CH2205 - Electroanalytical Techniques

Electrochemical Cells

Typically a cell is galvanic if it produces electrical energy or electrolytic if it consumes electrical

energy.

A cell is made of two conductors, or electrodes, each immersed in a suitable electrolyte. For

electricity to flow it is necessary that the electrodes are connected externally by a metal conductor

and that the two electrolyte solutions are in contact to permit movement of ions from one solution

to the other.

Two common cell set-ups are;

Electricity is conducted in various ways in the cell. There is a migration of cations and anions as

electrons move from one electrode to the other, at the electrode surfaces an oxidation or reduction

process provides a mechanism by which the ionic conduction of the solution is coupled with the

electron conduction of the electrodes, providing a complete circuit for the flow of electricity.

By definition;

An anode is where oxidation occurs.

A cathode is where reduction occurs.

The liquid junction (frit or salt bridge) is used to avoid the direct reaction of the components of the

two half-cells, there would be a direct deposition of one metal onto the other. There is a small

potential, the junction potential, arises at the interface between two electrolytic solutions of

differing compositions.

Schematic representations of cells are often used to simplify the diagrams using the following

notation;

M(s) І M2+

(aq)||Me2+(aq) І Me(s)

Where each line represents a phase boundary at which a potential may develop, two lines (||)

represents a salt bridge and a long single line (|) represents a frit.

Cell Potentials

Using the example cell;

2AgCl(s)+H2(g)⇋2Ag(s)+2Cl-+2H+

The equilibrium constant for this reaction is;

Salt Bridge

V

Frit

V

K =[H+]2[Cl−]2

p(H2)

To consider the value at any point during the reaction (using instantaneous concentration values)

using Q;

𝑄 =[H+]a

2[Cl−]a2

p(H2)a

The change of the free energy in the cell is given by;

∆G = ∆Geq − ∆Ga

∆G = −RTlnK and ∆Ga = −RTlnQ

∴ ∆G = RTlnQ − RTlnK

The magnitude of free energy for the system depends on how far away the system is from

equilibrium.

Since;

∆G = −nFEcell

−nFEcell = RTlnQ − RTlnK

Ecell =−RTlnQ

nF+

RT

nFlnK

RT

nFlnK = Ecell

o

∴ Ecell = Eo −RT

nFlnQ

Electrode Potentials

A cell is made of two half-cell reactions, conventionally written as reductions (a species gaining

electrons).

To obtain the cell, the second is subtracted from the first to cancel out electrons. To obtain the cell

potential, strictly speaking, it must be done via free energy of the system, however if n is the same

for each half-cell;

Ecell=Ecathode-Eanode

Calculating Half-Cell Potentials

It is rare that a cell will be ‘standard’ and its potential will often have to be calculated via the Nernst

equation.

For the reaction;

pP+qQ+ne-⇋rR+sS

Then;

E = E0 −RT

nFln

[R]r[S]s

[P]p[Q]q

The standard potential is often defined as the electrode potential of a half-cell when all reactants

and products exist at unit activity. It is a physical constant that gives a quantitative description of the

relative driving force of a half-cell reaction.

Potentiometry

Potentiometry has been used for a long time in detecting the end-point during titrimetric analysis.

More recently it has been used for the quantitative analysis of specific ions in solution.

The method simply involves the determination of the potential between two half-cells, if one half-

cell is constant then the potential of the other is effectively being measured and this potential is

related to the ion concentration through the Nernst equation.

The equipment required is similar to that of the above cells, one electrode is replaced with a

standard reference electrode and the other an indicator electrode.

Reference Electrodes

The reference electrode should have a known potential that is constant and insensitive to the

composition of the solution being studied. It is, however, important to remember that the potential

will vary with temperature.

-Calomel electrodes

Represented as follows;

||Hg2Cl2(sat), KCl(xM) ІHg

Where x is the molar concentration of KCl, the electrode reaction is;

Hg2Cl2(s)+2e-⇋2Hg(l)+2Cl-

The potential of this electrode will vary with [Cl-], hence the concentration of KCl must be quoted.

The saturated calomel electrode (SCE) is the most commonly used calomel electrode due to the ease

with which it can be prepared.

At 25oC the SCE has a potential of 0.244V

-Silver/silver chloride electrodes

Analogous to calomel the cell is written as;

||AgCl(s)(sat),KCl(xM)ІAg

The electrode reaction is;

AgCl(s)+e-⇋Ag(s)+Cl-

The potential of the saturated Ag/AgCl at 25OC is 0.199V.

