Ebook Analytical chemistry (7th edition) Part 2

Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 399

Chapter Thirteen
POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

Chapter 13 URLs

Learning Objectives
WHAT ARE SOME OF THE KEY THINGS WE WILL LEARN FROM THIS CHAPTER?
●

Types of electrodes and electrode potentials from the Nernst
equation (key equations: 13.3, 13.10, 13.16), pp. 400, 401, 402

●

The pH glass electrode (key equation: 13.42), p. 413

●

Liquid junctions and junction potentials, p. 405

●

Standard buffers and the accuracy of pH measurements,
pp. 418

●

Reference electrodes, p. 407

●

The pH meter, p. 421

●

Accuracy of potentiometric measurements (key equation:
13.36), p. 412

●

Ion-selective electrodes, p. 424

●

The selectivity coefficient (key equation: 13.46), p. 428

In Chapter 12, we mentioned measurement of the potential of a solution and described
a platinum electrode whose potential was determined by the half-reaction of interest.
This was a special case, and there are a number of electrodes available for measuring
solution potentials. In this chapter, we list the various types of electrodes that can be
used for measuring solution potentials and how to select the proper one for measuring
a given analyte. The apparatus for making potentiometric measurements is described
along with limitations and accuracies of potentiometric measurements. The important
glass pH electrode is described, as well as standard buffers required for its calibration.
The various kinds of ion-selective electrodes are discussed. The use of electrodes in
potentiometric titrations is described in Chapter 14.
Potentiometric electrodes measure activity rather than concentration, a unique

feature, and we will use activities in this chapter in describing electrode potentials.
An understanding of activity and the factors that affect it are important for direct
potentiometric measurements, as in pH or ion-selective electrode measurements. You
should, therefore, review the material on activity and activity coefficients in Chapter 6.
Potentiometry is one of the oldest analytical methods, with foundations of
electrode potentials and electrochemical equilibria laid down by J. Willard Gibbs
(1839–1903) and Walther Nernst (1864–1941). Inert electrodes are used as indicating
electrodes for redox titrations, and may be used in automatic titrators. The pH electrode
is the most widely used potentiometric electrode. Ion selective electrodes are now
more widely used than redox electrodes, for selectively measuring particular ions.
The measurement of fluoride, for example in toothpaste, is one of the more important
applications since fluoride is not easily measured otherwise. Clinical analyzers measure
the electrolytes sodium, potassium, lithium (used in the treatment of manic depression),
and calcium in blood using ion selective electrodes.

Review activities in Chapter 6, for
an understanding of
potentiometric measurements.

399


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 400

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

400

13.1 Metal Electrodes for Measuring the Metal Cation
An electrode of this type is a metal in contact with a solution containing the cation of
the same metal. An example is a silver metal electrode dipping in a solution of silver
nitrate.
For all electrode systems, an electrode half-reaction can be written from which
the potential of the electrode is described. The electrode system can be represented by
M/Mn+ , in which the line represents an electrode–solution interface. For the silver
electrode, we have
(13.1)
Ag|Ag+
and the half-reaction is

Ag+ + e−

Ag

(13.2)

The potential of the electrode is described by the Nernst equation:
0
E = EAg
+ ,Ag −

Increasing cation activity always
causes the electrode potential to
become more positive (if you
write the Nernst equation
properly).
The indicator electrode is the one
that responds to the analyte.

1
2.303RT
log
nF
aAg+

(13.3)

where aAg+ represents the activity of the silver ion (see Chapter 6). The value of n
here is 1. We will use the more correct unit of activity in discussions in this chapter
because, in the interpretation of direct potentiometric measurements, significant errors
would result if concentrations were used in calculations.
The potential calculated from Equation 13.3 is the potential relative to the normal
hydrogen electrode (NHE—see Section 13.3). The potential becomes increasingly
positive with increasing Ag+ (the case for any electrode measuring a cation). That is,
in a cell measurement using the NHE as the second half-cell, the voltage is
Emeasd. = Ecell = Eind vs. NHE = Eind − ENHE

(13.4)

where Eind is the potential of the indicator electrode (the one that responds to the test
solution, Ag+ ions in this case). Since ENHE is zero,
Ecell = Eind

(13.5)

Eref | solution | Eind

(13.6)

Ecell = Eright − Eleft = Eind − Eref = Eind − constant

(13.7)

corresponds to writing the cells as
and
The reference electrode completes
the cell but does not respond to the
analyte. It is usually separated
from the test solution by a salt
bridge.

Any pure substance does not
numerically appear in the Nernst
equation (e.g., Cu, H2 O); their
activities are taken as unity.

where Eref is the potential of the reference electrode, whose potential is constant. Note
that Ecell (or Eind ) may be positive or negative, depending on the activity of the silver
ion or the relative potentials of the two electrodes. This is in contrast to the convention
used in Chapter 12 for a voltaic cell, in which a cell was always set up to give a
positive voltage and thereby indicate what the spontaneous cell reaction would be. In
potentiometric measurements, we, in principle, measure the potential at zero current
so as not to disturb the equilibrium, i.e., don’t change the relative concentrations of
the species being measured at the indicating electrode surface—which establishes the
potential (see measurement of potential, below). We are interested in how the potential
of the test electrode (indicating electrode) changes with analyte concentration, as
measured against some constant reference electrode. Equation 13.7 is arranged so that
changes in Ecell reflect the same changes in Eind , including sign. This point is discussed

further when we talk about cells and measurement of electrode potentials.
The activity of silver metal above, as with other pure substances, is taken as
unity. So an electrode of this kind can be used to monitor the activity of a metal ion in
solution. There are few reliable electrodes of this type because many metals tend to
form an oxide coating that changes the potential.


Christian7e c13.tex

13.2 METAL–METAL SALT ELECTRODES FOR MEASURING THE SALT ANION

V2 - 08/13/2013 1:49 P.M. Page 401

401

13.2 Metal–Metal Salt Electrodes for Measuring the Salt Anion
The general form of this type of electrode is M|MX|Xn− , where MX is a slightly
soluble salt. An example is the silver–silver chloride electrode:
Ag|AgCl(s) |Cl−

(13.8)

The (s) indicates a solid, (g) is used to indicate a gas, and (l) is used to indicate a pure
liquid. A vertical line denotes a phase boundary between two different solids or a solid
and a solution. The half-reaction is
AgCl + e−

Ag + Cl−

(13.9)

where the underline indicates a solid phase and the potential is defined by
0
E = EAgCl,Ag
−

2.303RT
log aCl−
F

(13.10)

The number of electrons, n, does not appear in the equation because here n = 1.
This electrode, then, can be used to measure the activity of chloride ion in
solution. Note that, as the activity of chloride increases, the potential decreases. This
is true of any electrode measuring an anion—the opposite for a cation electrode. A
silver wire is coated with silver chloride precipitate (e.g., by electrically oxidizing it in
a solution containing chloride ion, the reverse reaction of Equation 13.9). Actually, as
soon as a silver wire is dipped in a chloride solution, a thin layer of silver chloride and
is usually not required.
Note that this electrode can be used to monitor either aCl− or aAg+ . It really
senses only silver ion, and the activity of this is determined by the solubility of the
slightly soluble AgCl. Since aCl− = Ksp /aAg+ , Equation 13.10 can be rewritten:
Ksp
2.303RT
log
F
aAg+

(13.11)

1
2.303RT
2.303RT
log Ksp −
log
F
F
aAg+

(13.12)

0
−
E = EAgCl,Ag
0
−
E = EAgCl,Ag

Comparing this with Equation 13.3, we see that
0
0
EAg
+ ,Ag = EAgCl,Ag −

2.303RT
log Ksp
F

(13.13)

◦
(see Chapter 6), since activities,
Ksp here is the thermodynamic solubility product Ksp
rather than concentrations, were used in arriving at it in these equations. We could
have arrived at an alternative form of Equation 13.10 by substituting Ksp /aCl− for aAg+
in Equation 13.3 (see Example 13.1).
In a solution containing a mixture of Ag+ and Cl− (e.g., a titration of Cl− with
+
Ag ), the concentrations of each at equilibrium will be such that the potential of
a silver wire dipping in the solution can be calculated by either Equation 13.3 or
Equation 13.10. This is completely analogous to the statement in Chapter 12 that the
potential of one half-reaction must be equal to the potential of the other in a chemical
reaction at equilibrium. Equations 13.2 and 13.9 are the two half-reactions in this case,
and when one is subtracted from the other, the result is the overall chemical reaction.

Ag+ + Cl−
−

+

AgCl

(13.14)

Note that as Cl is titrated with Ag , the former decreases and the latter increases.
Equation 13.10 predicts an increase in potential as Cl− decreases; and similarly,
Equation 13.12 predicts the same increase as Ag+ increases.

