Section 11

Potentiometry

Potentiometric Electrodes

• Potentiometric electrodes measure:

• Activity not concentration

• Concepts to review:

• Activity and affect factors

Potentiometric Electrodes• Electrode of the First Kind:• Metal in contact with Cation: M/Mn+

• Example: Ag/Ag+

• Ag+ + 1 e- Ag(s)

• Nernst Equation:• E = Eo

Ag+/Ag – 2.303RT/nF(log1/aAg+)• Where aAg+ = activity of silver ion• Emeasured = Ecell = Eind vs NHE = Eind – ENHE

• Why does Ecell = Eind ?• Eref solution Eind

Potentiometric Electrodes

• Electrode of the Second Kind:

• General Form: MMXXn-

• Example: AgAgCl(S)Cl-

• AgCl + 1 e- Ag(s) + Cl-

• Nernst Equation:

• E = EoAgCl/Ag – 2.303RT/nF(log aCl-)

• What observation can you make about potential for this electrode?

Potentiometric Electrodes

• Can the previous electrode be utilized to measure aAg+ ?

Potentiometric Electrodes • Electrode of the Second Kind:• General Form: MMXXn-

• Example: Calomel Electrode• Hg Hg2Cl2(S)Cl-

• Hg22+ + 2 e- Hg(s)

• E = EoHg22+/Hg – 2.303RT/nF(log 1/aHg22+

• Hg2Cl2(s) + 2e- 2 Hg + 2 Cl-

• E = EoHg2Cl2/Hg – 2.303RT/nF(log aCl-)2

• Kosp

= aHg ‘ (aCl-)2

Potentiometric Electrodes • REDOX ELECTRODES:

• Inert Metal Electrode: Platinum, Pt

• Ma+ + n e- M(a-n)+

• E = EoMa+/M(a-n)+ - 2.303RT/nF (log aM(a-n)+/ aMa+

• Example:

• MnO-4 + 8 H+ + 5e- Mn2+ + 4 H2O

• Can you write the potential for this half cell?

Potentiometric Electrodes• Normal Hydrogen Electrode:

• H+ + e- ½ H2(g)

• Can you write the Nernst Equation for this half cell?

Fig. 13.1. Hydrogen electrode.

H+ + e- = ½H2

E = Eo – 2.303RT/F log (PH2)1/2/aH+

Eo = 0.000 V. If PH2 = 1 atm., E = -2.303RT/F pH

H+ + e- = ½H2

E = Eo – 2.303RT/F log (PH2)1/2/aH+

Eo = 0.000 V. If PH2 = 1 atm., E = -2.303RT/F pH

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Potentiometric Electrodes

• Cells Without Liquid Junction:

• Two Half Cells Required

• Indicator Electrode

• Reference Electrode

• Example: Ptg), HCl (solution)gCl(s)g

• Can you write the cell potential Ecell for this electrode system?

Potentiometric Electrodes

• Cells With Liquid Junction:

• Two Half Cells Required

• Indicator Electrode

• Reference Electrode

• Example:

• HgHg2Cl2(sCl(saturated)Cl(solution), H2(g)Pt

• Can you write the cell potential Ecell for this electrode system?

• Ecell = (Eright – Eleft) + Ej

Fig. 13.2. Representation of liquid-junction potential.

Ej is due to unequal migration of cations and anions at the boundary of a liquid junction,

e.g., at a salt bridge interface.

This charge separation results in a potential at steady state.

Ej is minimized by a high concentration of a salt with nearly equal mobilities of cation and anion in the salt bridge, e.g., saturated KCl.

The flux of the migration of this electrolyte is much greater than more dilute ones, and largely determines Ej.

Ej is due to unequal migration of cations and anions at the boundary of a liquid junction,

e.g., at a salt bridge interface.

This charge separation results in a potential at steady state.

Ej is minimized by a high concentration of a salt with nearly equal mobilities of cation and anion in the salt bridge, e.g., saturated KCl.

The flux of the migration of this electrolyte is much greater than more dilute ones, and largely determines Ej.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

H+ and OH- have high mobilities and so pH markedly affects Ej.

Therefore, pH electrode calibration should be done near the sample pH.

H+ and OH- have high mobilities and so pH markedly affects Ej.

Therefore, pH electrode calibration should be done near the sample pH.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.3. Commercial saturated calomel electrode.

The S.C.E. is a common reference electrode.

The cell half-reaction is: ½Hg2Cl2 + e- = Hg + Cl-.

E = Eo – 0.0592 log 1/aCl- = 0.242 V for saturated KCl.

The S.C.E. is a common reference electrode.

The cell half-reaction is: ½Hg2Cl2 + e- = Hg + Cl-.

E = Eo – 0.0592 log 1/aCl- = 0.242 V for saturated KCl.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.4. Schematic representation of electrode potentialrelative to different reference electrodes.

Reference electrode potentials are all relative.

The measured cell potential depends on the one used.

Reference electrode potentials are all relative.

The measured cell potential depends on the one used.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.5. Cell for potentiometric measurements.

A complete cell consists of an indicating electrode that responds to the analyte and a reference electrode of fixed potential.

The potential difference between the two is measured.

A complete cell consists of an indicating electrode that responds to the analyte and a reference electrode of fixed potential.

The potential difference between the two is measured.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

 Fig. 13.6. Glass pH electrode.

