Analytical Electrochemistry : Potentiometry

1
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

Introduction This module provides an
introduction to the measurement technique of
potentiometry. It is intended to be a primary
learning tool for a student in a Quantitative
Analysis or Analytical Chemistry course and as a
review resource for a student in Instrumental
Analysis. It could also serve as a beginning
resource for new practitioners. If you have ever
used a pH meter, then you have already performed
potentiometry, an electrochemical method in which
the potential of an electrochemical cell is
measured while little to no current is passed
through the sample. In a potentiometric
measurement, an indicator electrode responds to
changes in the activity, or effective
concentration of the analyte. A potential, or
voltage, that develops at the interface between
the electrode and the analyte solution is
measured relative to a reference electrode. This
potential will be proportional to the amount of
analyte in the sample. An illustration of one
type of indicator electrode, the hydrogen ion or
pH electrode, is shown in the upper right-hand
corner of this page. Click here to continue
reading the introduction.
Erin M. Gross1, Richard S. Kelly2, and Donald M.
Cannon, Jr.3 1Department of Chemistry,
Creighton University, Omaha, NE 2Department of
Chemistry, East Stroudsburg University, East
Stroudsburg, PA 3Department of Chemistry,
University of Iowa, Iowa City, IA This work is
licensed under a Creative Commons
Attribution-Noncommercial-Share Alike 2.5 License
2
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

• Goals and Objectives
• After completion of this e-Module, you should be
  able to
• Describe the basic concepts of making a
  potentiometric measurement.
• Name some applications of potentiometry.
• Know the difference between a reference electrode
  and an indicator electrode.
• Describe the reactions of the typical reference
  electrodes.
• Define liquid junction potential and boundary
  potential.
• Describe how ion-selective electrodes (ISEs)
  function.
• Describe how both a pH electrode and a pH meter
  work.
• Describe the errors involved in pH measurements.
• Perform basic troubleshooting while making a pH
  measurement.
• Use the Nernst equation to perform calculations
  for potentiometric measurements.
• Click here to get started!

3
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

Potentiometry Timeline Shown below are
major milestones in the development of
potentiometry. Additional information is
available in the references cited.
Adapted from references 1- 7.
Click here for Potentiometric Theory.
4
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References

• Potentiometric Theory
• The origin of the measured potential at an
  indicator electrode is
• most generally the separation of charge across an
  interface
• between solutions of differing ionic strengths
  (an inner solution
• at fixed analyte activity and an outer solution
  with variable
• analyte activity).
• The mechanism leading to this charge separation
  varies with
• electrode type. After defining what is meant by
  a junction
• potential, we will consider two types of
  indicator electrodes
• the metallic direct indicator electrode, whose
  response involves a surface or solution redox
  reaction, and
• the membrane electrode, or ion-selective
  electrode (ISE).
• Click here to learn about junction potentials.

5
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References
Junction Potentials A potential develops at
any interface, or junction, where there is a
separation of charge. For example, a potential
can develop when a metal electrode comes in
contact with a solution containing its cation. A
potential of this type can be described using the
Nernst Equation. A potential can also develop
when electrolyte solutions of differing
composition are separated by a boundary, such as
a membrane or a salt bridge (a gel-filled tube
containing an inert electrolyte that connects
half-cells to allow charge neutrality to be
maintained). The two solutions may contain the
same ions, just at different concentrations or
may contain different ions altogether. These ions
have different mobilities, which means that they
move at different rates. Click here for more
about junction potentials.
6
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References
Direct Indicator Electrodes The simplest
type of direct indicator electrode is a metal, M,
in contact with a solution containing its own
cation, M. At the metal-solution interface, a
potential develops that is proportional to the
activity of the metal ion in solution. The
potential can be measured directly with respect
to a reference electrode using the
simple arrangement shown at right. Inert metal
electrodes like Pt or Au can be used as
indicator electrodes for ions involved in redox
reactions that occur in solution but do not
include the metallic form of the analyte. Click
here for more on direct indicator electrodes.
7
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References
Ion-Selective Electrodes Click
here to learn more about ion-selective
electrodes.
So far you have learned that in the technique of
potentiometry, the potential, or voltage, of an
electrochemical cell is measured. The cell
consists of both an indicator and reference
electrode. Since the potential of the reference
electrode is constant, it is the potential
developed at the indicator electrode that
contains information about the amount of analyte
in a sample. During the measurement, there is
little to no current flow. An electrochemical
cell for making a potentiometric measurement
with a membrane electrode (also known as an
ion-selective electrode, ISE) is shown in the
figure to the right. As you can see the main
difference between an ISE and the direct
indicator electrode is in the ISEs composition.
8
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References
Reference Electrodes It should be clear by
now that at least two electrodes are necessary to
make a potential measurement. As Kissinger and
Bott have so perfectly expressed,
electrochemistry with a single electrode is like
the sound of one hand clapping
(http//currentseparations.com/issues/20-2/20-2d.p
df). In potentiometry, those two electrodes are
generally called the indicator electrode and the
reference electrode. The indicator electrode
possesses some characteristic that allows it to
selectively respond to changes in the activity of
the analyte being measured. For the measured
potential to have meaning in this context, the
reference electrode must be constructed so that
its composition is fixed and its response is
stable over time, with observed changes in
measured potential due solely to changes in
analyte concentration. Click here to learn
more about reference electrodes.
9
Analytical Electrochemistry
Potentiometry
Introduction Goals and Objectives
Potentiometry Timeline Potentiometric Theory
Junction Potentials Direct Indicator
Electrodes Ion-Selective Electrodes
Reference Electrodes Nernst Equation
Instrumentation pH Electrodes Experiments
Common Troubleshooting Tips References
Nernst Equation The technique of
potentiometry involves the measurement of cell
potentials under conditions of no current flow.
In the electrochemical cell, if a high impedance
device like a voltmeter, is placed between the
two half cells, no current will flow between the
two compartments. As we have seen, it is
possible under these conditions to measure the
potential difference that exists between the two
electrodes. For cells with all reactants present
at unit activity, the measured cell potential
will be the standard cell potential, E0cell. In
real applications of potentiometry, reactant
activities are seldom (read never) equal to
unity, and measured cell potentials move away
from those that result from the tabulated values
of E0. A fundamental expression for
characterizing redox systems under equilibrium
conditions is the Nernst equation. One usually
has encountered this expression early in their
study of electrochemistry, perhaps in a general
chemistry course long ago. Click here to learn
more about the Nernst equation.
10
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