Increasing anion activity always

causes the electrode potential to
decrease.

The Ag metal really responds to
Ag+ , whose activity is determined
◦
and aCl− .
by Ksp


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 402

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

402

The silver electrode can also be used to monitor other anions that form slightly
soluble salts with silver, such as I− , Br− , and S2− . The E0 in each case would be that
for the particular half-reaction AgX + e−
Ag + X− .
Another widely used electrode of this type is the calomel electrode, Hg,
Hg2 Cl2(s) |Cl− . This will be described in more detail when we talk about reference
electrodes.

Example 13.1
Given that the standard potential of the calomel electrode is 0.268 V and that of the
Hg/Hg2 2+ electrode is 0.789 V, calculate Ksp for calomel (Hg2 Cl2 ), for 298K.
Solution

For Hg2 2+ + 2e−

Hg,
1
0.05916
log
2
aHg2 2+

(1)

0.05916
log(aCl− )2
2

(2)

Ksp
0.05916
log
2
aHg2 2+

(3)

1
0.05916
0.05916
log Ksp −

log
2
2
aHg2 2+

(4)

E = 0.789 −
For Hg2 Cl2 + 2e−

2Hg + 2Cl− ,
E = 0.268 −

Since Ksp = aHg2 2+ · (aCl− )2 ,
E = 0.268 −
E = 0.268 −
From (1) and (4),
0.789 = 0.268 −

0.05916
log Ksp
2

Ksp = 2.44 × 10−18

13.3 Redox Electrodes—Inert Metals
In the redox electrode, an inert metal is in contact with a solution containing the
soluble oxidized and reduced forms of the redox half-reaction. This type of electrode
was mentioned in Chapter 12.
The inert metal used is usually platinum. The potential of such an inert electrode

is determined by the ratio at the electrode surface of the reduced and oxidized species
in the half-reaction:
Ma+ + ne−
M(a−n)+
(13.15)
0
E = EM
a+ ,M(a−n)+ −

a (a−n)+
2.303RT
log M
nF
aMa+

(13.16)

An example is the measurement of the ratio of MnO4 − /Mn2+ :
MnO4 − + 8H+ + 5e−
0
E = EMnO

4

− ,Mn2+

−

Mn2+ + 4H2 O

aMn2+
2.303RT
log
5F
aMnO4 − · (aH+ )8

(13.17)
(13.18)


Christian7e c13.tex

13.3 REDOX ELECTRODES—
—INERT METALS

V2 - 08/13/2013 1:49 P.M. Page 403

403

H2

HCI solution

H2

Platinized Pt electrode

Fig. 13.1.

Hydrogen electrode.

The pH is usually held constant, and so the ratio aMn2+ /aMnO4 − is measured, as in a
redox titration.
A very important example of this type of electrode is the hydrogen electrode,
Pt|H2 , H+ :
1
(13.19)
H+ + e−
2 H2
0
E = EH
+ ,H −
2

(pH2 )1/2
2.303RT
log
F
aH+

(13.20)

The construction of the hydrogen electrode is shown in Figure 13.1. A layer of
platinum black must be placed on the surface of the platinum electrode by cathodically
electrolyzing in a H2 PtCl6 solution. The platinum black provides a larger surface area
for adsorption of hydrogen molecules and catalyzes their oxidation. Too much platinum
black, however, can adsorb traces of other substances such as organic molecules or
H2 S, causing erratic behavior of the electrode.
The pressure of gases, in atmospheres, is used in place of activities. If the
hydrogen pressure is held at 1 atm, then, since E0 for Equation 13.19 is defined

as zero,
1
2.303RT
2.303RT
log
pH
E=
=−
(13.21)
F
aH+
F

For gases, we will use pressures, p
(in atmospheres), in place of
activity (or the thermodymic
equivalent term for gases,
fugacity).

Example 13.2
Calculate the pH of a solution whose potential at 25◦ C measured with a hydrogen
electrode at an atmospheric pressure of 1.012 atm (corrected for the vapor pressure of
water at 25◦ C) is −0.324 V (relative to the NHE).
Solution

From Equation 13.20,
−0.324 = −0.05916 log

(1.012)1/2
aH+

= −0.05916 log(1.012)1/2 − 0.05916 pH
pH = 5.48

The vapor pressure of water above
the solution must be subtracted
from the measured gas pressure.


Christian7e c13.tex

404

V2 - 08/13/2013 1:49 P.M. Page 404

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

While the hydrogen electrode is very important for specific applications (e.g.,
establishing standard potentials or the pH of standard buffers—see below), its use for
routine pH measurements is limited. First, it is inconvenient to prepare and use. The
partial pressure of hydrogen must be established at the measurement temperature. The
solution should not contain other oxidizing or reducing agents since these will alter
the potential of the electrode.

13.4 Voltaic Cells without Liquid
Junction—For Maximum Accuracy

It is possible to construct a cell
without a salt bridge. For practical
purposes, this is rare because of

the tendency of the reference
electrode potential to be
influenced by the test solution.

This cell is used to accurately
measure the pH of “standard
buffers.” See Section 13.12.

To make potential measurements, a complete cell consisting of two half-cells must
be set up, as was described in Chapter 12. One half-cell usually is comprised of
the test solution and an electrode whose potential is determined by the analyte we
wish to measure. This electrode is the indicator electrode. The other half-cell is
any arbitrary half-cell whose potential is not dependent on the analyte. This half-cell
electrode is designated the reference electrode. Its potential is constant, and the
measured cell voltage reflects the indicator electrode potential relative to that of the
reference electrode. Since the reference electrode potential is constant, any changes
in potential of the indicator electrode will be reflected by an equal change in the
cell voltage.
There are two basic ways a cell may be set up, either without or with a salt
bridge. The first is called a cell without liquid junction. An example of a cell of this
type would be
Pt|H2 (g), HCl(solution)|AgCl(s)|Ag

(13.22)

The solid line represents an electrode–solution interface. An electrical cell such as
this is a voltaic one, and the cell illustrated above is written for the spontaneous
reaction by convention (positive Ecell —although we may actually measure a negative
cell voltage if the indicator electrode potential is the more negative one; we haven’t
specified which of the half-reactions represents the indicator electrode). The hydrogen

electrode is the anode, since its potential is the more negative (see Chapter 12 for a
review of cell voltage conventions for voltaic cells). The potential of the left electrode
would be given by Equation 13.20, and that for the right electrode would be given
by Equation 13.10, and the cell voltage would be equal to the difference in these
two potentials:
⎛
⎞
1
2
pH2
2.303RT
2.303RT
0
0
⎠
log aCl− − ⎝EH
log
−
Ecell = EAgCl,Ag
+ ,H −
2
F
F
aH+
(13.23)
Ecell =

0
EAgCl,Ag

−

0
EH
+ ,H
2

a +a −
2.303RT
log H Cl1
−
F
(pH2 ) 2

(13.24)

The cell reaction would be (half-reaction)right − (half-reaction)left (to give a positive
Ecell and the spontaneous reaction), or
AgCl + e−
−(H+ + e−
AgCl + 12 H2

Ag + Cl−

(13.25)

1
2 H2 )

(13.26)

Ag + Cl− + H+

(13.27)


Christian7e c13.tex

13.5 VOLTAIC CELLS WITH LIQUID JUNCTION—
—THE PRACTICAL KIND

V2 - 08/13/2013 1:49 P.M. Page 405

405

Equation 13.23 would also represent the voltage if the right half-cell were used
as an indicating electrode in a potentiometric measurement of chloride ion and the left
cell were the reference electrode (see Equations 13.6 and 13.7). That is, the voltage
(and hence the indicator electrode potential) would decrease with increasing chloride
ion. If we were to use the hydrogen electrode as the indicating electrode to measure
hydrogen ion activity or pH, we would reverse the cell setup in Equation 13.22 from
left to right to indicate what is being measured. Equation 13.23 will be reversed as
well, and the voltage (and indicator electrode potential) would increase with increasing
acidity or decreasing pH (Ecell = Eind − Eref , Equation 13.7).
Cells without liquid junction are always used for the most accurate measurements
because there are no uncertain potentials to account for and were used for measuring
the pH of NIST standard buffers (see below). However, there are few examples of
cells without liquid junction (sometimes called cells without transference), and they
are inconvenient to use. Therefore, the more convenient (but less accurate) cells with
liquid junction are commonly used.

