Potentiometry

Common types of electrochemical measurements: 1. Potentiometry: Measurement of a potential (voltage) at an electrode (relative to some reference) in the absence of current flow.

R eference Ind ica to r

H igh im pedance vo ltm ete r

G alvan ic (o r vo lta ic ) cell 2. Amperometry: Measurement of a limiting current at a constant potential.

Reference Indicator

Low impedance current meter

+ -

Fixed potential

Electrolytic cell

3. Coulometry: Measurement of the quantity of electrical charge needed to convert an analyte from one oxidation state to another.

Reference Indicator

Coulometer

+ -

Fixed potential

Electrolytic cell

?Idt=Charge (Q)

4. Voltammetry: Measurement of current as a function of applied potential.

Reference Indicator

Low impedance current meter

+ -

Variable potential

1. Potentiometry: Measurement of a potential (voltage) at an electrode (relative to some reference) in the absence of current flow.

R eference Ind ica to r

H igh im pedance vo ltm ete r

G alvan ic (o r vo lta ic ) cell

reference electrode|salt bridge|analyte solution|indicator electrode

Direct potentiometry involves the direct measurement of a voltage generated at an electrode in solution relative to a potential of a reference electrode. Potentiometry is used to measure concentrations of specific anions and cations, sometimes in conjunction with a titration… Ecell = EInd – Eref + EJ

[ ]Xn

LpXn

LEInd log0592.00592.0+=−=

for metal indicator electrodes, L is usually E°, while for membrane electrodes, L is a collection of constants

[ ] ( )( )n

LEEEXpX refJcell

/0592.0log

+−−−=−=

[ ] ( ) ( )

0592.0/0592.0log KEn

nKEXpX cellcell −

−=−

−=−= for cations

[ ] ( ) ( )0592.0/0592.0

log KEnn

KEApA cellcell −=

−=−= for anions

Potentiometric titrations: e.g. acid/base with a pH electrode

or a redox titration

Ecell = Eind - Eref + Ej Reference electrodes: All Eº are reported relative to the Standard Hydrogen Electrode (SHE). However, this electrode is often inconvenient to use. Common alternatives: The saturated calomel electrode (SCE): Hg|Hg2Cl2(sat’d), KCl (x M)|| E°=0.2444 V @ 25 °C

Silver/silver chloride reference electrode: Ag|AgCl(sat’d), KCl (sat’d)|| E°=0.199 V @ 25 °C

The liquid junction potential, Ej: A potential is developed across a boundary between electrolyte solutions of different composition:

KCl is usually used in salt bridges because the mobilities of K+ and Cl- are similar. Therefore, only very modest liquid junction potentials are developed at the interfaces between the salt bridge and the other solutions.

Indicator electrodes: Metal indicators electrodes: 1. Pure metal electrode that is in equilibrium with its cation in solution Xn+(aq) + ne- X(s) such that,

[ ] [ ]++° +=−= ++

nXXnXXInd X

nE

XnEE nn log0592.01log0592.0

//o

[ ] pXn

EXn

EE XXn

XXInd nn

0592.0log0592.0// −=+= ++

+ oo

2. Pure metal electrode that responds to anions that form sparingly soluble precipitates or stable complexes with the electrode cations. e.g., AgCl(s) + e- Ag(s) + Cl-(aq) E°=0.222 V

[ ] pClECln

EE AgAgClAgAgClInd 0592.0log0592.0// +=−= −° o

3. Inert metallic electrodes: Inert conductors that themselves do not engage in electrochemical reactions under the conditions in which a redox reaction of interest occurs. e.g. platinum, gold, palladium, carbon

Membrane electrodes: Example: the glass pH electrode

The glass membrane permits H+(aq) to exchange with Na+ in the silicate structure: H+(aq) + Na+Gl- Na+(aq) + H+Gl-

The boundary potential arises from the equilibria established at the interior and exterior surfaces of the glass electrode: H+Gl1(s) H+(aq) + Gl1

-(s) where surface 1 is between the exterior glass and the analyte solution H+Gl2(s) H+(aq) + Gl2

-(s) where surface 2 is between the interior glass and the internal solution Both surfaces develop a negative charge but the net charge at each surface is dependent upon the pH of the analyte solution at the exterior of the glass and the pH of the interior solution at the interior of the glass. Even when the pH values of the solutions on either side of the glass are equal, a small potential, referred to as the Asymmetry Potential, is observed. (Arises due to the fact that the two surfaces are not identical.) Due to the existence of the asymmetry potential, pH electrodes must be calibrated frequently.

SCE||[H+]analyte|Glass membrane|[H+]reference, [1 M Cl-], AgCl|Ag ESCE, EJ E1 E2 EAg,AgCl SCE = reference electrode 1 Ag,AgCl = reference electrode 2 glass membrane + reference electrode 2 = indicator electrode Changes between EInd and ESCE arise from changes in [H+].

[ ][ ]reference

analyteb H

HEEE +

+

=−= log0592.021

[ ][ ]reference

analyteb H

HEEE +

+

=−= log0592.021

[H+]reference is constant, so

[ ] pHLHLE analyteb 0592.0log0592.0 −′=+′= +

where L’ = -0.0592log[H+]reference Potential of the glass electrode includes the small asymmetry potential, Easy, EInd = Eb + EAg/AgCl + Easy EInd = L – 0.0592pH where L = L’ + EAg/AgCl + Easy L cannot be determined theoretically because Easy is unknown… Therefore, use of calibrations are essential for pH measurements with a pH electrode.

NIST as well as international scientific organizations have established an operation definition of pH. One or more NIST buffers are used to calibrate the pH electrode at various specified pH values. The meter reading Es is adjusted in accord with the pHs value (pH of standard). The glass electrode in the unknown solution gives a potential, Eu, and the pH of the unknown solution (pHu) is: pHu = pHs – (Eu-Es)/0.0592 Note that the pH electrode is inaccurate at low and high pH

A variety of other ion-selective electrodes are in use: Liquid-membrane electrodes:

polyvalent cations, some anions

Crystalline-based membranes: mostly for anions

Ion-sensitive field effect transistors: solid state semiconductor electrodes for various ions Gas permeable membranes for determination of dissolved gases (e.g., O2 and CO2 in blood):