Potentiometry

Dr Hisham E Abdellatef 2010

• It is a method of analysis in which we determine the concentration of an ion or substance by measuring the potential developed when a sensitive electrode is immersed in the solution of the species to be determined.

Mo = Mn++ neApplying Nernest equation.

Determination of the substances by potentiometric technique can be carried out by two ways:

1.Direct potentiometry and2.potentiometric titrations

• The potential of the indicator electrode cannot be measured alone;

• For any potentiometric measurement

we must have: 1.Reference electrode2.Indicator electrode.3.Potentiometer4.Salt bridge to connect the two

electrode solutions and complete the circuit.

A- Reference electrodeReference electrode must:1. Have a constant potential 2. Its potential must be definite

To express any electrode we have to mention:1. Redox reaction at the electrode surface.2. Half cell and Nernst equation.3. Sketch of its design.4. Any necessary conditions for its preparation.5. Any necessary precautions for its use.

It’s a primary reference electrode. Its potential is considered to be zero.

Electrode reaction: half cell: pt/ H2 , H+ (1N) Eo = zero d-Limitation1. It is difficult to be used and

to keep H2 gas at one atmosphere during all determinations.

2. It needs periodical replating of Pt. Sheet with Pt. Black

Standard Hydrogen Electrode

Saturated calomel electrode (S.C.E.)

Hg | Hg2Cl2 (sat’d), KCl (sat’d) | |

electrode reaction in calomel hal-cell

Hg2Cl2 + 2e = 2Hg + 2Cl–

Eo = + 0.268V

E = Eo – (0.05916/2) log[Cl–]2 = 0.244 V

Temperature dependent

Potential of the electrode depends on the chloride ion

Hg2Cl2 ⇌ 2 Hg22+ + 2Cl-

Sp Hg2Cl2 = [Hg22+]2 [Cl-]2

Ksp = 1.8 ×10–18

E = Eo – (0.0591/2) log[Cl–]2 = 0.244 V

Hg2Cl2 Hg22+

+ 2Cl– Ksp = 1.8 ×10–18

Saturated KCl = 4.6 M KCl

The crystal structure of calomel(Hg2Cl2), which has limited solubility in water (Ksp = 1.8 ×10–18).

KCl E volt

Saturated 0.241

1M 0.280

0.1 M 0.334

Silver-silver chloride electrode

Ag(s) | AgCl (sat’d), KCl (xM) | |

AgCl(s) + e = Ag(s) + Cl–

Eo = +0.244V

E = Eo – (0.05916/1) log [Cl–]

E (saturated KCl) = + 0.199V (25oC)

• SpAgCl = [Ag+] [Cl-]

E = Eo – (0.05916/1) log [Cl–]

Disadvantage of silver-silver chloride electrode • It is more difficult to prepare than SCE.• AgCI in the electrode has large solubility in

saturated KCl

Advantage of Ag/AgCI electrodes over SCE. • It has better thermal stability.• Less toxicity and environmental problems with

consequent cleanup and disposal difficulties.

B- Indicator electrode

its potential is sensitive to the concentration of analyte

Ecell=Eindicator-EreferenceIt must be: (a) give a rapid response and(b) its response must be reproducible. Metallic electrodes: where the redox reaction takes

place at the electrode surface.Membrane (specific or ion selective) electrodes:

where charge exchange takes place at a specific surfaces and as a result a potential is developed.

1. Electrodes for precipitemetry and complexometry

a- First-order electrodes for cations:e.g. in determination of Ag+ a rode or wire of

silver metal is the indicator electrode, it is potential is:

It is used for determination of Ag+ with Cl-, Br-

and CN-. Copper, lead, cadmium, and mercury

Example of First-order electrode

Ag+ + e = Ag(s) Eo = + 0.800V

E = 0.800 – (0.05916/1) log {1/[Ag+]}

b) Second order electrodes for anions

A metal electrode can sometimes be indirectly responsive to the concentration of an anion that forms a precipitate or complex ion with cations of the metal.

Ex. 1. Silver electrode

The potential of a silver electrode will accurately reflect the concentration of

iodide ion in a solution that is saturated with silver iodide.