13.5 Voltaic Cells with Liquid Junction—The Practical Kind
An example of this type of cell is
Hg|Hg2 Cl2 (s)|KCl(saturated)||HCl(solution), H2 (g)|Pt

(13.28)

The double line represents the liquid junction between two dissimilar solutions and
is usually in the form of a salt bridge. The purpose of this is to prevent mixing of
the two solutions. In this way, the potential of one of the electrodes will be constant,
independent of the composition of the test solution, and determined by the solution
in which it dips. The electrode on the left of cell 13.28 is the saturated calomel
electrode, which is a commonly used reference electrode (see below). The cell is set
up using the hydrogen electrode as the indicating electrode to measure pH.
LIQUID-JUNCTION POTENTIAL——WE CAN’T IGNORE THIS
The disadvantage of a cell of this type is that there is a potential associated with
the liquid junction, called the liquid-junction potential. The potential of the above
cells is
Ecell = (Eright − Eleft ) + Ej

(13.29)

where Ej is the liquid-junction potential; Ej may be positive or negative. The liquidjunction potential results from the unequal diffusion of the ions on each side of
the boundary. A careful choice of salt bridge (or reference electrode containing a
suitable electrolyte) can minimize the liquid-junction potential and make it reasonably
constant so that a calibration will account for it. The basis for such a selection is
discussed as follows.
A typical boundary might be a fine-porosity sintered-glass frit with two different
solutions on either side of it; the frit prevents appreciable mixing of the two solutions.
The simplest type of liquid junction occurs between two solutions containing the same

electrolyte at different concentrations. An example is HCl (0.1 M)||HCl (0.01 M),
illustrated in Figure 13.2. Both hydrogen ions and chloride ions will migrate across the
boundary in both directions, but the net migration will be from the more concentrated
to the less concentrated side of the boundary, the driving force for this migration
being proportional to the concentration difference. Hydrogen ions migrate about five

The presence of a liquid-junction
potential limits the accuracy of
potentiometric measurements.


Christian7e c13.tex

406

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

H + (0.1 M )

H + (0.01 M )
Equilibrium
Cl − (0.01 M )

Cl − (0.1 M )

Fig. 13.2.

Representation of
liquid-junction potential.

We minimize the liquid-junction
potential by using a high
concentration of a salt whose ions
have nearly equal mobility, for
example, KCl.

V2 - 08/13/2013 1:49 P.M. Page 406

Boundary

−
−
−
−

+
+
+
+

+40 mV

times faster than chloride ions. Therefore, a net positive charge is built up on the
right side of the boundary, leaving a net negative charge on the left side; that is,
there is a separation of charge, and this represents a potential. A steady state is
rapidly achieved by the action of this build-up positive charge in inhibiting the further
migration of hydrogen ions; the converse applies to the negative charge on the lefthand side. Hence, a constant potential difference is quickly attained between the
two solutions.
The Ej for this junction is +40 mV, and Ecell = (Eright − Eleft ) + 40 mV. This Ej
is very large, owing to the rapid mobility of the hydrogen ion. As the concentration of

HCl on the left side of the boundary is decreased, the net charge built up will be less,
and the liquid-junction potential will be decreased until, at equal concentration, it will
be zero, because equal amounts of HCl diffuse in each direction.
A second example of this type of liquid junction is 0.1 M KCl/0.01 M KCl. This
situation is completely analogous to that above, except that in this case the K+ and
Cl− ions migrate at nearly the same rate, with the chloride ion moving only about 4%
faster. So a net negative charge is built up on the right side of the junction, but it will
be relatively small. Thus, Ej will be negative and is equal to −1.0 mV.
HOW DO WE MINIMIZE THE LIQUID-JUNCTION POTENTIAL?
The nearly equal migration of potassium and chloride ions makes it possible to
significantly decrease the liquid-junction potential. This is possible because, if an
electrolyte on one side of a boundary is in large excess over that on the other side,
the flux of the migration of the ions of this electrolyte will be much greater than that
of the more dilute electrolyte, and the liquid-junction potential will be determined
largely by the migration of this more concentrated electrolyte. Thus, Ej of the junction
KCl (3.5 M)|| H2 SO4 (0.05 M) is only −4 mV, even though the hydrogen ions diffuse
at a much more rapid rate than sulfate.
Some examples of different liquid-junction potentials are given in Table 13.1.
(The signs are for those as set up, and they would be the signs in a potentiometric
measurement if the solution on the left were used for the salt bridge and the one
on the right were the test solution. If solutions on each side of the junction were
reversed, the signs of the junction potentials would be reversed.) It is apparent that
the liquid junction potential can be minimized by keeping a high concentration of a
salt such as KCl, the ions of which have nearly the same mobility, on one side of the
boundary. Ideally, the same high concentration of such a salt should be on both sides
of the junction. This is generally not possible for the test solution side of a salt bridge.
However, the solution in the other half-cell in which the other end of the salt bridge
forms a junction can often be made high in KCl to minimize that junction potential. As
noted before, this half-cell, which is connected via the salt bridge to form a complete

Christian7e c13.tex

13.6 REFERENCE ELECTRODES: THE SATURATED CALOMEL ELECTRODE

V2 - 08/13/2013 1:49 P.M. Page 407

407

Table 13.1

Some Liquid-Junction Potentials at 25◦ Ca
Boundary

Ej (mV)

0.1 M KCl||0.1 M NaCl
3.5 M KCl||0.1 M NaCl
3.5 M KCl||1 M NaCl
0.01 M KCl||0.01 M HCl
0.1 M KCl||0.1 M HCl
3.5 M KCl||0.1 M HCl
0.1 M KCl||0.1 M NaOH
3.5 M KCl||0.1 M NaOH
3.5 M KCl||1 M NaOH

+6.4
+0.2
+1.9
−26

−27
+3.1
+18.9
+2.1
+10.5

a Adapted from G. Milazzo, Electrochemie. Vienna: Springer, 1952; and D. A. MacInnes and Y. L. Yeh, J. Am.
Chem. Soc., 43 (1921) 2563.

cell, is the reference electrode. See the discussion of the saturated calomel electrode
below.
As the concentration of the (dissimilar) electrolyte on the other side of the
boundary (in the test solution) increases, or as the ions are made different, the
liquid-junction potential will get larger. Very rarely can the liquid-junction potential
be considered to be negligible. The liquid-junction potential with neutral salts is
less than when a strong acid or base is involved. The variation is due to the
unusually high mobilities of the hydrogen ion and the hydroxyl ion. Therefore, the
liquid-junction potential will vary with the pH of the solution, an important fact to
remember in potentiometric pH measurements. A potassium chloride salt bridge, at
or near saturation, is usually employed, except when these ions may interfere in a
determination. Ammonium chloride or potassium nitrate may be used if the potassium
or chloride ion interferes.
Various types of electrolyte junctions or salt bridges have been designed, such
as a ground-glass joint, a porous glass or ceramic plug, or a fine capillary tip. The
reference electrode solution then contains saturated KCl solution, which slowly leaks
through the bridge to create the liquid junction with the test solution.

Liquid-junction potentials are
highly pH dependent because of
the high mobilities of the proton

and hydroxide ions.

13.6 Reference Electrodes: The Saturated Calomel Electrode
A requirement of a reference electrode is that its potential be fixed and stable, unaffected
by the passage of small amounts of current required in making potentiometric
measurements (ideally, the current in the measurement is zero, but in practice some
small current must be passed—see below). Metal–metal salt electrodes generally
possess the needed properties.
A commonly used reference electrode is the saturated calomel electrode (SCE).
The term “saturated” refers to the concentration of potassium chloride; and at 25◦ C,
the potential of the SCE is 0.242 V versus NHE. An SCE consists of a small amount
of mercury mixed with some solid Hg2 Cl2 (calomel), solid KCl, and enough saturated
KCl solution to moisten the mixture. This is contacted with a saturated KCl solution
containing some solid KCl to maintain saturation. A platinum electrode is immersed
in the paste to make contact with the small mercury pool formed, and the connecting
wire from that goes to one terminal of the potential measuring device. A salt bridge
serves as the contact between the KCl solution and the test solution and is usually

Reference electrodes are usually
metal–metal salt types. The two
most common are the Hg/Hg2 Cl2
(calomel) and the Ag/AgCl
electrodes.


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 408

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

408

Pt wire sealed in the
inner tube to make
contact with the paste
Paste of Hg,
Hg2Cl2, and KCl

Hole for filling
with KCl solution

Pinhole for contact
of paste with the
KCI solution

To potentiometer

Fig. 13.3.

Saturated KCI solution

Commercial saturated
calomel electrode. (Source:
Courtesy of Arthur H. Thomas
Company.)

Porous ceramic
junction (salt bridge)

a fiber or porous glass frit wetted with the saturated KCl solution. If a different salt
bridge is needed to prevent contamination of the test solution (you can’t use the SCE
for chloride measurements!), then a double-junction reference electrode is used in
which the KCl junction contacts a different salt solution that in turn contacts the test
solution. This, of course, creates a second liquid-junction potential, but it is constant.
A commercial probe-type SCE is shown in Figure 13.3. This contains a porous
fiber or frit as the salt bridge in the tip that allows very slow leakage of the saturated
potassium chloride solution. It has a small mercury pool area and so the current it
can pass without its potential being affected is limited (as will be seen below, a small
current is usually drawn during potential measurements). The fiber salt bridge has a
resistance of ∼ 2500 satisfactory for use with any modern high input impedance
voltmeter, including a pH meter.

