Chapter 15 Electrode Measurements
Problems: 2, 4, 8, 11, 18
In the last chapter we learned how to use oxidation-reduction reaction to make
cells or batteries. In this chapter we will learn how to harness the same math and
reactions to make indicator electrodes
An electrode is a cell designed specifically to measure the solution concentration
of a chemical. The whole area of using cell voltages to gain chemical information is
called potentiometry.
Electrodes and potentiometry are an important aspect and everyday application
of chemistry. Table 15-1 from your text lists several chemical species that must be
determined in a normal blood work. Most, if not all of these compounds are determined
using electrodes.
In the last chapter we found that there are two types of indicator electrodes,
inert, where the metal of the electrode does not take part in any chemical reactions,
active, where the metal of the electrode takes part in a redox reaction that helps to
determine the electrode’s potential. This chapter focuses on these types of electrodes
14-1 The Silver Indicator electrode
In the last chapter we used we used a silver/AgCl electrode as a reference electrode.
Now lets learn how to use a silver electrode as an indicator electrode. In its simplest
form all we have to do is to dip a silver wire into a solution, and this becomes a silver
indicator electrode. Lets see how this works.
Ag+ + e- WAg(s)
E0 = .799V
we need to measure this potential against something else so lets use a calomel
electrode, and get a cell that resembles the figure 15-1 on page 312
The E of the calomel electrode is fixed so it gives EL = .241 V
ER = Esilver ½ reactioon = .799- .059log(1/[Ag+])
Ecell = ER-EL = .799- .059log(1/[Ag+])- .241 V
=.558 -.059 log(1/[Ag+])
= .588 + .059 log [Ag+]
you can see that the V of this system will change by .059V each time the
concentration changes by a power of 10
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Now let’s look at the electrode you will be making and using in the lab. Figure
from lab manual. How are these electrodes similar or different?
Similar, have silver wire in solution as an indicator
Different, book example electrode has El as a calomel electrode as the reference
(V=.241), lab electrode has some copper wire in a solution of 0.1M CuSO4
What is going on? Copper wire in CuSO4 solution is a simple reference
electrode, much simpler and chemically safer than getting out mercury and making a
calomel electrode
What is the Voltage of this reference?
Reaction:
Cu2++2e-6Cu0
E0 = .339
Nernst Equation:
E = .339 + -.059/2 log (1/0.1)
= .339-.0295
= .3095
Since calculating the voltage of this set of electrodes at given silver
concentrations is part of you lab exercise, I’ll leave it to you to put the pieces together
from here.
Titration of a Halide Ion with Ag+
Now let’s link this electrode with a precipitation titration we are doing in the lab
Say we were using AgNO3 to titrate the I- in a solution
Ag+ + I- WAgI(s)
Ksp = 8.3x10-17
Ag+ + Cl- WAgCl(s)
Ksp = 1.0x10-10
Rather than going through all the calculations, lets just cut to the chase. Two
different species will precipitate out in this titration, AgI, and AgCl. Which one
precipitates first? (AgI with the smaller Ksp). Before the first equivalence point the Ag
and the I combine and ppt out, so you calculate [Ag+] using the concentration if I- and
the Ksp of AgI. After the first equivalence point Ag+ is now being precipitated by the
Cl- so the [Ag+] is calculated from the Cl- concentration and the Ksp of AgCl. Finally,
after the second equivalence point you have excess Ag+ in the solution and the [Ag+] is
calculated from this excess and the total volume. You then plug the [Ag+] into the E
equation of the silver/copper reference cell we were just working with and come up with
a voltage. The result should look like figure 6-5 from your text, although the actual
voltage may go up instead of down, depending on how you hooked up your volt meter
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between the electrodes.
In the above system, the potential we measured comes directly from a redox
reaction taking place on the metal of the electrode. Now let’s explore another
way we can get a potential in a solution that does NOT involve redox chemistry.