AgI(s) + e = Ag(s) + I– Eo = – 0.151V

E = – 0.151 – (0.05916/1) log [I–] = – 0.151 + (0.05916/1)pI

2. Mercury electrode for measuring the concentration of the EDTA anion Y4–.

Mercury electrode responds in the presence of a small concentration of the

stable EDTA complex of mercury(II).

HgY2– + 2e = Hg(l) + Y4– Eo = 0.21V E = 0.21 – (0.05916/2) log ([Y4–] /[HgY2–])

K = 0.21 – (0.05916/2) log (1 /[HgY2–])E = K – (0.05916/2) log [Y4–] = K +(0.05916 / 2) pY

Examples:

Ag(s) | AgCl[sat’d], KCl[xM] | | Fe2+,Fe3+) | Pt

Fe3++e = Fe2+ Eo = +0.770V

Ecell = Eindicator – Ereference

= {0.770 – (0.05916/1) log [Fe2+]/[Fe3+]} – {0.222 – (0.05916/1) log [Cl–]}

2. Inert electrodes (Indicators electrodes for redox reaction) Chemically inert conductors such as gold, platinum, or carbon that do not participate, directly, in the redox process are called inert electrodes. The potential developed at an inert electrode depends on the nature and concentration of the various redox reagents in the solution.

2) Membrane indicator electrodes

The potential developed at this type of electrode results from an unequal charge buildup at opposing surface of a special membrane. The charge at each surface is governed by the position of an equilibrium involving analyte ions, which, in turn, depends on the concentration of those ions in the solution.

The electrodes are categorized according to the type of membrane they employ :

glass,

polymer,

crystalline,

gas sensor. The first practical glass electrode. (Haber and Klemensiewcz, Z. Phys. Chem, 1909, 65, 385.

3. indicator electrodes for neutralization reaction

Glass Membrane Electrode

Composition of glass membranes

70% SiO2

30% CaO, BaO, Li2O, Na2O,

and/or Al2O3

Ion exchange process at glass membrane-solution interface:

Gl– + H+ = H+Gl–

(a) Cross-sectional view of a silicate glass struture. In addition to the three Si│O bonds shown, each silicon is bonded to an additional oxygen atom, either above or below the plane of the paper. (b) Model showing three-dimensional structure of amorphous silica with Na+ ion (large dark blue) and several H+ ions small dark blue incorporated.

Glass Membrane Electrode

E = K + 0.059 (pH1 - pH2)K= constant known by the

asymmetry potential.PH1 = pH of the internal solution 1.PH2 = pH of the external solution 2.

The final equation is: E = K - 0.059 pH

Standardization at pH=7.00 , E = 0 V.

pH 4.00, E= 59.16 mV/pH unit

•Asymmetry potential•E of the 2 reference electrodes

•pH of the internal solution•Liquid junction potential

Measurement of pH (cont.)

Ecell = E°cell - (0.0591)log[H+] + constant

• Ecell is directly proportional to log [H+]

electrode

pH Meters

Glass Membrane Electrode

• Advantages of glass electrode: It can be used in presence of oxidizing, reducing, complexing

• Disadvantage: 1. Delicate, it can’t be used in presence of dehydrating agent e.g. conc.

H2SO4, ethyl alcohol….2. Interference from Na+ occurs above pH 12 i.e Na+ excghange

together with H+ above pH 12 and higher results are obtained.3. It takes certain time to come to equilibrium due to resistance of

glass to electricity.

Junction potential :

a small potential that exists at the interface between two electrolyte solutions that differ in composition.

Development of the junction potential caused by unequal mobilities of ions.