Example 13.3
Calculate the potential of the cell consisting of a silver electrode dipping in a silver
nitrate solution with aAg+ = 0.0100 M and an SCE reference electrode.
Solution

Neglecting the liquid-junction potential,
Ecell = Eind − Eref
0
Ecell = EAg
+ ,Ag − 0.05916 log

= 0.799 − 0.05916 log
= 0.439 V

1
aAg+

− ESCE

1
− 0.242
0.0100


Christian7e c13.tex

13.7 MEASUREMENT OF POTENTIAL

V2 - 08/13/2013 1:49 P.M. Page 409

409

Example 13.4
A cell voltage measured using an SCE reference electrode is −0.774 V. (The indicating electrode is the more negative half-cell.) What would the cell voltage be
with a silver/silver chloride reference electrode (1 M KCl; E = 0.228 V) or with
an NHE?
Solution

The potential of the Ag/AgCl electrode is more negative than that of the SCE by
0.242 − 0.228 = 0.014 V. Hence, the cell voltage using the former electrode is less
negative by this amount:
Evs. Ag/AgCl = Evs. SCE + 0.014
= −0.774 + 0.014 = −0.760 V
Similarly, the cell voltage using the NHE is 0.242 V less negative:
Evs. NHE = Evs. SCE + 0.242 V

Reference electrode potentials are

all relative. The measured cell
potential depends on which one is
used.

= −0.774 + 0.242 = −0.532 V
Potentials relative to different reference electrodes may be represented schematically
on a scale on which the different electrode potentials are placed (see Reference 2).
Figure 13.4 illustrates this for Example 13.4.
We should note that although calomel electrodes were once the gold standard,
many laboratories now wish to limit the use of toxic mercury, and therefore use
Ag/AgCl electrodes.

13.7 Measurement of Potential
We create a voltaic cell with the indicator and reference electrodes. We measure the
voltage of the cell, giving a reading of the indicator electrode potential relative to the
reference electrode. We can relate this to the analyte activity or concentration using
the Nernst equation.
THE pH METER
pH measurements with a glass (or other) electrode involve the measurement of
potentials (see Sections 13.11–13.16 below for detailed descriptions of how we
measure pH). The pH meter is essentially a voltmeter.
The pH meter is a voltage measuring device designed for use with high-resistance
glass electrodes and can be used to measure potential in both low- and high-resistance
circuits. pH meters are typically built with very high input impedance operational
amplifiers called electrometers as the front end.

Arnold Beckman (1900–2004)
developed the first commercial pH
meter to measure citrus acidity. It
was the first fully integrated

analytical instrument to combine
electronics with chemistry, and
Beckman Instruments was founded
in 1935 to produce it. (Courtesy of
Beckman Coulter, Inc. For a
fascinating history of his invention
and company formation, see
/>LXVII2/beckman.html)

−0.760 V(vs.Ag/AgCl)
+0.228 V (Ag/AgCl)
+0.242 V (SCE)

0 V (NHE)

−0.532 V(vs.NHE)
−0.774 V(vs.SCE)

Fig. 13.4.

Schematic
representation of electrode potential
relative to different reference
electrodes.


Christian7e c13.tex

410

A pH meter or electrometer draws
very small currents and are well
suited for irreversible reactions
that are slow to reestablish
equilibrium. They are also
required for high-resistance
electrodes, like glass pH or
ion-selective electrodes.
Impedance in an ac circuit is
comparable to resistance in dc
circuits, although aside from
resistance it involves
frequencydependent components.
While no ac measurements are
made by a pH meter, the term high
input impedance adjective for the
opeartional amplifiers simply
indicate that they are compatible
with ac measurements. Although
in principle no current is (or
should be) drawn by an ideal
voltmeter, in practice even the
high impedance voltmeters draw a
finite current, albeit very small, in
the 1-100 fA range.

V2 - 08/13/2013 1:49 P.M. Page 410

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

A pH meter is a high-input impedance voltmeter that senses the cell voltage,
provides a digital readout (either in terms of voltage or pH) and often provides an
amplified output to be acquired by an external data system device. Because very little
current is drawn, chemical equilibrium is not perceptibly disturbed. This is vital for
monitoring irreversible reactions that do not return to the prior state if an appreciable
amount of current is drawn. The resistance of a typical glass pH electrode is of the
order of 108 .
Sufficiently sensitive pH meters are available that will measure the potential with
a resolution of 0.1 mV. These are well suited for direct potentiometric measurements
with both pH electrodes and other ion-selective electrodes.
THE CELL FOR POTENTIAL MEASUREMENTS
In potentiometric measurements, a cell of the type shown in Figure 13.5 is set up.
For direct potentiometric measurements in which the activity of one ion is to be
calculated from the potential of the indicating electrode, the potential of the reference
electrode will have to be known or determined. The voltage of the cell is described by
Equation 13.7, and when a salt bridge is employed, the liquid-junction potential must
be included. Then,
Ecell = (Eind − Eref ) + Ej

(13.30)

The Ej can be combined with the other constants in Equation 13.30 into a single
constant, assuming that the liquid-junction potential does not differ significantly from
one solution to the next. We are forced to accept this assumption since Ej cannot
0
be evaluated under most circumstances. Eref , Ej , and Eind
are lumped together into a
constant k:
0
k = Eind

− Eref + Ej

(13.31)

Then (for a 1:1 reaction),
Ecell = k −

2.303RT
a
log red
nF
aox

(13.32)

To voltmeter

Reference electrode
Indicating electrode

Fig. 13.5.

Cell for potentiometric
measurements.


Christian7e c13.tex

13.9 RESIDUAL LIQUID-JUNCTION POTENTIAL—
—IT SHOULD BE MINIMIZED

V2 - 08/13/2013 1:49 P.M. Page 411

411

The constant k is determined by measuring the potential of a standard solution in
which the activities are known.

13.8 Determination of Concentrations
from Potential Measurements
Usually, we are interested in determining the concentration of a test substance rather
than its activity. Activity coefficients are not generally available, and it is inconvenient
to calculate activities of solutions used to standardize the electrode.
If the ionic strength of all solutions is held constant at the same value, the activity
coefficient of the test substance remains nearly constant for all concentrations of the
substance. We can then write for the log term in the Nernst equation:
−

2.303RT
2.303RT
2.303RT
log fi Ci = −
log fi −
log Ci
nF
nF
nF

(13.33)

Under the prescribed conditions, the first term on the right-hand side of this equation
is constant and can be included in k, (we will call this k ), so that at constant ionic
strength,
Ecell = k −

C
2.303RT
log red
nF
Cox

(13.34)

If the ionic strength is maintained
constant, activity coefficients are
constant and can be included in k
(to be called a new constant, k ).
So concentrations can be
determined from measured cell
potentials.

In other words, the electrode potential changes by ±2.303RT/nF volts for each 10-fold
change in concentration of the oxidized or reduced form.
It is best to construct a calibration curve of potential versus log concentration;
this should have a slope of ±2.303RT/nF. In this way, any deviation from this
theoretical response will be accounted for in the calibration curve. Note that the
intercept of the plot would represent the constant, k , which includes the standard
potential, reference electrode potential, liquid junction potential, and the activity
coefficient.
Since the ionic strength of an unknown solution is usually not known, a high

concentration of an electrolyte is added both to the standards and to the samples to
maintain about the same ionic strength. The standard solutions should have the same
matrix as the test solutions, notably any species that will change the activity of the
analyte, such as complexing agents. However, since the complete sample composition
is often not known, this is frequently not possible.

13.9 Residual Liquid-Junction Potential—It Should Be Minimized
We have assumed above in Equations 13.32 and 13.34 that k or k is the same in
measurements of both standards and samples. This is so only if the liquid-junction
potential at the reference electrode is the same in both solutions. But the test solution
will usually have a somewhat different composition from the standard solution, and
the magnitude of the liquid-junction potential will vary. The difference in the two
liquid-junction potentials is called the residual liquid-junction potential, and it will
remain unknown. The difference can be kept to a minimum by keeping the ionic
strength of both solutions as close as possible, and especially by keeping the pH of the
test solution and the pH of the standard solution as close as possible.

If the liquid-junction potentials of
the calibrating and test solutions
are identical, no error results (the
residual Ej = 0). Our goal is to
keep residual Ej as small as
possible.