15-2 Junction Potentials
Whenever you have two different solution in contact, a junction potential will
develop between the liquids because the ions making the solution can’t diffuse at the
same speed. Examine the diagram below:
If you go back to chapter table 12-1
(Chapter 12) you will find that the ionic radius of
Cl- is 300 pM, while that of Na+ is 450 pM. The
slightly large size means that Na will travel
through the solution a little more slowly. (The
mobilities are given in table 15-2) What
happens as the Cl- ions pass the Na+ ions?
You get a charge imbalance, and anytime you
have a charge imbalance, you have a potential.
In this case the Junction Potential because it
occurs at the junction of two solutions
Implication of Junction potentials
I. You may have wondered why I always used KCl in my salt bridges. If you check you
will find that K+ and Cl- have similar mobilities, so salt bridges made with KCl have
negligible junction potentials.
II. All electrodes contain solvent /solvent junctions, and all have these built in junction
potentials that will make a measured voltage different than its theoretical value. As a
result we ALMOST NEVER determine a concentration directly from a voltage. We
either use an electrode in a titration where the absolute value of the potential is not as
important as the big change that occurs in the potential at the equivalence point, or we
first calibrate an electrode using standard of known concentrations
15-3 How Ion-Selective Electrodes work.
First, what and why is an ion-selective electrode.
An Ion-selective electrode is one that is sensitive only to a particular ion. In the
hospital setting you the most likely place you will see ion-selective electrode is in
determination of blood concentration of different ions. For instance you can make
electrode to measure blood pH, O2, CO2, glucose, etc.
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The ion selective electrodes we will discuss here all work on the same principle.
We will have some king of barrier between two solutions. If there is a difference in
concentration of an ion across the two sides of the barrier we will get a potential,
analogous to the junction potential we just discussed. The only trick will be go make
our concentration dependent electrode sensitive to just one ion in the solution.
Let’s say we have two solutions, one that Is 1M HCl (On the left) and one that is
water (on the right). Let’s separate these two solution by an impermeable membrane,
is nothing can cross. Because nothing can pass between the cell there is a 0 potential
difference.
No let’s manipulate the membrane so only H+ can pass. What happens? The
1M H+ from the left side wants to move to the 1x10-7 M H- concentration on the right
side. Since the + ions are moving and the negative ions are staying behind, you
quickly build up a + charge and a junction potential. If you stick electrodes into these
solution you can now see this potential and, by the Nernst Equation E= .059/n log
1M/1x10-7 = 6x.059 = .354 volts
(Another way to look at it. Initially both solution are electrically neutral. If you
allow H+ to move, then you have a charge imbalance on the solution on both sides of
the membrane, and the junction potential represents this imbalance)
The whole trick of ion selective electrodes is to use clever chemistry to
manipulate the membrane so that only 1 kind of ion can pass through it.
Thus we now have two types of indicator electrodes
Metal electrodes : where the metal of the electrode is involved in redox reaction
that creates a potential
Ion-selective electrodes: No redox chemistry is involved, the potential comes
from creating a barrier between two solutions that allows a single type of ion to pass (a
semi-permeable barrier)
15-4 pH measurement with a glass electrode
Perhaps the most familiar example of an ion selective electrode is the pH
electrode. Let’s see how this works
Well just as I have outlined above. The tricky part is the membrane. In reality it
is a glass bulb about 0.1 mm thick (This is the part that I warn you about breaking in the
lab)
In reality the membrane does not actually let H+ ions across the glass, but what
seem to happen is that the glass is a polymeric mass of Si and O (See figure 14-9) in
which + sodium atoms are caught. As you have H+ concentration differences, the Na+
within the electrode glass moves in response and has creates the same junction
potential you would have if H+ has moved.
