Transcript for Potentiometric Theory.

Analytical Electrochemistry :
Potentiometry
Introduction
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Instrumentation
pH Electrodes
Experiments
Common Troubleshooting Tips
References
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
of Chemistry, East Stroudsburg University, East Stroudsburg, PA
3Department of Chemistry, University of Iowa, Iowa City, IA
2Department
This work is licensed under a
Creative Commons Attribution-Noncommercial-Share Alike 2.5 License
Analytical Electrochemistry :
Potentiometry
Goals and Objectives
Introduction
Goals and Objectives
After completion of this e-Module, you should be able to:
Potentiometry Timeline
•
Describe the basic concepts of making a potentiometric
measurement.
Potentiometric Theory
•
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.
Instrumentation
pH Electrodes
Experiments
Common Troubleshooting Tips
References
Click here to get started!
Analytical Electrochemistry :
Potentiometry
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometry Timeline
Shown below are major milestones in the development of potentiometry.
Additional information is available in the references cited.
Potentiometric Theory
Instrumentation
pH Electrodes
Experiments
Common Troubleshooting Tips
References
Adapted from references 1- 7.
Click here for Potentiometric Theory.
Analytical Electrochemistry :
Potentiometry
Introduction
Potentiometric Theory
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Junction Potentials
Direct Indicator Electrodes
Ion-Selective Electrodes
Reference Electrodes
Nernst Equation
Instrumentation
pH Electrodes
Experiments
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:
1) the metallic direct indicator electrode, whose response
involves a surface or solution redox reaction, and
2) the membrane electrode, or ion-selective electrode (ISE).
Common Troubleshooting Tips
References
Click here to learn about junction potentials.
Analytical Electrochemistry :
Potentiometry
Junction Potentials
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Junction Potentials
Direct Indicator Electrodes
Ion-Selective Electrodes
Reference Electrodes
Nernst Equation
Instrumentation
pH Electrodes
Experiments
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.
Common Troubleshooting Tips
References
Click here for more about junction potentials.
Analytical Electrochemistry :
Potentiometry
Direct Indicator Electrodes
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
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.
0.319 V
Voltmeter
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.
Reference
electrode
M
Click here for more on direct
indicator electrodes.
M+ (aq)
Analytical Electrochemistry :
Potentiometry
Ion-Selective Electrodes
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
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 ISE’s composition.
Click here to learn more about ion-selective electrodes.
Analytical Electrochemistry :
Potentiometry
Introduction
Reference Electrodes
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
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.pdf). 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.
Analytical Electrochemistry :
Potentiometry
Nernst Equation
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
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.
Analytical Electrochemistry :
Potentiometry
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Instrumentation
pH Electrodes
Experiments
Common Troubleshooting Tips
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.
References
Click here to learn more about instrumentation.
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/quartz.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.
Analytical Electrochemistry :
Potentiometry
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Instrumentation
pH Electrodes
Experiments
Common Troubleshooting Tips
References
Experiments
1. Potentiometric Titration of an Unknown Monoprotic
Weak Acid
2. Determination of Chloride Using Potentiometry
3. Fluoride Ion by Direct Potentiometry/Standard
Addition
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.
Analytical Electrochemistry :
Potentiometry
References
Introduction
Goals and Objectives
Potentiometry Timeline
Potentiometric Theory
Instrumentation
pH Electrodes
Experiments
1.
Cremer, M,. Z. Biol. 1906, 47, 562.
2.
Buck, R.P. and Lindner, E. Anal. Chem. 2001, 73, 88A.
3.
Frant, M.S., Analyst, 1994, 119, 2293.
4.
Frant, M.S. and Ross, J.W., Science, 1966, 154, 1553.
5.
Ross, J.W., Science, 1967, 156, 1378.
6.
Simon, W., Swiss Pat., 479870, 1969.
7.
Frant, M.S. and Ross, J.W., Science, 1970, 167, 987.
8.
Bakker, E. and Pretsch, E., Anal. Chem., 2002, 74, 420A.
9.
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.
Common Troubleshooting Tips
10. Bakker, E.; Buhlmann, P.; Pretsch, E. Talanta, 2004, 63, 3.
References
11. Wang, J. Analytical Electrochemistry, 3rd ed.; Wiley: Hoboken, NJ
2006, pp. 165-200.
12. Skoog, D.A., Holler, F.J., Crouch, S.R., Principles of Instrumental
Analysis, 6th ed.; Thomson Brooks/Cole: Belmont, CA, 2007.
13. Buhlmann, P.; Pretsch, E.; Bakker, E. Chem Rev. 1998, 98, 1593.