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 412

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

412

13.10 Accuracy of Direct Potentiometric
Measurements—Voltage Error versus Activity Error
We can get an idea of the accuracy required in potentiometric measurements from
the percent error caused by a 1-mV error in the reading at 25◦ C. For an electrode
responsive to a monovalent ion such as silver,
Ecell = k − 0.05916 log

1
aAg+

(13.35)

and
aAg+ = antilog

For a dilute or poorly poised
solution, stirring the solution helps
achieve an equilibrium reading.
For a quantitative discussion of
poising capacity, see E. R.
Nightengale, “Poised
Oxidation-Reduction Systems. A
Quantitative Evaluation of Redox
Poising Capacity and Its Relation
to the feasibility of Redox
Titrations,” Anal. Chem., 30(2)
(1958) 267–272.

Ecell − k
0.05916

(13.36)

A ±1-mV error results in an error in aAg+ of ±4, or in pAg units, an error of ∼0.017.
The absolute accuracy of most electrode based measurements is no better than 0.2 mV;
this limits the maximum accuracy attainable in direct potentiometric measurements.
The same percent error in activity will result for all activities of silver ion with a 1-mV
error in the measurement. The error is doubled when n is doubled to 2. So, a 1-mV
error for a copper/copper(II) electrode would result in an 8% error in the activity of
copper(II). It is obvious, then, that the residual liquid junction potential can have an
appreciable effect on the accuracy.
The accuracy and precision of potentiometric measurements are also limited by
the poising capacity of the redox couple being measured. This is analogous to the
buffering capacity in pH measurements. If the solution is very dilute, the solution is
poorly poised and potential readings will be sluggish. That is, the solution has such
a low ion concentration that it takes longer for the solution around the electrode to
rearrange its ions and reach a steady state, when the equilibrium is disturbed during
the measurement process. This is why a high input impedance voltmeter that draws
very small current is preferred for potentiometric measurements in such solutions. To
maintain a constant ionic strength, a relatively high concentration of an inert salt (ionic
strength “buffer”) can be added; this also helps reduce solution resistance, helpful
when physically separate reference and indicator electrodes are used. Stirring helps
speed up the equilibrium response.
In very dilute solutions, the potential of the electrode may be governed by
other electrode reactions. In a very dilute silver solution, for example, − log(1/aAg+ )
becomes very negative and the potential of the electrode is very reducing. Under these
conditions, an oxidizing agent in solution (such as dissolved oxygen) may be reduced

at the electrode surface, setting up a second redox couple (O2 /OH− ); the potential
will be a mixed potential.
Usually, the lower limit of concentration that can be measured with a degree
of certainty is 10−5 to 10−6 M, although the actual range should be determined
experimentally. As the solution becomes more dilute, a longer time should be
allowed to establish the equilibrium potential reading because of slower approach to
equilibrium. An exception to this limit is in pH measurements in which the hydrogen
ion concentration of the solution is well poised, either by a buffer or by excess acid or
base. At pH 10, the hydrogen ion concentration is 10−10 M, and this can be measured
with a glass pH electrode (see Section 13.11). A neutral, unbuffered solution is poorly
poised, however, and pH readings are sluggish. Pure water is a speciailly difficult
sample to measure the pH of, both because of poor buffering and very high resistance.
Potassium chloride is often deliberately added prior to pH measurement. The best
choice is a refillable, liquid-filled electrode, ideally made of low-resistance glass.


Christian7e c13.tex

13.11 GLASS PH ELECTRODE—
—WORKHORSE OF CHEMISTS

V2 - 08/13/2013 1:49 P.M. Page 413

413

A flowing reference junction has a higher flow rate to minimize junction potentials.
A fast leak rate is desirable with pure water so that the equilibrium potential can be
established more quickly.

13.11 Glass pH Electrode—Workhorse of Chemists

The glass electrode, because of its convenience, is used almost universally for pH
measurements today. Its potential is essentially not affected by the presence of oxidizing
or reducing agents, and it is operative over a wide pH range. It is fast responding and
functions well in physiological systems. No other pH-measuring electrode possesses
all these properties.
PRINCIPLE OF THE GLASS ELECTRODE
A typical construction of a pH glass electrode is shown in Figure 13.6. For measurement, only the bulb need be submerged. There is an internal reference electrode and
electrolyte (Ag|AgCl|Cl− ) for making electrical contact with the glass membrane; its
potential is necessarily constant and is set by the concentration of HCl. A complete
cell, then, can be represented by
reference
electrode H+ (unknown)
(external)

reference
glass
H+ (internal) electrode
membrane
(internal)

The double line represents the salt bridge of the reference electrode. The glass electrode
is attached to the indicating electrode terminal of the pH meter while the external
reference electrode (e.g., SCE) is attached to the reference terminal.
The potential of the glass membrane is given by
Eglass = constant −

a +
2.303RT
log H int
F

aH+ unk

(13.37)

and the voltage of the cell is given by
Ecell = k +

2.303RT
log aH+ unk
F

(13.38)

Internal filling solution (HCl)

Glass membrane

Ag/AgCl reference electrode

Fig. 13.6.

Glass pH electrode.


Christian7e c13.tex

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

414

The glass pH electrode must be
calibrated using “standard
buffers.” See Section 13.12.

V2 - 08/13/2013 1:49 P.M. Page 414

where k is a constant that includes the potentials of the two reference electrodes, the
liquid-junction potential, a potential at the glass membrane due to H+ (internal), and a
term known as the asymmetry potential.
The asymmetry potential is a small potential across the membrane that is present
even when the solutions on both sides of the membrane are identical. It is associated
with factors such as nonuniform composition of the membrane, strains within the
membrane, mechanical and chemical attack of the external surface, and the degree of
hydration of the membrane. It slowly changes with time, especially if the membrane
is allowed to dry out, and is unknown. For this reason, a glass pH electrode must be
calibrated at least once a day. The asymmetry potential will vary from one electrode
to another, owing to differences in construction of the membrane.
Since pH = − log aH+ , Equation 13.38 can be rewritten1
Ecell = k −

2.303RT
pHunk
F

(13.39)

k − Ecell
2.303RT/F

(13.40)

or
pHunk =

It is apparent that the glass electrode will undergo a 2.303RT/F-volt response for each
change of 1 pH unit (10-fold change in aH+ ); k must be determined by calibration with
a standard buffer (see below) of known pH:
k = Ecell +

2.303RT
pHstd
F

(13.41)

Substitution of Equation 13.41 into Equation 13.39 yields
pHunk = pHstd +
We usually don’t resort to this
calculation in pH measurements.
Rather, the potential scale of the
pH meter is calibrated in pH units
(see Section 13.14).

Ecell std − Ecell unk
2.303RT/F

(13.42)

Note that since the determination involves potential measurements with a very highresistance glass membrane electrode (50 to 500 M ), it is essential to use a high input
impedance voltmeter.

Example 13.5
A glass electrode–SCE pair is calibrated at 25◦ C with a pH 4.01 standard buffer, the
measured voltage being 0.814 V. What voltage would be measured in a 1.00 × 10−3 M
acetic acid solution? Assume aH+ = [H+ ].
Solution

From Example 6.7 in Chapter 6, the pH of a 1.00 × 10−3 M acetic solution is 3.88;
0.814 − Ecell unk
∴ 3.88 = 4.01 +
0.0592
Ecell unk = 0.822 V
Note that the potential increases as the H+ (a cation) increases, as expected.
1 We will assume the proper definition of pH as − log a
H+ in this chapter since this is what the glass electrode

measures.


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 415

13.11 GLASS PH ELECTRODE—
—WORKHORSE OF CHEMISTS

415

To reference electrode terminal

To indicator electrode terminal

Solution level
Ag/AgCl
reference
electrode

KCl solution
Porous plug salt bridge
Glass bulb shield
Ag/AgCl internal
reference electrode
HCl solution

Fig. 13.7.

Combination
pH–reference electrode.

Glass electrode

COMBINATION pH ELECTRODES——A COMPLETE CELL
Both an indicating and a reference electrode (with salt bridge) are required to make a
complete cell so that potentiometric measurements can be made. It is convenient to
combine the two electrodes into a single probe, so that only small volumes are needed
for measurements. A typical construction of a combination pH–reference electrode is
shown in Figure 13.7. It consists of a tube within a tube, the inner one housing the pH
indicator electrode and the outer one housing the reference electrode (e.g., a Ag/AgCl
electrode) and its salt bridge. There is one lead from the combination electrode, but
it is split into two connectors at the end, one (the larger) going to the pH electrode

terminal and the other going to the reference electrode terminal. It is important that
the salt bridge be immersed in the test solution in order to complete the cell. The salt
bridge is often a small plug in the outer ring rather than a complete ring as illustrated
here. Combination electrodes are convenient, and therefore the most commonly used.
WHAT DETERMINES THE GLASS MEMBRANE POTENTIAL?
The pH glass electrode functions as a result of ion exchange on the surface of a hydrated
layer. The membrane of a pH glass electrode consists of chemically bonded Na2 O and
−SiO2 . The surface of a new glass electrode contains fixed silicate groups associated
with sodium ions, −SiO− Na+ . For the electrode to work properly, it must first be
soaked in water. During this process, the outer surface of the membrane becomes
hydrated. The inner surface is already hydrated. The glass membrane is usually 30 to
100 μm thick, and the hydrated layers are 10 to 100 nm thick.
When the outer layer becomes hydrated, the sodium ions are exchanged for
protons in the solution:
−SiO− Na+ + H+
solid

solution

−SiO− H+ + Na+
solid

solution

(13.43)

Other ions in the solution can exchange for the Na+ (or H+ ) ions, but the equilibrium
constant for the above exchange is very large because of the affinity of the glass for
protons. Thus, the surface of the glass is made up almost entirely of silicic acid, except

A combination electrode is a
complete cell when dipped in a
test solution.