Precaution in the uses of Reference Electrodes

The level of the internal solution must always be above that of the solution in which it is immersed.

This prevents contamination of the electrode or plugging of the junction by reaction of analytes with

Ag(I) or Hg(I) salts. Junction plugging is the most common cause of erratic cell behaviour.

With the level of the electrode solution above the analyte solution some sample contamination may

occur. This is normally so slight it is of little concern, however for the determination of Cl-, K+ or Ag+

the precaution of having a second salt bridge may be necessary. This bridge should contain a high

concentration of non-interfering electrolyte.

Indicator Electrodes

-Metallic electrodes of the first kind

These are in direct equilibrium with the cation derived from the metal electrode;

e.g. Cu2++2e-⇋Cu(s)

Therefore E=E0Cu – 0.0591/2 log10 (1/[Cu2+])

Thus the copper electrode provides a direct measurement of copper concentration in solution.

Similarly Hg, Ag, Cd, Zn and Pb may be used, however not all metals behave reversibly and hence not

all metals are suitable.

-Metallic electrodes of the second kind

A metal electrode can be made responsive to the concentration of anion with which it forms a

precipitate or a stable complex ion.

e.g. AgCl(s)+e-⇋Ag(s)+Cl-

Therefore E=E0AgCl – 0.0591/2 log10 [Cl-]

Thus a convenient way of preparing a Cl- sensitive electrode is to make a pure silver wire the anode

in an electrolytic cell containing KCl. The wire is coated with an adherent AgCl deposit which will

rapidly equilibrate with the surface of the solution in which it is immersed. As the AgCl solubility is

low an electrode may be used for numerous measurements.

An important electrode is that used for the measurement of [edta4-] based on the response of a Hg

electrode when a small amount of Hg(edta)2- is present.

[edta= ]

The half-reaction is;

Hg(edta)2-+2e-⇋Hg(l)+edta4-

Thus E=E0 – 0.0591/2 log ([edta4-]/[Hg(edta)2-])

Hg(edta)2- is very stable and its concentration will not change, giving rise to a new constant.

K=E0+0.0591/2 log10[Hg(edta)2-]

-Metallic electrodes of the third kind

A metal electrode can, under certain circumstances, be made to respond to a different cation. An

example is the determination of the calcium concentration using a mercury electrode.

A small amount of Hg(edta)2- complex is introduced into the solution, hence the potential is given by;

E=K – 0.0591/2 log10[edta4-]

In addition if a small volume of Ca(edta)2- is introduced a new equilibrium is established;

Ca(edta)2-⇋Ca2++edta4-

Kf = Ca2+ [edta4−]

Ca edta 2−

∴ edta4− = Kf

Ca edta 2−

Ca2+

Substituting this into the previous equation for E;

E = K − 0.0591

2log10Kf

Ca edta 2−

Ca2+

𝐸 = 𝐾 −0.0591

2log10Kf Ca edta 2− −

0.0591

2log10(

1

Ca2+ )

Hence if a constant concentration of Ca(edta)2- is used;

E = K′ −0.0591

2log10 1/[Ca2+]

Membrane Electrodes

These may be broadly classified into ion selective and molecular selective electrodes.

-Ion selective electrodes (ISE’s)

Properties;

Minimal solubility – the solubility in analyte solutions should approach zero. Typically membranes

are formed from large molecules or molecular aggregates (e.g. glasses or polymeric resins). Inorganic

compounds of low solubility (e.g. silver halides) can be converted into membranes.

Electrical conductivity – the membrane must have some electrical conductivity (generally via the

migration of ions)

Selective reactivity with analyte – membrane must be able to selectively bind the analyte of interest.

Typical bindings include ion exchange and complexation.

Principles and design;

Cell design – the cell consists of a reference electrode and a membrane electrode (both of which are

dipped into the analysed solution). These two electrodes are connected to a voltmeter. The

membrane electrode consists of an active membrane sealed at one end of a tube, the tube holds a

standard solution of the species to be analysed, X+, and immersed in this standard solution is

reference electrode 2.

The standard solution within the membrane electrode serves a dual purpose. It bathes the internal

surface of the membrane with a fixed concentration of X+ and it also serves as part of the reference

electrode.

For example the internal solution of a calcium selective electrode is typically CaCl2 that is saturated

with AgCl. When a silver wire is dipped into this solution a Ag/AgCl reference electrode is formed. In

addition the Ca2+ ions expose the inner membrane to a constant concentration of analyte.