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The actual construction of a pH electrode is actually fairly complicated because
it contains not only the special glass electrode, but it has two reference electrodes built
into it, one for each side of the glass membrane (See overhead)
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The line diagram of this electrode is:
Ag(s)| AgCl(s)|Cl-(ag)|| H+(aq, outside) :Special glass:
H+(aq,inside),Cl-(aq,inside)|AgCl(s)|Ag(s)
IN theory the response of an electrode should be
E= constant + $(.059) log (H+outside)/(H+inside)
The actual response of each electrode is different anc varies from day to day. Thus
they have to be calibrated at least daily. This is done by using buffer solutions of two
known pH/s and tuning the electronics of the pH meter to match those values.
Errors in pH Measurements
pH measurement can have errors here are some of the common ones
1. The pH standards are only good to ± .01 pH units, so your absolute pH
measurement is only good to that level as well
2. There is a salt bridge in the pH electrode, that will have a junction potential so
this add an uncertainty of ± .01
3. The electrode glass works by passing Na+ Thus if [Na+] is high and [H+] is low
errors will occur because the electrode can sense Na+
as well.
4. In strong acids the measured pH is high, but we don’t know why
5. You need to let the electrode equilibrate. This can take a few second in
buffered solutions or it can take minutes in unbuffered solutions (ie the end point of a
titration)
6. Dry electrodes must be soaked for several hours before they are ready to use.
7. Measurements and calibration should take place at the same temperature.
Errors 1 and 2 make the total absolute error in a pH measurement ± sqrt(.012 + .012)
.014 pH units. Relative accuracy between two measurements (say in a titration curve
is about .002 pH units
Solid State pH sensors Figure 15-11
New - High tech
Use a field effect transistor
transistor that binds H+ and that varies current flow through transistor
look at pictures in text
15-5 Ion-Selective Electrodes
There are 4 major types of ion selective electrodes
1. Glass membrane
2. Solid-state
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3. Liquid-based electrodes
4. Combination electrodes
1. Glass membrane we have already talked about in reference to measuring pH.
By manipulating the formula of the glass you can make electrodes for Li+, Na+, K, and
Ag+. None of the electrodes is perfectly selective, and the amount that an electrode
responds to another ion can be quantitated in the selectivity coefficient, but we won’t
get into that. Other tah to say that your standard pH electrode tends to give bad results
when [H+] <10-12 or [Na+] >.01
2. Solid-State electrodes (fig 15-12)
In a solid state electrode an inorganic crystal is used at the interface between
the reference and test solutions. In this case the ion to be sensed is part of the crystal,
and actually migrates across the crystal during the measurement.
Electrodes for
F-, Cl-, Br-, I-, SCN- ,CN and S-2 have been made
in this manner
3.Liquid-based Ion-Selective Electrodes (fig 15-14)
Here the barrier between liquids is a hydrophobic membrane that is saturated
with a metal ion chelator. Because the membrane is hydrophobic, no ions can pass
though it. However metals that bind to the chelator can enter the membrane and
establish a junction potential just as in out other electrodes
Electrode for Ca+2, NO3-, ClO4- BF4- and Cl- have been made in this way.
4. Compound electrodes
Diagram Figure 15-17
In compound electrodes you have one of the above electrodes surrounded by a
membrane that will let a different analyte ion through, and then somehow covert it into
an ion that the electrode senses
The simplest example of this is a pH electrode surrounded by a membrane that
only allows CO2 to pass though it. As CO2 passes the outer membrane, you have the
reaction CO2 + H2O W H2CO3 an acid, and this lowers the pH of the solution which you
then detect with the inner electrode. Electrodes of this kind are used for the detection
of acidic and basic gasses both in the gas phase and in solution
Another trick that can be done is to make electrodes for various enzymatic
reactions that generate H+. There are lots of biochemical reaction that either generate
or utilize H+. You can sense for only one of these compounds by surrounding a pH
type electrode with a semipermeable membrane the enzyme that runs the reaction in
the liquid around the membrane. As one of the compounds that run the reaction enters
the space, the enzyme run the reaction, and H+ is used or generated, which the pH
electrode then senses