Christian7e c13.tex

416

The pH of the test solution
determines the external boundary
potential.

Does the glass electrode sense H+
or OH− in alkaline solutions?

V2 - 08/13/2013 1:49 P.M. Page 416

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

in very alkaline solution, where the proton concentration is small. The −SiO− sites
are fixed, but the protons are free to move and exchange with other ions. (By varying
the glass composition, the exchange for other ions becomes more favorable, and this
forms the basis of electrodes selective for other ions—see below.)
The potential of the membrane consists of two components, the boundary
potential and the diffusion potential. The former is almost the sole hydrogen ion
activity-determining potential. The boundary potential resides at the surface of the
glass membrane, that is, at the interface between the hydrated gel layer and the external
solution. When the electrode is dipped in an aqueous solution, a boundary potential is
built up, which is determined by the activity of hydrogen ions in the external solution

and the activity of hydrogen ions on the surface of the gel. One explanation of the
potential is that the ions will tend to migrate in the direction of lesser activity, much
as at a liquid junction. The result is a microscopic layer of charge built up on the
surface of the membrane, which represents a potential. Hence, as the solution becomes
more acidic (the pH decreases), protons migrate to the surface of the gel, building
up a positive charge, and the potential of the electrode increases, as indicated by
Equations 13.37 and 13.38. The reverse is true as the solution becomes more alkaline.
The diffusion potential results from a tendency of the protons in the inner part
of the gel layer to diffuse toward the dry membrane, which contains −SiO− Na+ , and a
tendency of the sodium ions in the dry membrane to diffuse to the hydrated layer. The
ions migrate at a different rate, creating a type of liquid-junction potential. But a similar
phenomenon occurs on the other side of the membrane, only in the opposite direction.
These in effect cancel each other, and so the net diffusion potential is very small, and
the potential of the membrane is determined largely by the boundary potential. (Small
differences in diffusion potentials may occur due to differences in the glass across the
membrane—these represent a part of the asymmetry potential.)
Cremer described the first predecessor of the modern glass electrode [Z. Biol.
47 (1906) 56]. More than a hundred years later, exactly how a glass electrode works
is still not eminently clear. Pungor has presented evidence that the establishment
of an electrode potential is caused by charge separation, due to chemisorption of
the primary ion (H+ ) from the solution phase onto the electrode surface, that is,
a surface chemical reaction. Counter ions of the opposite charge accumulate in the
solution phase, and this charge separation represents a potential. A similar mechanism
applies to other ion-selective electrodes (below). [See E. Pungor, “The New Theory of
Ion-Selective Electrodes,” Sensors, 1 (2001) 1–12 (this is an open access electronic
journal: —Pungor author.)]
K. L. Cheng has proposed a theory of glass electrodes based on a capacitor model
in which the electrode senses the hydroxide ion in alkaline solution (where aH+ is very
small), rather than sensing protons. [K. L. Cheng, “Capacitor Theory for Nonfaradaic
Potentiometry,” Microchem. J., 42 (1990) 5.] Nonfaradaic here refers to reactions that

do not involve a redox process. Cheng has performed isotope experiments that suggest
the generally accepted ion exchange reaction between H+ and Na+ does not occur. He
argues that the electrode actually responds to OH− ions in alkaline solution (remember,
[H+ ] at pH 14 is only 10−14 M!) [C.-M. Huang et al., “Isotope Evidence Disproving
Ion Exchange Reaction Between H+ and Na+ in pH Glass Electrode,” J. Electrochem.
Soc., 142 (1995) L175]. While this theory is not generally accepted, Cheng et al. present
some compelling arguments and experimental results that make this an interesting
hypothesis. It has some commonality with Pungor’s double-layer hypothesis.
ALKALINE ERROR
Two types of error occur that result in non-Nernstian behavior (deviation from the
theoretical response). The first is called the alkaline error. Such error is due to the


Christian7e c13.tex

13.11 GLASS PH ELECTRODE—
—WORKHORSE OF CHEMISTS

417

capability of the membrane for responding to other cations besides the hydrogen
ion. As the hydrogen ion activity becomes very small, these other ions can compete
successfully in the potential-determining mechanism. Although the hydrated gel layer
prefers protons, sodium ions will exchange with the protons in the layer when the
hydrogen ion activity in the external solution is very low (reverse of Equation 13.43).
The potential then depends partially on the ratio of aNa+ external /aNa+ gel ; that is, the
electrode becomes a sodium ion electrode.
The error is negligible at pH less than about 9; but at pH values above this, the
+
H concentration is very small relative to that of other ions, and the electrode response

to the other ions such as Na+ , K+ , and so on, becomes appreciable. In effect, the
electrode appears to “see” more hydrogen ions than are present, and the pH reading
is too low. The magnitude of this negative error is illustrated in Figure 13.8. Sodium
ion causes the largest errors, which is unfortunate, because many samples, especially
alkaline ones, contain significant amounts of sodium. Commercial general-purpose
glass electrodes are usually supplied with a graphically represented alkaline error
correction values, and if the sodium ion concentration is known, these electrodes are
useful up to pH about 11.
By a change in the composition of the glass, the affinity of the glass for sodium
ion can be reduced. If the Na2 O in the glass membrane is largely replaced by Li2 O,
then the error due to sodium ions is markedly decreased. This is the so-called lithium
glass electrode, high-pH electrode, or full-range electrode (0 to 14 pH range). Most pH
electrodes in use today have glass membranes formulated to be capable of measurement
up to pH 13.5 with reasonable accuracy if sodium error is corrected for. But if you
need to make pH measurements in very alkaline solutions, the specially formulated
electrodes are recommended. As mentioned before, it was the discovery that variation
in the glass composition could change its affinity for different ions that led to the
development of glasses selective for ions other than protons, that is, of ion-selective
electrodes, that extended eventually to materials altogether different from glass.

Sample
pH reading
pH correction
to be added
8C 20
25
30
13

0.1 M

0.5 M

50

Fig. 13.8.

The sodium error of a
good general purpose pH electrode.
The example shows how to use this
“nomogram” to correct the apparent
measurement. Imagine that I have a
solution, 0.5 M in sodium and the
apparent pH read at 50◦ C is 12.10.
We draw a line from the pH 12.10
point on the x-axis through the
intersection point of the 50◦ C line and
the 0.5 M line and find the line
intersects the error axis at 0.01. The
actual pH is therefore
12.10 + 0.01 = 12.11. (Courtesy
Thermo Fisher Scientific Inc.)

0.005

sample
1M
sodium
concentration
2M

40

0.01
0.02

60
12

70

5M
208C
25

80

30

0.1 M
11

V2 - 08/13/2013 1:49 P.M. Page 417

40

0.5 M

50
1M
60

2M

70

10

80
5M

0.03
0.04
0.05
0.06
0.08
0.10
0.15
0.20
0.25
0.30
0.40
0.50
0.60
0.70

The glass electrode senses other
cations besides H+ . This becomes
appreciable only when aH+ is very
small, as in alkaline solution. We
can’t distinguish them from H+ ,
so the solution appears more

acidic than it actually is.


Christian7e c13.tex

418

V2 - 08/13/2013 1:49 P.M. Page 418

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

Error, pH units

0.4

Fig. 13.9.

Error of glass electrode
in hydrochloric acid solutions.
(From L. Meites and L. C. Thomas,
Advanced Analytical Chemistry.
Copyright © 1958, McGraw-Hill,
New York. Used with permission of
McGraw-Hill Company.)