Electrical conduction – unlike metallic electrodes where conduction occurs via redox processes,

conduction in membrane electrodes occurs via ion transfer. This is made possible by the ionic nature

of the membrane or a species within the membrane. A common ion exchanger for IS membranes is a

silicate glass which consists of a 3D infinite network of oxygen atoms held by silicon atoms. If the

glass has a negative charge this is balanced out by cations residing in open regions of the structure.

Singly charged cations within the glass are mobile enough to provide a mechanism by which

electricity may be carried through the glass.

Conduction across the two glass-solution interfaces of M+ (e.g. H+) will constitute the current. In the

absence of a current the equilibria, on the two glass surfaces, is determined by the pH of the internal

and external solutions. When these equilibria are different the surface where greater dissociation

has occurred will be negative with respect to the other surface. Thus a potential develops (the

boundary potential) whose magnitude depends on the difference in pH on the two sides of the

membrane, this provides the basis for potentiometric pH measurement.

Non-Crystalline Ion Selective Membrane Electrodes

-Glass electrodes

A glass electrode is made by sealing a thin, pH sensitive glass tip to the end of a piece of heavy

walled glass tubing. The resulting bulb is filled with HCl (0.1M) which is saturated in AgCl. A silver

wire is immersed in the solution to form reference electrode 2.

Composition of glass membrane – originally Corning 015 glass (22%Na2O, 6%CaO, 72%SiO2) was

used, this glass shows strong specificity towards H+ upto pH 9, above this value it will start to

respond to Na+.

Hygroscopity of glass membrane – the surface of the glass membrane must be hydrated in order to

have pH activity. Non-hygroscopic glasses, such as a Pyrex and quartz show no pH function.

Dehydrated Corning 015 glass shows little activity but its activity is restored after a few hours

standing in water.

The difference in pH on either side of the glass gives rise to the reading (since the pH inside is fixed

and only the outside varies, it forms a useable electrode).

-Liquid membrane electrodes

These are formed from immiscible liquids that selectively bond certain ions. Membranes of this type

are important as they permit the direct potentiometric determination of several polyvalent cations

and certain singly charged anions and cations.

Early liquid membranes were prepared from immiscible ion exchangers, which were retained in a

porous inert solid support. A porous hydrophobic plastic disk serves to hold the organic layer

Ag

AgCl

HCl

Inner surface of glass;

Binding Site+H+⇋BSH+

Outer surface of glass;

Binding Site+H+⇋BSH+

between the two aqueous layers and a wick action caused the pores of the disk to be filled with

organic liquid contained in an outer reservoir;

Again a potential is set-up, the porous material has a known MXn concentration above binding to

Ca2+ and a varying MXn concentration in the solution outside.

For divalent cation determinations, the inner tube contains a standard solution of MCl2 where M is

the ion to be determined. The solution is also saturated with AgCl to allow the formation of an

Ag/AgCl reference electrode.

As an alternative to the porous disk it is possible to immobilise the ion exchangers in rigid PVC

membranes. Currently most liquid membrane electrodes are of this type.

The active compounds in membranes are of three types;

-cation exchanger; ammonium salts

-anion exchanger; sulphate salts

-macrocyclic compound which complexes a specific ion

One of the most important liquid membrane electrodes is selective towards Ca2+ (in neutral media).

The active species is a cation exchanger consisting of an aliphatic diester of phosphoric acid

dissolved in a polar solvent. The diester contains a single acidic proton and two molecules of it react

with Ca2+ to form a complex in which the two molecule form donor bonds to the Ca2+.

Crystalline Ion Selective Membrane Electrodes

These may be single crystal membranes of polycrystalline heterogeneous membranes. Typically

single crystal homogeneous membranes have better reproducibility, selectivity, linearity of response

and have longer lifetimes.

-Single crystal membranes

Most ionic crystals are insulators or do not have sufficient electrical conductivity to serve as

membrane electrodes. Those that conduct are characterised by having a small ion that is mobile in

the solid phase, for example fluorides of rare earths, Ag+ in silver halides or sulphides and copper I in

Cu2S. Conduction typically occurs by ions jumping to holes (defects) within the crystal lattice. As

these holes are very size specific, these single crystal membranes exhibit excellent specificity.

-The fluoride electrode

LaF3 is an ideal compound for the preparation of membrane electrodes for F- determination.

Membranes are prepared by cutting disks from a single crystal of the compound doped with

Ag

AgCl

Liquid ion exchanger

Porous, plastic hydrophobic

disk

MXn

i.e. CaCl2

europium difluoride to create defects or ‘holes’. Again the mechanism by which a potential is

developed is analogous to the glass electrode.

