Electro-Analytical Chemistry

Preliminaries

Electrochemistry involves redox systems.

Therefore species amenable for analysis primarilyare red-ox systems.

Terminology:Red-ox reaction one species undergoes a lossof electrons another gains electrons.

Fe(III) + V(II) Fe(II) + V(III)oxidant reductant

e-

Further the reaction is carried out so that oxidation andreduction occur at different locations - electrodes - inelectrochemical set-ups.

Fe(III) + V(II) Fe(II) + V(III)

Two electrodes when coupled constitute anelectrochemical cell.

The electron movement; q = quantity of charge,i = rate of movement of charge. The electricalpotentials of the electrodes, Eel and the potentialdifference of the electrodes Ecell that are involved canbe measured.

Two types of cells are studied:

a.Galvanic (Voltaic) cell: Galvanic cell usesspontaneous red-ox chemical reactions to produceelectrical energy; that would result in a flow of electrons.G < 0.b. Electrolytic cell: An electrolytic cell decomposeschemical compounds by red-ox processes usingelectrical energy - electrolysis. This is an energydemanding process non-spontaneous. G > 0.In any type of cell: Anode oxidation occurs

Cathode reduction occurs

AnalyticalChemistry

Qualitative Analysisproperty characteristicto analyte

Quantitative analysisproperty related toconcentration of analyte

E

i, q

Cd + 2 AgCl(s) Cd+2 + 2Ag + 2Cl-

Oxidation reduction

i rate of reactionEcell G of reactionG = -nFEcell

q = quantity of current= i t= n F

n = #moles electrons

1 mol e- = 96485C = 1F

Galvanic Cell

i A

Cd + 2Ag+ Cd+2 + 2Ag

Notation: Cd|Cd+2(aq)||Ag+1(aq)|Agphase boundary

V Half Cell Reactions:

Cd (s) + 2 AgCl (s) Cd+2 (aq) + 2Ag (s) + 2Cl- (aq)Cd(s) Cd+2(aq) + 2e anode; oxidation2AgCl(s) + 2e 2Ag(s) + 2Cl-(aq) cathode; reductionThe above two equations (half reactions) involve aphysical transfer of electrons (Faradaic Process)

By convention: anode (oxidation half reaction) - left

i, ampere; A t, sec F = 96485 coulombs/moleR resistance, 96485 C/molE, volts, V

i = q/t q = charge; C

1 C/sec = 1 A (quite a large charge flow rate)

Work (J) = Free energy from reaction = q ECoulombs volts

G = -nF Ecelli = E/R

Construction of electrodes:

The electrode is made out of species involved inthe half reaction. If a metal is not involved, Ptprovides electrical connectivity.

ConventionNegative terminalLeft BlackANODE

ConventionPositive terminalRight RedCATHODE

Standard Hydrogen Electrode: SHE Standard Hydrogen Electrode: SHE

In any cell, when electrons move (current flows) betweenelectrodes the potential difference drops to zero.

Ecell = Difference in electrical potentials of the twoelectrodes.

Negative terminalCOM - Black

Positive terminalRed

Measured Ecell positiveif anode connected tonegative terminal,

pH meter is very closeto an ideal voltmeter.Impedance Draws negligible current.

Negative terminal- reference electrode

slot.- BNC outer connector

Bayonet Neill Concelman connector

Electrodes (Equilibrium) Electrode Potentials:

M(s)

M+

e

M+ (aq) M(s) M+(aq) + e

if Metal acquires a negativepotential w.r.t. solution

The equilibrium set up at the electrodes is not theconventional chemical equilibrium, rather it is anelectrochemical equilibrium.

The potential difference developed across theinterface also controls the equilibrium positionof the half reaction.

This type of equilibrium is also referred to asfrustrated equilibrium.

M(s)

M+

e

M+ (aq) M(s) M+(aq) + e

Depending on the position of equilibriumthe metal acquires a negative/positivepotential w.r.t. solution.

CM+ (aq) is one determinant of theposition of equilibrium.

Faradaic processes: an interfacialphenomenon.

How electrode potentials develops at an electrode:

Ex: Cu/Cu+2(aq)

Electro-chemical potentials of the species Cu+2 ion isnot the same in the two phases, that has not attainedequilibrium. Natural tendency is to equalize theelectrochemical potentials of the two species.In order to achieve such a state, ion concentrationmust change in the two phases at the interface.

Cu(s) Cu+2(aq) + 2e

If the electrochemical equilibrium shifts to right;the excess e- would remain on the Cu metal makingit more negative w. r. t. solution and vice versa.

Electrode Interface :

IHL OHL Bulk

The higher the tendency for the oxidation processto occur, the higher would be the electron density onthe metal.

For cases where the oxidation is dominant itselectrode potential is more negative.

The electrode potentials of red-ox systems aretabulated, relative to the standard hydrogenelectrode (SHE).

Equilibrium Electrode potential:

The magnitude of the electrode potential depends onthe excess charge that exists above that of themetal alone.

If the excess of (negative charge) electrons ispresent in the metal of the electrode the potential ofthe electrode is negative, with the energy of theelectrons in the electrode high and vice versa.