0.2

0

−1.0

−0.5

0

+0.5

pH

ACID ERROR
At very low pH values (pH < 1), the gel layer of the pH-sensitive glass membrane
absorbs acid molecules. This absorption decreases the activity of hydrogen ions and
results in a lower potential at the outer membrane phase boundary. The pH measurement
therefore shows a higher pH value than the actual pH value of the sample solution. A
second and possibly greater contributor to the acid error is more aptly described as the
water activity error, is the second type causing non-Nernstian response. Such error
occurs because the potential of the membrane depends on the activity of the water
with which it is in contact. If the activity is unity, the response is Nernstian. In very
acidic solutions, the activity of water is less than unity (an appreciable amount is used
in solvating the protons), and a positive error in the pH reading results (Figure 13.9).
A similar type of error will result if the activity of the water is decreased by a high
concentration of dissolved salt or by addition of nonaqueous solvent such as ethanol.
In these cases, a large liquid-junction potential may also be introduced and another
error will thereby result, although this is not very large with small amounts of ethanol.
Similiar to specially formulated electrodes for use at strongly alkaline pH,
specialized electrodes are available for use in strongly acid solutions that exhibit
considerably less acid error. Acid error, in general, is smaller than alkaline error.

13.12 Standard Buffers—Reference for pH Measurements

Only the phosphate mixtures and
borax are really buffers. Disodium
tetraborate is effectively an
equimolar mixture of orthoboric
acid and its fully neutralized salt,
and so is a buffer. The pH values
change with temperature due to
the temperature dependence of the
Ka values.

Because we cannot measure the activity of a single ion (but only estimate it by
calculation using the Debye–H¨uckel equation), operational definitions of pH have
been proposed. One of these is that developed at the National Bureau of Standards
(NBS), now called the National Institute of Standards and Technology (NIST), under
the direction of Roger Bates. He developed a series of certified standard buffers for
use in calibrating pH measurements. The pH values of the buffers were determined
by measuring their pH using a hydrogen-indicating electrode in a cell without liquid
junction (similar to the cell given by Equation 13.22). A silver/silver chloride reference
electrode was used. From Equation 13.24, we see that the activity of the chloride ion
must be calculated (to calculate the potential of the reference electrode) using the
Debye–H¨uckel theory; this ultimately limits the accuracy of the pH of the buffers
to about ±0.01 pH unit. Faced with the problem of choosing a convention for the
ionic activity coefficient of a single species (chloride) to be used for the purpose of
assigning pH values, Bates chose values for chloride that were similar to those for the
mean activity coefficients (which can be measured) of HCl and NaCl in their mixtures.
Hence, this is the basis for the operational definition of pH. This convention is known


Christian7e c13.tex

13.12 STANDARD BUFFERS—
—REFERENCE FOR PH MEASUREMENTS

V2 - 08/13/2013 1:49 P.M. Page 419

419

Table 13.2

pH Values of NIST Buffer Solutionsa
Buffer
t (◦ C)
0
5
10
15
20
25
30
35
38
40
45
50
55
60
70
80
90
95

Tetroxalateb

Tartratec

Phthalated

Phosphatee

Phosphatef

Boraxg

Calcium
Hydroxideh

1.666
1.668
1.670
1.672
1.675
1.679
1.683
1.688
1.691
1.694
1.700
1.707
1.715
1.723

1.743
1.766
1.792
1.806

—
—
—
—
—
3.557
3.552
3.549
3.549
3.547
3.547
3.549
3.554
3.560
3.580
3.609
3.650
3.674

4.003
3.999
3.998
3.999
4.002
4.008

4.015
4.024
4.030
4.035
4.047
4.060
4.075
4.091
4.126
4.164
4.205
4.227

6.984
6.951
6.923
6.900
6.881
6.865
6.853
6.844
6.840
6.838
6.834
6.833
6.834
6.836
6.845
6.859
6.877

6.886

7.534
7.500
7.472
7.448
7.429
7.413
7.400
7.389
7.384
7.380
7.373
7.367
—
—
—
—
—
—

9.464
9.395
9.332
9.276
9.225
9.180
9.139
9.102
9.081

9.068
9.038
9.011
8.985
8.962
8.921
8.885
8.850
8.833

13.423
13.207
13.003
12.810
12.627
12.454
12.289
12.133
12.043
11.984
11.841
11.705
11.574
11.449
—
—
—
—

a From

R. G. Bates, J. Res. Natl. Bur. Std., A66 (1962) 179. (Reprinted by permission of the U.S. Government Printing Office.)
m potassium tetroxalate (m refers to molality, but only small errors result if molarity is used instead).
c Satd. (25◦ C) potassium hydrogen tartrate.
d 0.05 m potassium hydrogen phthalate.
e 0.025 m potassium dihydrogen phosphate, 0.025 m disodium monohydrogen phosphate.
f
0.008695 m potassium dihydrogen phosphate, 0.03043 m disodium hydrogen phosphate.
g
0.01 m borax.
h
Satd. (25◦ C) calcium hydroxide.
b 0.05

as the Bates–Guggenheim convention (Edward A. Guggenheim of Reading University
in the U. K. and Bates were charged by IUPAC to come up with a recommendation,
and Guggenheim went along with Bates’ suggestion). The partial pressure of hydrogen
is determined from the atmospheric pressure at the time of the measurement (minus
the vapor pressure of the water at the temperature of the solution).
The compositions and pH of NIST standard buffers are given in Table 13.2.
The NIST pH scale is a “multi-standard” scale with several fixed points. The British
Standards Institute, however, has developed an operational pH scale based on a single
primary standard, and pH of other “standard” buffers are measured relative to this,
these being secondary standards, rather than having a series of standard reference
solutions. In practice, when accurate pH measurement is necessary, the pH meter is
calibrated with two standard buffers that bracket the sample pH as closely as possible.
Although the absolute value of the pH accuracy is no better than 0.01 unit, the buffers
have been measured relative to one another to 0.001 pH. The potentials used in
calculating the pH can be measured reproducibly this closely, and the discrimination
of differences of thousandths of pH units is sometimes important (i.e., an electrode

may have to be calibrated to a thousandth of a pH unit). The pH of the buffers is
temperature dependent because of the dependence of the ionization constants of the
parent acids or bases on temperature.

Roger Bates (1912–2007) developed the NIST certified standard
buffers while at the old National
Bureau of Standards (NBS), now
NIST. From SEAC Communications. For a fascinating historical
account in Bates’s own words of the
development of the NIST
operational pH scale, see “Why
Students (and Others) Don’t Know
pH,” SEAC Communications 18 (3),
December 2002, written at the age
of 90. />SEACcom/SEACcom-dec02.pdf.


Christian7e c13.tex

420

V2 - 08/13/2013 1:49 P.M. Page 420

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

Note that several of these solutions are not really buffers, and they are actually
standard pH solutions whose pH is stable since we do not add acid or base. They are
resistant to pH change with minor dilutions (e.g., H+ ≈ Ka1 Ka2 ).
It should be pointed out that if a glass electrode–SCE cell is calibrated with
one standard buffer and is used to measure the pH of another, the new reading will

not correspond exactly to the standard value of the second because of the residual
liquid-junction potential.
The KH2 PO4 − Na2 HPO4 buffer (pH 7.384 at 38◦ C) is particularly suited for
calibration for blood pH measurements. Many blood pH measurements are made at
38◦ C, which is near body temperature; thus, the pH of the blood in the body is
indicated.
For a discussion of the above NIST pH standard and other proposed definitions
of pH, see the letters by W. F. Koch (Anal. Chem., December 1, 1997, 700A; Chem.
& Eng. News, October 20, 1997, 6).

13.13 Accuracy of pH Measurements
The residual liquid-junction
potential limits the accuracy of pH
measurement. Always calibrate at
a pH close to that of the test
solution.

Potentiometric measurements of
aH+ are only about 5% accurate.

The accuracy of pH measurements is governed by the accuracy to which the hydrogen
ion activity of the standard buffer is known. As mentioned above, this accuracy is not
better than ±0.01 pH unit because of several limitations. The first is in calculating the
activity coefficient of a single ion.
A second limitation in the accuracy is the residual liquid-junction potential. The
cell is standardized in one solution, and then the unknown pH is measured in a solution
of a different composition. We have mentioned that this residual liquid-junction
potential is minimized by keeping the pH and compositions of the solutions as near as
possible. Because of this, the cell should be standardized at a pH close to that of the
unknown. The error in standardizing at a pH far removed from that of the test solution

is generally within 0.01 to 0.02 pH unit but can be as large as 0.05 pH unit for very
alkaline solutions.
The residual liquid-junction potential, combined with the uncertainty in the
standard buffers, limits the absolute accuracy of measurement of pH of an unknown
solution to about ±0.02 pH unit. It may be possible, however, to discriminate between
the pH of two similar solutions with differences as small as ±0.004 or even ±0.002 pH
units, although their accuracy is no better than ±0.02 pH units. Such discrimination
is possible because the liquid-junction potentials of the two solutions will be virtually
identical in terms of true aH+ . For example, if the pH values of two blood solutions
are close, we can measure the difference between them accurately to ±0.004 pH. If
the pH difference is fairly large, however, then the residual liquid-junction potential
will increase and the difference cannot be measured as accurately. For discrimination
of 0.02 pH unit, changes in the ionic strength may not cause serious errors, but for
smaller pH changes than this, large changes in ionic strength will cause errors.
An error of ±0.02 pH unit corresponds to an error in aH+ of ±4.8% (±1.2 mV),2
and a discrimination of ±0.004 pH unit would correspond to a discrimination of
±1.0% in aH+ (±0.2 mV).
If pH measurements are made at a temperature other than that at which the
standardization is made, other factors being equal, the liquid-junction potential will
change with temperature. For example, in a rise from 25◦ to 38◦ C, a change of
+0.76 mV has been reported for blood. Thus, for very accurate work, the cell must be
standardized at the same temperature as the test solution.
2 The electrode response is 59 mV/pH at 25◦ C.