Instrumentation You have realized by now that
potentiometric measurements are fairly easy to
make from the standpoint of instrumentation. In
addition to the indicator electrode and the
reference electrode, the only remaining component
is a device used to measure the potential
difference that exists between the two
electrodes. If you have been with us to this
point, you should remember that potentiometric
measurements are ideally made under conditions of
very little current flow. This means that the
resistance (impedance to current flow) in the
electrochemical cell must be very high (up to 100
MW). This is usually not a problem due to the
nature of the indicator electrode, but the
measurement of potential under these conditions
requires the use of a device whose input
resistance is even larger than the cell
resistance. Click here to learn more about
instrumentation.
11
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

pH Electrodes The most widely used
ion-selective electrode is the glass pH
electrode, which utilizes a thin glass membrane
that is responsive to changes in H activity. F.
Haber, in 1901, was the first person to observe
that the voltage of a glass membrane changed with
the acidity of a solution. In 1906, M. Cremer
observed the pH dependence of measured potential
across a thin glass membrane. Today, pH
sensitive glasses are manufactured primarily from
SiO2 which are connected via a tetrahedral
network with oxygen atoms bridging two silicon
atoms (see an interactive 3d structure at
(http//www.geo.ucalgary.ca/tmenard/crystal/quart
z.html). In addition, the glasses are made to
contain varying amounts of other metal oxides,
like Na2O and CaO. Oxygen atoms within the
lattice that are not bound to two silicon atoms
possess a negative charge, to which cations can
ion pair. In this way, ions (primarily Na) are
able to diffuse slowly in the lattice, moving
from one charge pair site to another. While the
membrane resistance is very high (100 MW), this
movement of cations within the glass allows a
potential to be measured across it.
Click here to learn more about pH electrodes.
12
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

• Experiments
• Potentiometric Titration of an Unknown Monoprotic
  Weak Acid
• Determination of Chloride Using Potentiometry
• Fluoride Ion by Direct Potentiometry/Standard
  Addition

13
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

Common Troubleshooting Tips While
understanding the underlying general concepts of
potentiometry is a useful first step at becoming
a regular "potentiometric practitioner",
experience is also a great resource for
effectively conducting these types of
measurements.  Through experience comes
familiarity with common "problem areas" of this
field.  This page is intended to present some
troubleshooting tips.  It is not our intention to
replace recommendations outlined in manufacturer
literature. Before specific discussion on common
problem areas, the subtle nuance differences
between efforts in  calibration methods and
quality control (QC) must be highlighted. 
Calibration and QC methods are complementary to
one another and are often integrated into a
method validation program that defines the
overall reliability.  Calibrations give
analytical methods an initial quantitative
starting point, whereas QC validates the
developed calibration model. Click here to
learn more about troubleshooting.
14
Analytical Electrochemistry
Potentiometry

• Introduction
• Goals and Objectives
• Potentiometry Timeline
• Potentiometric Theory
• Instrumentation
• pH Electrodes
• Experiments
• Common Troubleshooting Tips
• References

• References
• Cremer, M,. Z. Biol. 1906, 47, 562.
• Buck, R.P. and Lindner, E. Anal. Chem. 2001, 73,
  88A.
• Frant, M.S., Analyst, 1994, 119, 2293.
• Frant, M.S. and Ross, J.W., Science, 1966, 154,
  1553.
• Ross, J.W., Science, 1967, 156, 1378.
• Simon, W., Swiss Pat., 479870, 1969.
• Frant, M.S. and Ross, J.W., Science, 1970, 167,
  987.
• Bakker, E. and Pretsch, E., Anal. Chem., 2002,
  74, 420A.
• Meyerhoff, M.E. and Opdycke, W.N. In Advances in
  Clinical Chemistry, vol. 25 Spiegel, H.E., Ed.
  Academic Press, Inc., Orlando, 1986, pp. 1-47.
• Bakker, E. Buhlmann, P. Pretsch, E. Talanta,
  2004, 63, 3.
• Wang, J. Analytical Electrochemistry, 3rd ed.
  Wiley Hoboken, NJ 2006, pp. 165-200.
• Skoog, D.A., Holler, F.J., Crouch, S.R.,
  Principles of Instrumental Analysis, 6th ed.
  Thomson Brooks/Cole Belmont, CA, 2007.
• Buhlmann, P. Pretsch, E. Bakker, E. Chem Rev.
  1998, 98, 1593.