Mobilties of ions in water at 25oC:

Na+ : 5.19 × 10 –8 m2/sV K+ : 7.62 × 10 –8 Cl– : 7.91× 10 –8

To reduce the liquid junction potential to only few millivolts one has to:

1.Use a sat for preparation of the junction which its cation and anion have very near mobilities, so that

they move by the same rate e.g. KCl and KNO3. (K+

=74, Cl- = 73 and NO3- = 76)

2.Use high concentration of the salt for preparation of the bridge, to reduce the effect of difference in rates of migration of other ions in the electrode solutions.

electrode reaction:

Nernst equation

E = -0.059 pHWhen it is connected with NHE as reference

electrode the e.m.f. of the cell :Ecell = zero –(–0.059 pH)

= 0.059 pHpH = E / 0.059

2. Standard Hydrogen Electrode

2balck Pt. H 2e H 2

1

][H

1 log 0.059 - zero E

Disadvantages:-1. It cannot be used in solution containing oxidising agent which

will oxidiose [ ½ H 2 = H+ + e ] or reducing substances which will reduce [ H+ + e = ½ H 2 ] especially in presence of platinum black

2. It cannot be used in reactions involving volatile constituent’s e.g. CO2, as it will be bubbled out by the H2 gas.

3. It cannot be used in presence of catalytic poisons which will affect Pt black which catalyses the electrode reaction.

4. It needs repletion with Pt black.

5. It is not easy to keep H2 gas at one atmospheric pressure during all measurements.

3. Antimony electrode Sb/Sb2O3

Electrode reaction:Sb2O3 + 6 H+ + 6 e = 2 Sb + 3 H2O

Nernst equation

]H][OSb[

[Sb] log

6

0.059 - E E

632

2o

OSb/Sb 32

6

32

2025

][

1log

6

059.0

][

][log

6

0.059 - E E

HOSb

Sbor

E25 = E0 – 0.059 pH

Advantages Easy to use, cheep and

durable.

Disadvantages1. can only be used within

pH range 2 – 8 at lower pH Sb2O3 dissolves and at higher pH Sbo dissolves.

2. It cannot be used in presence of oxidizing agents, reducing agents, complexing agents and noble metals

Quinhydrone electrode

• Advantages1. It is not affected by catalytic poisons.2. Easy to prepare and use.3. It comes to equilibrium rapidly.

• Disadvantages:1. It cannot be used in presence of oxidising agents and

reducing agents 2. The upper limit of the electrode use is pH 8 3. It needs to be used freshly.

Application of potentiometry

1. Direct potentiometric measurementsEobs = Eref + Eja - Eind

sulfate calibration curvey = 14427x - 12024R2 = 0.999

0

200000

400000

600000

800000

1000000

1200000

1400000

0 10 20 30 40 50 60 70 80 90

concentration (ppm )

pe

ak

are

a

by area Linear (by area)

2. Potentiometric titration

It is used for all types of volumetric analysis: acid base, precipitimetry, complexometry and redox

It is used when it is not easy or impossible to detect the end point by ordinary visual methods i.e:

1. For highly coloured or turbid solutions.2. For very dilute solutions 10-3, 10-6 M.3. When there is no available indicator

Potentiometric titration

Titration of 2.433mmol of chloride ion with 0.1000M silver nitrate.

(a)Titration curve. (b)First-derivative curve.(c)Second-derivative curve.

Application of potentiometric titration in

a) Neutralization reactions: glass / calomel electrode for determination of Ph

b) Precipitation reactions: Membrane electrodes for the determination of the halogens using silver nitrate reagent

c) Complex formation titration: metal and membrane electrodes for determination of many cations (mixture of Bi3+, Cd2+ and Ca2+ using EDTA)

d) Redox titration: platinum electrode For example for reaction of Fe3+/ Fe2+ with Ce4+/Ce3+

a) Neutralization reactions:

glass / calomel electrode for determination of pH

• Precipitation reactions: Membrane electrodes for the determination of the halogens using silver nitrate reagent

• Complex formation titration: metal and membrane electrodes for determination of many cations (mixture of Bi3+, Cd2+ and Ca2+ using EDTA)

E,

V vs

Ag

/Ag

Cl,

1M

KC

l

Volume, mL0.0 0.1 0.2 0.3 0.4 0.5 0.6

0.06

0.08

0.10

0.12

E/VE

• Redox titration: platinum electrode For example for reaction of Fe3+/ Fe2+ with Ce4+/Ce3+