Christian7e c13.tex

13.14 USING THE PH METER—
—HOW DOES IT WORK?

V2 - 08/13/2013 1:49 P.M. Page 421

421

13.14 Using the pH Meter—How Does It Work?
We have already mentioned that owing to the high resistance of the glass electrode,
a high input impedance voltmeter (all pH meters qualify) must be used to make the
potential measurements. If voltage is measured directly, Equation 13.40 or 13.42 is
applied to calculate the pH. The value of 2.303RT/F at 298.16 K (25◦ C) is 0.05916;
if a different temperature is used, this value should be corrected in direct proportion to
the absolute temperature.
A digital pH meter is shown in Figure 13.10. The potential scale is calibrated
in pH units, with each pH unit equal to 59.16 mV at 25◦ C (Equation 13.39). The pH
meter is adjusted to indicate the pH of the standard buffer or the calibration function
will cause it to calibrate itself with the known pH of the calibrant buffer. Then, the
standard buffer is replaced by the unknown solution and the pH is read. This procedure,
in effect, sets the constant k in Equation 13.40 and adjusts for the asymmetry potential
as well as the other constants included in k.
Most pH meters contain a temperature adjustment dial, which changes the
sensitivity response (mV/pH) so that it will be equal to 2.303RT/F. For example, it is
54.1 mV at 0◦ C and 66.0 mV at 60◦ C.
Electrodes and meters are designed to have a point in calibrations lines, in the
midrange of activity measurements, where the potential essentially has no variation
with temperature. For pH glass electrodes, this is set at pH 7 (Figure 13.11). This is
called the isopotential point, and the potential is zero. (pH meters actually measure
potential which is converted to pH reading, and potentials can be recorded directly.)
Any potential reading different from 0 mV for a pH 7.0 standard buffer is called
the offset of that electrode. When the temperature is changed, the calibration slope
changes, and the intersection of the curves establishes the actual isopotential point.
If the isopotential point of the electrode differs from pH 7, then the temperature of

the calibration buffer and the test solution should be the same for highest accuracy
because a slight error will occur in the slope adjustment at different temperatures. For
more details of the isopotential point and its quantitative interpretation, see A. A. S.
C. Machado, Analyst, 19 (1994) 2263.

The temperature setting on the pH
meter adjusts T in the RT/nF
value, which determines the slope
of the potential versus pH buffers.

Fig. 13.10.

Typical pH meter.
(Courtesy of Denver Instrument
Company.)


Christian7e c13.tex

V2 - 08/13/2013 1:49 P.M. Page 422

CHAPTER 13 POTENTIOMETRIC ELECTRODES AND POTENTIOMETRY

422

500
400
60°C (66.10 mV/pH)
300
200

Potential, mV

25°C (59.16 mV/pH)
100

Isopotential point, pH 7

0°C (54.20 mV/pH)
0

0

2

4

6

8

10

12

−100
−200
−300
−400
−500
pH

Fig. 13.11.

Isopotential point.

In calibrating the pH meter, the electrodes are inserted in a pH 7.0 standard
buffer. The temperature of the buffer is checked and the temperature adjustment knob
is adjusted to that temperature. Using the standardized or calibration knob, the meter
is adjusted to read 7.00. Many pH meters have microprocessors that recognize specific
pH values, e.g., pH calibration standards of 4, 7, and 10. If you put an electrode in a
pH 7 calibration standard for example, and press the calibration button (or equivalent),
it will automatically calibrate the instrument to read the pH at 7.02, the pH of the
NIST standard NaH2 PO4 -Na2 H2 PO4 buffer. Next the slope is set by repeating the
calibration with either a pH 4 (potassium hydrogen phthalate pH 4.01) or 10 (Na2 CO3 NaHCO3 , pH 10.00) calibration standard, depending on the pH of the sample to be
measured. Most pH meters now have temperature measurement capability and include
a separate temperature probe so that temperature compensation can be automatic. The
temperature probe may be incorporated in the electrode itself.
Most pH meters are precise to ±0.01 pH unit (±0.6 mV) with a full-meter
scale of 14 pH units (about 840 mV). The meters can be set to read millivolts directly
(usually with a sensitivity of 1400 mV full scale). Higher-resolution pH meters are
capable of reading to ±0.001 pH unit; to accomplish this, the potential must be read
to closer than 0.1 mV.
When the pH of an unbuffered solution near neutrality is measured, readings will
be sluggish because the solution is poorly poised and a longer time will be required
to reach a stable reading. The solution should be stirred because a small amount
of the glass tends to dissolve, making the solution at the electrode surface alkaline
(Equation 13.43, where H2 O—the source of H+ —is replaced by NaOH solution).
See the discussion on measuring the pH of pure water at the end of Section 13.10.

13.15 pH Measurement of Blood—Temperature Is Important

The pH measurement of blood
samples must be made at body
temperature to be meaningful.

Recall from Chapter 7 that, because the equilibrium constants of the blood buffer
systems change with temperature, the pH of blood at the body temperature of
37◦ C is different than at room temperature. Hence, to obtain meaningful blood pH


Christian7e c13.tex

13.16 PH MEASUREMENTS IN NONAQUEOUS SOLVENTS

measurements that can be related to actual physiological conditions, the measurements
should be made at 37◦ C and the samples should not be exposed to the atmosphere.
(Also recall that the pH of a neutral aqueous solution at 37◦ C is 6.80, and so the acidity
scale is changed by 0.20 pH unit.)
Some useful rules in making blood pH measurements are as follows:

1. Calibrate the electrodes using a standard buffer at 37◦ C, making sure to select
the proper pH of the buffer at 37◦ C and to set the temperature on the pH meter
at 37◦ C (slope = 61.5 mV/pH). It is a good idea to use two standards for
calibration, narrowly bracketing the sample pH; this assures that the electrode
is functioning properly. Also, the electrodes must be equilibrated at 37◦ C
before calibration and measurement. The potential of the internal reference
electrode inside the glass electrode is temperature dependent, as may be
the potential-determining mechanism at the glass membrane interface; and
the potentials of the SCE reference electrode and the liquid junction are
temperature dependent. (We should note here that if pH or other potential
measurements are made at less than room temperature, the salt bridge or

the reference electrode should not contain saturated KCl, but somewhat less
concentrated KCl, because solid KCl crystals will precipitate in the bridge
and increase its resistance.)
2. Blood samples must be kept anaerobically to prevent loss or absorption
of CO2 . Make pH measurements within 15 min after sample collection,
if possible, or else keep the sample on ice and make the measurements
within 2 h. The sample is equilibrated to 37◦ C before measuring. (If a pCO2
measurement is to be performed also, do this within 30 min.)
3. To prevent coating of the electrode, flush the sample from the electrode
with saline solution after each measurement. A residual blood film can be
removed by dipping for only a few minutes in 0.1 M NaOH, followed by 0.1
M HCl and water or saline.

Generally, venous blood is taken for pH measurement, although arterial blood
may be required for special applications. The 95% confidence limit range (see Chapter
3) for arterial blood pH is 7.31 to 7.45 (mean 7.40) for all ages and sexes. A range
of 7.37 to 7.42 has been suggested for subjects at rest. Venous blood may differ from
arterial blood by up to 0.03 pH unit and may vary with the vein sampled. Intracellular
erythrocyte pH is about 0.15 to 0.23 unit lower than that of the plasma.

13.16 pH Measurements in Nonaqueous Solvents
Measurement of pH in a nonaqueous solvent when the electrode is standardized
with an aqueous solution has little significance in terms of possible hydrogen ion
activity because of the unknown liquid-junction potential, which can be rather large,
depending on the solvent. Measurements made in this way are usually referred to as
“apparent pH.” pH scales and standards for nonaqueous solvents have been suggested
using an approach similar to the one for aqueous solutions. These scales have no
rigorous relation to the aqueous pH scale, however. You are referred to the book
by Bates (Reference 3) for a discussion of this topic. Some efforts have since been
made to establish reference electrode potentials in mixed aqueous solvents at different

V2 - 08/13/2013 1:49 P.M. Page 423

423