Makes electrical contact with glass,

and sets potential of ref. electrode.

Eglass = constant – 2.303RT/F log (aH+ int/ aH+ ext)

The hydrated glass responds to aH+.

The asymmetry potential of the glass membrane is unknown, so the electrode must be calibrated with a standard buffer.

Eglass = constant – 2.303RT/F log (aH+ int/ aH+ ext)

The hydrated glass responds to aH+.

The asymmetry potential of the glass membrane is unknown, so the electrode must be calibrated with a standard buffer.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.7. Combination pH-reference electrode.

This is two electrodes in one.

The porous plug salt bridge must be immersed in the solution.

This is two electrodes in one.

The porous plug salt bridge must be immersed in the solution.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.8. Error of Corning 015 glass electrode in strongly alkaline solutions containing various cations.

 When aH+ is very small, the glass electrode senses other cations.

The solution appears more acidic than it really is.

 When aH+ is very small, the glass electrode senses other cations.

The solution appears more acidic than it really is.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.9. Error of glass electrode in hydrochloric acid solutions.

In very acid solutions, the activity of water is less than unity (it solvates the proton).

The aH+ is decreased, and the pH reading is increased.

High concentrations of dissolved salts or adding a nonaqueous solvent does the same.

In very acid solutions, the activity of water is less than unity (it solvates the proton).

The aH+ is decreased, and the pH reading is increased.

High concentrations of dissolved salts or adding a nonaqueous solvent does the same.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

These are really accurate to 0.01 pH, due to Debye-Huckel limitations in calculating activity of the chloride ion in the Ag/AgCl reference electrode used in measurements.

But they are reproducible to 0.001 pH unit.

Only the phosphate mixtures are really buffers.

Ka values change with temperature, and so the pH changes.

These are really accurate to 0.01 pH, due to Debye-Huckel limitations in calculating activity of the chloride ion in the Ag/AgCl reference electrode used in measurements.

But they are reproducible to 0.001 pH unit.

Only the phosphate mixtures are really buffers.

Ka values change with temperature, and so the pH changes.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.10. Typical pH meter.

The potential scale is calibrated in pH units (59.16 mV/pH at 25o C).

A temperature adjustment feature changes the slope by 2.303RT/F.

The potential scale is calibrated in pH units (59.16 mV/pH at 25o C).

A temperature adjustment feature changes the slope by 2.303RT/F.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.11. Isopotential point.

The isopotential point (usually at pH 7) of a glass electrode is temperature insensitive.

The potential is set at zero here (if different than zero, this is the offset).

The isopotential point (usually at pH 7) of a glass electrode is temperature insensitive.

The potential is set at zero here (if different than zero, this is the offset).

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.12. Crystal membrane electrode.

The most common ion-selective electrode of this type is the fluoride electrode: LaF3 crystal doped with Eu(II) to increase conductivity.

Plastic membrane-ionophore electrodes have a similar design, with a soft plastic PVC membrane containing a neutral lipophilic ionophore that selectively complexes with the test ion.

The most common ion-selective electrode of this type is the fluoride electrode: LaF3 crystal doped with Eu(II) to increase conductivity.

Plastic membrane-ionophore electrodes have a similar design, with a soft plastic PVC membrane containing a neutral lipophilic ionophore that selectively complexes with the test ion.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.13. Liquid-membrane electrode.

The potential determining “membrane” is a layer of water-immiscible liquid ion exchanger at the porous membrane face.

The potential determining “membrane” is a layer of water-immiscible liquid ion exchanger at the porous membrane face.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

There are numerous plastic membrane-ionophore electrodes also, e.g., for Na+, K+, Li+, and Ca2+.

There are numerous plastic membrane-ionophore electrodes also, e.g., for Na+, K+, Li+, and Ca2+.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.14. 14-Crown-4 ether that selectively binds lithium ion.

The crown ether cavity size is just right for complexing lithium ion.

It is placed in a PVC plastic membrane.

The crown ether cavity size is just right for complexing lithium ion.

It is placed in a PVC plastic membrane.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

 Fig. 13.15. Ionophores for H+, Na+, and Ca2+.

 Amide-based ionophores in PVC membranes are good complexers of these ions.  Amide-based ionophores in PVC membranes are good complexers of these ions.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

The separate solution method of determining selectivity coefficient for ion A with respect to ion B.

The larger the potential difference, the more selective the electrode. The larger the potential difference, the more selective the electrode.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.16. Fixed interference calibration curve.

The calibration curve for A is prepared in the presence of fixed activity of interfering ion.

zA and zB are the charges of the ions.

The intersection is where the electrode responds equally to both ions.

The calibration curve for A is prepared in the presence of fixed activity of interfering ion.

zA and zB are the charges of the ions.

The intersection is where the electrode responds equally to both ions.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

Fig. 13.17. Solid-state ISFET electrode.

These are ion-selective field-effect transistor electrodes.

A semiconductor transistor serves as the base for electrical contact.

It is coated with insulating layers of SiO2 and Si3N4, and then an ion-sensitive membrane.

This one is a non-glass ISFET pH electrode.

These are ion-selective field-effect transistor electrodes.

A semiconductor transistor serves as the base for electrical contact.

It is coated with insulating layers of SiO2 and Si3N4, and then an ion-sensitive membrane.

This one is a non-glass ISFET pH electrode.

©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)