An external power supply is capable of forcing anexcess or a depletion electrons from the metal ofthe electrode, could lead to a non-equilibriumcondition at the interface - electrolysis.

In electrolysis, there is a net reaction forced by theapplied power source later topic.

Absolute individual electrode potentials cannot bemeasured, only potential differences; i.e. only relativevalues can be measured.

Electrode potentials are measured against a standardelectrode and tabulated ; electrode potential of thestandard hydrogen electrode (SHE) is defined as0.00V at 250C.

SHE|| Test electrode

Pt|H2(g) (p=1atm)|H+(aq) (a=1)

Potentiometer

Etest or Eel- +

Ecell = ERHS ELHS

Ecell = Etest ELHS

Ecell = Etest ESHE

Ecell = Etest for this set up.

Ecell = Etest = Eel

Measured Ecell for this set up = Electrode potentialof test electrode

Potentiometermeasures the potentialdifference betweentwo electrodes. = 0, definition

Hg(EDTA)-2(aq, 0.005M)+2eHg(l)+EDTA-4 (aq,0.015M)

aH=1

pH2=1atm

SHE

Potentiometer

Standard Electrodes and Standard ElectrodePotential Eoel:

When the activity (~concentration) of all speciesinvolved in the half reaction is unity in an electrode standard electrode.

Potential of such electrodes are defined as its standardelectrode potential.

SHE|| (Test) Standard electrode

Potentiometer

- +

High Impedance Voltmeter

Electrode reaction E /V Electrode reaction E /V

Li+ + e Li . . . . . 3.045 AgI + e Ag + I . . . . . 0.152 2

K+ + e K . . . . . 2.925 Sn2+ + 2e Sn . . . . . 0.136

Cs+ + e Cs . . . . . 2.923 2H+ + 2e H2 . . . . . 0 exactly

Ba2+ + 2e Ba . . . . 2.92 AgBr + e Ag + Br . . . . + 0.071 1

Al3+ + 3e Al . . . . 1.67 I3 + 2e 3I . . . . . . . + 0.536

Zn2+ + 2e Zn . . . . 0.762 6 Fe3+ + e Fe2+ . . . . . . + 0.771

Ga3+ + 3e Ga . . . . 0.529 Hg22+ + 2e 2Hg . . . . . + 0.796 0

Fe2+ + 2e Fe . . . . 0.44 Ag+ + e Ag . . . . . . . + 0.799 1

Cr3+ + e Cr2+ . . . . 0.424 Hg2+ +2e Hg22+ . . . . . + 0.911 0

Cd2+ + 2e Cd . . . . 0.042 5 Pd2+ + 2e Pd . . . . . . + 0.915

V3+ + e V2+ . . . . 0.255 Cl2 + 2e 2Cl . . . . . . + 1.358 3

Ni2+ + e Ni . . . . . 0.257 Au3+ + 3e Au . . . . . . + 1.52

Note: Large and negative electrode potential meansless tendency for reduction, tendency is to oxidize.Electrode potential, by convention is ameasure of the ability to undergo reduction.

a=1

Reduction assumedTo occur at test electrode

If so Eoel = positive

Hg(EDTA)-2(a=1), EDTA(a=1)/Hg

V>0E-E+

Electrode potential is an interfacial phenomenon.

Separation of charges across metal/solution interfacebrings about the potential difference between thesolution and metal.

The position of equilibrium governed by theactivities (concentrations) of species (by way ofreaction quotient Q) involved in the half reaction;determines the electrode potential and the inherenttendency to undergo reduction.

By convention electrode potentials are expressed forreduction reactions.

If potential is positive, all what it means is that thereduction reaction of the test electrode has a higherpropensity to happen compared to the reduction atthe standard hydrogen electrode.

Calculation of Electrode Potential:

R = 8.314 J/K molT = temperature, Kn= number of electrons involved in half reactionF = 96487 CQel = reaction quotient in activity

pure liquids, solvents, solids; a=1

0 ln el el elRTnF

QE E @ 25o C; volts

Calculation of Electrode Potential (Nernst equation): Calculation of Electrode Potential:

Electrode potentials are calculated for the reductionprocess as reduction potentials Write the halfreaction as a reduction reaction, balanced in massand charge. Write the expression for Q, determinethe # electrons involved;

Substitute in the Nernst Equation.

E.g. MnO4- +8H+ +5e = Mn+2 + 4H2O

2

2

4

24 4

0/ / 8

ln5MnO Mn MnO M

Mn

Mn

nO H

RTEF

a

aE

a

0 ln el el elRTnF

QE E

Tabulation

Eg. Hg(EDTA)2-(aq)+ 2e = Hg(l) + EDTA4- (aq)

4 4

2 2

0 0ln ln2 2

Hg EDTA EDTAel el el

HgEDTA HgEDTA

a a aRT RTE E EF a F a

Calculation of Electrode Potential (Nernst equation):Calculation of Cell Potential:

Cell potentials are calculated for redox reaction.Write the reaction, balanced.Write the expression for Q of the reaction, recognizethe # electrons involved; for the reaction.

Substitute in the Nernst Equation.