The magnitude of the charge is dependent on the F- concentration in solution. Thus the side of the

membrane with the lowest F- concentration is positive with respect to the other surface. It is this

difference which allows the measurement in the difference of concentration between the two

solutions.

Most commercial electrodes are rugged, may be used between 0-80oC and give a linear response

down to 20ppb, the only ion that interferes with the electrode is the hydroxide ion and

measurements above a pH of 8 are problematic.

Molecular Selective Electrodes

-Gas sensing probes

A gas sensing probe is not an electrode but a cell containing both ion specific and reference

electrodes. A thin gas permeable membrane allows the gas to pass into the electrode and so be

analysed.

Gas effuses into the pores of the membrane and rapidly establishes an equilibrium. Within the pores

it is also in equilibrium with the internal solution of the electrode.

CO2 ⇋ CO2 ⇋ CO2

External⇋Membrane⇋Internal

Here another equibrium is established that causes the internal pH to change.

CO2+H2O⇋H++HCO3-

The glass electrode immersed in the internal solution detects this change.

The concentration of H+ will be equal to the rate constant times concentration of CO2 divided by the

concentration of HCO3-. If the concentration of HCO3

- is made relatively high (so that it doesn’t

change significantly);

K’=[H+]/[CO2(aq)]

Hence [H+]=K’[CO2(aq)]

Ag/AgCl2

NaCl,NaHCO3

Ag/AgCl2

HCl

pH electrode

Thin glass

Glass permeable membrane

Thus the pH of the internal solution (and the potential across the glass membrane) to the amount of

dissolved CO2 in the sample can be related.

Enzyme Electrodes

By combining the specificity of an ISE and the selectivity of enzyme catalysed reactions we may

obtain apparatus for the determination of compounds of biological and biochemical interest.

To achieve this the enzymes must be immobilised on the surface of the electrode and ideally should

be long-lived and useable for several measurements.

An example is the measurement of Blood Urea Nitrogen (BUN), a routine clinical test. In neutral

media, urea is hydrolysed in the presence of the enzyme urease;

(NH2)2CO + 2H2O + H+ ⇋ 2NH4+ + HCO3

-

The products may then be determined via an ammonium selective electrode. Other electrodes exist

for the determination of glucose, lactose, sucrose, galactose, cholesterol and insecticides.

Applications of Ionic/Molecular Selective Electrodes

-Calibration curves for concentration measurement

An obvious method of correcting potentiometric measurements to give measurements in

concentration is to make an empirical calibration curve (Voltage against log[X]). For this approach to

be a success, it is desirable that the matrix of the standards is as similar as possible to those of the

samples. Sometimes it is helpful to swamp out both standards and samples with an excess of inert

electrolyte. Under these circumstances the additional electrolyte in the sample is negligible. A

commercially available solution, Total Ionic Strength Adjusting Buffer (TISAB), is used for this

purpose.

-Activity vs. concentration

Electrode response is in fact related to activity not concentration. At higher concentrations activity

‘falls off’

-Standard addition

While typically used for colorimetric, atomic absorption and atomic emission spectroscopy analysis

via standard addition in potentiometry is also possible. Standard addition works by measuring the E

value of the unknown solution (with a known volume), adding a small volume of standard solution

(with a concentration much higher than that of the analyte), then measuring the change in E.

Assuming a linear change in log[M+] vs. E the original concentration can then be calculated by;

For the original solution;

E=Eo + 59.1/n log[M+]

For the solution with added standard solution;

E=Eo + 59.1/n log (v1/v2 [M+] + v3/v2 x [Standard])

where v1 is the original volume of solution, v2 is the volume plus the volume of standard and v3 is the

volume of added standard solution.

By subtracting the first equation from the second the concentration of M+ can be calculated.

Voltammetry

Current-Voltage Measurements at Micro-Electrodes

-Micro-electrodes; typically a micro-electrode has a surface area of 1-10mm2, combined with the

fact the typical concentration of an active species is low (<10-4M), the current normally observed at

the electrode is quite small (μA scale). As the microelectrode is small with respect to the volume of

the solution analysed, the quantity of species oxidised or reduced is negligible to the total quantity

and so during the course of analysis the concentration of the species of interest does not change.

Voltammetry is defined as the measurement of a current at a microelectrode as a function of the

potential (vs. a reference electrode) of the micro-electrode.

