic strength = 0.2 M.
At high pH (>12), the half-wave
potential does not level off to a
slope of 0, as Equation 14-34
predicts. Instead, a hydrolysis
reaction of ascorbic acid
occurs and the chemistry is
more complex.
•
Example. Cadmium electrode immersed in a solution that is 0.015M in
Cd2+ .
Solution:
Cd2+ + 2e  Cd(s) 0 = -0.403
ε = εo -
0.0591
1
log
2
Cd 2+ 
0.0591
1
ε = -0.403 log
  0.457 V
2
0.0150
• Example: Calculate the potential for a platinum electrode
immersed in a solution prepared by saturating a 0.015 M
solution of KBr with Br2
Br2(l) + 2e  2BrE0 = 1.065 V
Solution:
 2

Br 
0.0591

E= 1.065 log
2
1.0
E = 1.065 -
0.0591
2
log  0.0150   1.173V
2
Important notes
1. When you multiply a reaction by a coefficient, the potential of
such a reaction does not change!
2. When you subtract or add reactions, the Gibbs energy of those reactions
can always be directly added or subtracted, not their potentials!
Chapter 15: Electrodes and
potentiometry
A few definition
• Potentiometry: the use of electrodes to measure
voltages that provide chemical information.
• Electroactive species: a substance that can accept or
donate electrons at an electrode.
• Indicator electrode: one that develops a potential whose
magnitude depends on the activity of one or more
species in contact with the electrode.
• Reference electrode: one that maintains constant
potential against which the potential of another half-cell
may be measured.
15-1 Reference electrodes
• A galvanic cell that can monitor
the quotient [Fe2+]/[Fe3+] in the
right half-cell.
The electrode potential
E+ = 0.771 – 0.05916 log([Fe2+]/[Fe3+])
E- = 0.222 – 0.05916 log([Cl-])
E = E + - E-
• The Pt wire is the indicator
electrode.
• Silver-silver chloride electrode:
a reference electrode.
Reference electrodes
Silver-Silver Chloride
Reference Electrode
Calomel electrode
•
Figure shows a double-junction
electrode that minimizes
contact between analyte
solution and KCl from the
electrode.
• A problem with reference electrodes is that porous plugs become
clogged, thus causing sluggish, unstable electrical response.
Some designs incorporate a free-flowing capillary in place of the
porous plug.
Potential for the solution saturated with KCl (is not 0.222
V):
AgCl(s) + e- ↔ Ag(s) + ClEo = -0.197 V
• A Hg electrode saturated with KCl is called a saturated calomel
electrode, abbreviated S.C.E. The advantage in using saturated
KCl is that [Cl−] does not change if some liquid evaporates.
Ecalomel  saturated with KCl   0.241 V
Voltage conversions between different
reference scales
• If an electrode has a potential of −0.461 V with
respect to a calomel electrode, what is the
potential with respect to a silver-silver chloride
electrode? What would be the potential with
respect to the standard hydrogen electrode?
15-2 Indicator electrode
• Metal electrodes: such as Pt, Au, Ag, etc.
• Ion-selective electrodes:
• Carbon electrode
Example: Use of Ag and calomel electrodes
to measure [Ag+].

 1 


   0.241
E  0.799  0.0591 log 

  Ag   

 


Why do we need double junction
calomel reference electrode here?
15-3 What is a junction potential
• Whenever dissimilar electrolyte solutions are in contact, a voltage
difference called the junction potential develops at their interface.
This small voltage (usually a few millivolts) is found at each end of a
salt bridge connecting two half-cells. The junction potential puts a
fundamental limitation on the accuracy of direct potentiometric
measurements.
• The steady-state junction potential represents a balance between
the unequal mobilities.
Development of the junction potential
caused by unequal mobilities of Na+ and
Cl−.
•
Junction potential:
An electric potential
that exists at the
junction between two
different electrolyte
solutions or
substances. It arises in
solutions as a result of
unequal rates of
diffusion of different
ions.
• A 0.1 M NaCl
solution was placed
in contact with a 0.1
M NaNO3. Which
side of the junction
is positive?
15-4 How ion-selective electrodes
work
• Respond selectively to one ion.
• Do not involve redox processes.
• Features a thin membrane capable of
binding only the intended ion.
• The electric potential difference across the
ion-selective membrane is measured with
two reference electrodes.
15-5 pH measurements with a Glass
electrode
•
The glass electrode used to
measure pH is the most common
ion-selective electrode.
•
A typical pH combination
electrode, incorporating both
glass and reference electrodes in
one body.
•
Glass combination electrode with
a silver-silver chloride reference
electrode. The glass electrode is
immersed in a solution of
unknown pH so that the porous
plug on the lower right is below
the surface of the liquid. The two
silver electrodes measure the
voltage across the glass
membrane.
• The potential difference between inner and outer silver-silver
chloride electrodes depends on the chloride concentration in each
electrode compartment and on the potential difference across the
glass membrane.
• Because [Cl−] is fixed in each compartment and because [H+] is
fixed on the inside of the glass membrane, the only variable is the
pH of analyte solution outside the glass membrane.
• The voltage of the ideal pH electrode changes by 59.16 mV for
every pH-unit change of analyte activity at 25°C.
Errors in pH measurement
1.
2.
3.
4.
5.
6.
7.
8.
9.
Standards.
Junction potential
Junction potential drift.
Sodium error.
Acid error.
Equilibration time.
Hydration of glass.
Temperature.
Cleaning.
Ion-selective electrodes-example
Most ion-selective electrodes fall into one of the following classes:
1. Glass membranes for H+ and certain monovalent cations
2. Solid-state electrodes based on inorganic salt crystals
3. Liquid-based electrodes using a hydrophobic polymer membrane
saturated with a hydrophobic liquid ion exchanger
4. Compound electrodes with a species-selective electrode enclosed by
a membrane that is able to separate that species from others or that
generates the species in a chemical reaction.
Ion-selective electrodes-example
• Electrodes that respond
selectively to specific analytes
in solution or in the gas phase.
• Ion exchange between heparin
and Cl− associated with
tetraalkylammonium ions in the
membrane of the ion-selective
electrode.
Rotating ion-selective heparin electrode
Potentiometry with an Oscillating Reaction
• The Belousov-Zhabotinskii reaction is a cerium-catalyzed oxidation
of malonic acid by bromate, in which the quotient [Ce3+]/[Ce4+]
oscillates by a factor of 10 to 100.
•
•
•
•
•
Oscillation between yellow and colorless is set up in a 300-mL beaker with
the following solutions:
160 mL of 1.5 M H2SO4
40 mL of 2 M malonic acid
30 mL of 0.5 M NaBrO3 (or saturated KBrO3)
4 mL of saturated ceric ammonium sulfate, (Ce(SO4)2 · 2(NH4)2SO4 · 2H2O)
Trace b shows two different cycles
superimposed in the same solution.
This unusual event occurred
in a reaction that had been oscillating
normally for about 30 min