0 ln rxcell cell n

RTF

QE En

E0cell = E0red E0oxd

E0cell = E0+ E0-E0cell = E0MnO4-/Mn+2 E0Fe+3/Fe+2

4

Tabulated Electrode Elect Pot Reaction Pot

8 5MnO H 2 24 4

3 2 3 2

2 0 02

2 3 0 0

4E EE - E/ /

/ /

MnO Mn MnO Mn

Fe Fe Fe Fe

e Mn H O

Fe Fe e

Calculation of E0cell

MnO4- +8H+ +5 Fe+2 = Mn+2 + 5Fe+3 + 4H2O

Calculation of Cell Potential:

Eg. MnO4- +8H+ +5e = Mn+2 + 4H2OFe+2 = Fe+3 + e

Overall (all aqueous species)MnO4- +8H+ +5 Fe+2 = Mn+2 + 5Fe+3 + 4H2O

0 lncell cell rxnRTE E QnF

2 3

24

50

8 5ln5Mn Fe

cell cellMnO H Fe

a aRTE EF a a a

Ecell = Ered process Eoxd process

Ecell = E+ E-

Ecell can be calculated by calculation each electrodepotential (reduction) separately and subtracting thepotential at oxidation half from the reduction half.

0 lncell cell rxnRTE E QnF

2 3

24

50

8 5ln5Mn Fe

cell cellMnO H Fe

a aRTE EF a a a

Electrons move from less positive electrode potentialelectrode to more positive electrode potentialelectrode, when connected by a conductor.

Alternate view of some electrodes:

Ag/AgCl(s), HCl(aq)

AgCl(s) + e = Ag(s) + Cl-(aq)

Ag+(aq) + e = Ag(s)

0/

0.05916 log1

el AgCl Ag ClE E a

,0 0/ /

0.05916 1 0.05916log log1 1

sp AgClel Ag Ag Ag AgAg Cl

KE E E

a a

Alternate view of some electrodes:

Pb/PbF2(s), HF(aq)

PbF2(s), + 2e = Pb(s) + 2F-(aq)

Pb+2(aq) + 2e = Pb(s)

0 22/

0.05916 log2

el PbF Pb FE E a

, 20 02/ 2 / 2

0.05916 1 0.05916log log2 2

sp PbFel Pb Pb Pb PbPb F

KE E E

a a

The fact that the electrode potential (and the cellpotential) is dependent on the concentration ofspecies allows the use of electrodes as chemicalprobes.

The electrode potential at an electrode measuresthe propensity of a reduction reaction to occur;at the concentrations (activities) of the species thathas attained an electrochemical equilibrium atthe interface.

Use of electrodes in Analytical Chemistry

Not a viablecell configuration.

Zn Cu

ZnCl2 (aq)

Cu (NO3)2 (aq)

Electrode Interface : Study of charge transfer reactions is the goal of mosttechniques electrochemistry. The interfacial system ofis very complex and even in the absence of electrontransfer and processes other than electron transfer dooccur. Such processes can affect the electrical doublelayer and therefore the electrode behavior.

Such processes include phenomena such as adsorption,desorption, and charging of the interface as a result ofchanging electrode potential.

These are called non-faradaic processes.

The SHE is cumbersome to construct. Other half-cellsare being used as secondary standards.

Silver/Silver chloride/KCl referenceAgCl(s) +e- = Ag(s) + Cl- (aq)

Silver/silver chloride:AgCl(s) +e- = Ag(s) + Cl- (aq)

Potential @ 25vs. SHE vs. SCE

Ag/AgCl, KCl (0.1M) 0.288 0.047Ag/AgCl, KCl (3M) 0.210 -0.032

Ag/AgCl, KCl (3.5M) 0.205 -0.039Ag/AgCl, KCl (sat'd) 0.199 -0.045Ag/AgCl, NaCl (3M) 0.209 -0.035

Ag/AgCl, NaCl (sat'd) 0.197 -0.047

Calomel:Hg2Cl2(s) + 2e- = 2Hg + 2Cl-(aq)

Hg

Hg2Cl2(s)

KCl, Hg2Cl2 (aq,sat),

KCl(s)

frit

Calomel:Hg2Cl2(s) + 2e- = 2Hg + 2Cl-(aq)

Potential @ 25vs. SHE vs. SCE

Hg/Hg2Cl2, KCl (0.1M) 0.334 0.0925Hg/Hg2Cl2, KCl (1M)

NCE (Normal Calomel) 0.280 0.0389

Hg/Hg2Cl2, KCl (3.5M) 0.250 0.006Hg/Hg2Cl2, KCl (sat'd)SCE (Sat'd Calomel) 0.241 0

Hg/Hg2Cl2, NaCl (sat'd)SSCE 0.2360 -0.0052

Mercury/mercurous sulfate:Hg2SO4 (s) + 2e- = 2Hg + SO4-2 (aq)

Potential @ 25

vs. SHE vs. SCE

Hg/Hg2SO4, H2SO4 (0.5M) 0.682 0.441

Hg/Hg2SO4, H2SO4 (1M) 0.674 0.430

Hg/Hg2SO4, K2SO4(sat'd) 0.64 0.40