The measurement is normally carried out in an electrolysis cell containing the analyte solution. A

potential, E, is applied to a microelectrode called the working electrode. This potential is regulated

very precisely and accurately versus a reference electrode. A current flows between the working

electrode and an auxiliary electrode.

The Current-Voltage Curve

The output from a voltammetric instrument is a form of current-voltage curve;

Conventionally, E is plotted with more negative values to the right and a cathodic current

corresponding to a reduction is plotted with increasing values upward. The general name for such a

plot is a voltammogram.

Voltammetry has many forms, the most well known being polarography, another useful analytical

form is anodic stripping voltammetry. Cyclic voltammetry, differential pulsed voltammetry and

square wave voltammetry are routinely used to probe the electrochemistry of newly discovered

Measurement of

Current i Reference electrode

Auxiliary (counter) electrode

Working electrode

Electrolysis cell

Measurement of

Potential

I

V

materials. In addition AC polarography, fast linear sweep voltammetry and differential pulsed anodic

stripping voltammetry are new techniques being increasingly used.

-Polarography

A branch of voltammetry in which the working electrode consists of drops of mercury issuing from

the bottom of a piece of capillary tube (a dropping mercury electrode, DME).

The Hg drop of the DME falls off after every few seconds and is replaced by a new drop. This occurs

naturally or by the help of a drop knocker, the current-voltage curve (or polarogram) is

discontinuous and has a saw-tooth shape.

This shape is explained by the behaviour of the background solution, in which only the solvent and

electrolyte are present (the electrolyte is added to prevent electrostatic forces acting on the

electroactive species in solution ensuring only diffusional processes are involved in the migration of

analyte to working electrode). Only a small residual current will be observed as a negative potential

is applied potentially due to traces of impurity in the cell. A sharp rise will occur at a voltage in which

the first species is reduced, when a second species is added an s-shaped curve is obtained. This

reaction occurs at a potential sufficiently negative to reach a decomposition potential, at which i is

clearly higher than the residual current. This current will then reach a plateau and the different

between this current and the residual current is referred to as the limiting current. The current is

limited by the rate at which the ions of interest are transported to the electrode surface.

When diffusion is the only mechanism of motion (as is preferred), the limiting current is called the

diffusion current, id.

The voltage on the polarogram where i= ½ id is characteristic of the analyte and the media it is

dissolved in. This voltage is the half-wave potential, E1/2. When the analyte forms an amalgam E1/2 is

independent from the concentration of analyte.

-Analysis using polarography

-Qualitative; the value of E1/2 is indicative of the identity of a species and id provides quantitative

information;

Tl (I) Cd(II) Zn(II)

V

I

-Quantitative; The values of id can be used to generate calibration curves for ions and so give

concentrations of species. By measuring the potential at different concentrations a calibration curve

can be formed of id vs. [A], any unknown solution can then have its id measured and read against the

curve to determine concentration.

Standard addition can again be used in polarography, this time however since the relationship is

linear and idα[M+] and so c (in y=mx+c) will equal 0, y=mx. m is equal to the change in y over the

change in x and can be calculated leading to a calculation of the concentration of the ion in question.

-Stripping voltammetry

The analyte is deposited on the surface of a mercury drop by an electrochemical reaction occurring

at a specifically selected potential. This is done for several minutes, the analyte thus concentrated on

the working electrode surface is then stripped off in a few seconds by a rapid shift in potential.

Using a mercury drop, maintained at a potential more negative than E1/2 for ~5 minutes the metal

may be deposited in a concentrated form on the surface layer of the drop. During deposition the

solution is stirred to maximise the metal that is reduced (note only a small percentage of analyte is

reduced so it is important to have control and reproducibility to get accurate results).

After preconcentration, the potential is scanned to a more positive potential at a relatively rapid rate

(25mV/s). Near E1/2 the reverse reaction occurs. As a relatively large amount of analyte has been

preconcentrated a large anodic (negative i) current is observed. It reaches a maxima and decays back

to a low background level.

Stripping voltammetry is much more sensitive than ordinary polarography, and with very long

deposition times (1 hour) analytes wit a concentration of 10-9M can be determined.

Coulometry and Electrogravimetric Determination

-Electrogravimetric

If a well defined product is deposited at the electrode, the weight of the substance may be

measured, which if carried out exhaustively constitutes an electrogravimetric analysis.

In this method the quantity of electricity needed is not measured, only the mass of the product is

important.

This form of analysis is conceptually the most fundamental and simple method of analysis known as

it simply involves the isolation and weighing of the analyte.

-Coulometry

An electrolytic method of analysis where the quantity of electricity required to carry out an

electrolysis is measured. As one mole of electrons is 96485C (a Faraday), for a given half-equation,

we can relate the number of moles of analyte to the number of coulombs required to carry out a

process.

There are two forms of Coulometry, constant-potential coulometry and constant-current

coulometry.

The general equipment required for this process is the same as with voltammetry.

In all cases, the analytical reaction will occur at the working electrode. A reference electrode is

required for the regulation of the potential. Since the solution cannot be used again this form of

analysis must be achieved in one attempt.

The Feasibility of Electrolytic Analysis

An optimum potential for an analysis may have to be determined. This is especially true when other

electroactive species are present in solution with the analyte.

For accurate results the analyte must be reduced (or oxidised) ~100% while the interfering ion must

not be reduced to any significant degree.

If this is the case the Nernst equation can be used. Consider a solution of 0.0005M Cd2+ and an

unknown molarity of Pb2+.

Given the half-equations;

Pb2++2e-⇋Pb(s) Eo=-0.368V (vs. SCE)

Cd2++2e-⇋Cd(s) Eo=-0.645V (vs. SCE)

The lead ion is more easily reduced, thus it is necessary to calculate the percentage of Pb2+

remaining in solution as Cd2+ begins to deposit.

First the potential at which cadmium ions will deposit is calculated;

E = Eo +0.0591

2log Cd2+

E = −0.645 +0.0591

2log 0.005 = −0.713V

At this potential the value of [Pb2+] is;

E = −0.368 +0.0591

2log Pb2+ = −0.713

Pb2+ = 2.11x10−12M

And so the percentage of Pb2+ remaining is;

2.11x10−12

0.005x100 = 4.22x10−8%

Thus there will be a complete separation of lead ions from cadmium ions in any electrolytic assay

performed.

Alternatively voltammetry can give an idea of how viable electrolytic analysis will be, if a plot of

current against voltage shows rises too close together the method will not be feasible, if the steps

are well separated then one can be cleanly reduced before the other, making the method feasible.

-Electrogravimetry

Several metals can be determined by this method; Cu, Cd, Ni, Co, Ag, Bi, Zn and Pb.

Normally a platinum gauze electrode in a cylindrical configuration is employed as a working

electrode. The nature of the deposit is important to the success of the method, the deposit must be

strongly adherent. It should be a fine-grained smooth deposit with a metallic lustre, indicative of a

high quality metal deposit that can withstand subsequent washing, drying and weighing operations

without loss or decomposition.

Several factors influence the quality of the deposit. Stirring rate is one, another being that

complexes of CN- or NH3 give better quality deposits. A low current density (<0.1A/cm2) is beneficial

and the optimum temperature must be experimentally determined.

Hydrogen production at the cathode must be avoided, this is achieved by the addition of anitrate ion

(a cathode depolariser) which is reduced at the cathode to a non-gaseous product;

NO3- + :OH+ + 8e- → NH4

+ + 3H2O

Nitrite ions lead to poor quality deposits and must also be avoided and removed from solution, this

can be done chemically by adding sulphamic acid;

H+ + NH2S(O)2O- + NO2- → N2(g) + HSO4

- + H2O

-Constant-Potental Coulometry

This measures the analyte by determining the quantity of electricity that is involved in reducing or

oxidising the analyte at a working electrode set a constant potential (vs. a reference electrode).

The current is not constant and decreases during the electrolysis (with time less material exists in

the solution to reach the electrode quickly enough to maintain the current). The total change is

obtained by integrating the current-time curve over the time, t, of the reaction.

C = idt

t

0

-Constant-Current Coulometry

This involves an electrolytic reaction that occurs at a constant rate. It is impossible for the analyte to

react with 100% current efficiency at a relatively high constant current. A mediator species is added

at high concentrations that can be oxidised or reduced at the working electrode. This species can

then oxidise or reduce the analyte in solution, away from the electrode.

Consider the analysis of Fe2+ by oxidation to Fe3+, to maintain the current with decreasing Fe2+

concentration a mediator is added, in this case Ce4+ in high concentration.

This reacts with the Fe2+ to form Fe3+ and Ce3+, at the end-point the Ce4+ will no longer be used up,

using an ion-selective electrode to detect the presence of Ce4+ this end-point can be recognised. The

reaction is then halted, the solution stirred and the Ce4+ given a chance to ‘find’ more Fe2+ and

continue reacting, this is repeated until the Ce4+ is no longer used up after stirring.