Intro to Electroanalytical Chemistry pot

Electro-gravimetry m Coulometric titrations Q = It Main Branches of Electroanalytical Chemistry  Potentiometry: measure the potential of electrochemical cells without drawing substantia

Nov 16, 2004

Introduction to Electroanalytical

Chemistry

Lecture Date: April 27 h , 2008

Reading Material

● Skoog, Holler and Crouch: Ch 22 (An Introduction to

Electroanalytical Chemisty)

● See also Skoog et al Chapters 23-25

● Cazes: Chapters 16-19

● For those using electroanalytical chemistry in their work,

the following reference is recommended:

A J Bard and L R Faulkner, “Electrochemical Methods”, 2 nd

Ed., Wiley, 2001.

Advantages of Electroanalytical Methods

 Matched against a wide range of spectroscopic

and chromatographic techniques, the techniques

of electroanalytical chemistry find an important

role for several reasons:

– Electroanalytical methods are often specific for a

particular oxidation state of an element

– Electrochemical instrumentation is relatively

inexpensive and can be miniaturized

– Electroanalytical methods provide information about

activities (rather than concentration)

History of Electroanalytical Methods

 Michael Faraday: the law

of electrolysis

– “…the amount of a substance deposited

from an electrolyte by the action of a

current is proportional to the chemical

equivalent weight of the substance.”

 Walter Nernst: the Nernst

equation (Nobel Prize

1920)

 Jaroslav Heyrovsky: the

invention of polarography:

(Nobel Prize 1959)

Walter Nernst (1864-1941)

Michael Faraday (1791-1867)

Jaroslav Heyrovsky

Main Branches of Electroanalytical Chemistry

 Key to measured quantity: I = current, E = potential, R = resistance, G =

conductance, Q = quantity of charge, t = time, vol = volume of a standard solution,

m = mass of an electrodispensed species

Interfacial

Static methods

(I = 0)

Dynamic methods (I > 0)

Potentiometry

(E)

Conductometry (G = 1/R)

Controlled potential

Constant current

Voltammetry

(I = f(E))

Amperometric titrations (I = f(E))

Based on Figure 22-9 in Skoog, Holler and Crouch, 6th ed.

Electro-gravimetry (m)

Coulometric titrations (Q = It)

Main Branches of Electroanalytical Chemistry

 Potentiometry: measure the potential of electrochemical

cells without drawing substantial current

– Examples:pH measurements, ion-selective electrodes,

titrations (e.g KF endpoint determination)

 Coulometry: measures the electricity required to drive an

electrolytic oxidation/reduction to completion

– Examples: titrations (KF titrant generation),

“chloridometers” (AgCl)

 Voltammetry: measures current as a function of applied

potential under conditions that keep a working electrode

polarized

– Examples: cyclic voltammetry, many biosensors

Electrochemical Cells

Voltmeter

Salt bridge (KCl)

Zn electrode

Cu electrode

e-Zn  e-Zn 2+ (aq) + 2e

-a Zn 2+= 0.010

Anode

Cu 2+ (aq) + 2e-  Cu(s)

a Cu 2+= 0.010 Cathode

0.010M ZnSO4

solution

0.010M CuSO4 solution

 Zinc (Zn) wants to ionize more than copper (Cu).

 We can use this behavior to construct a cell:

e-reference electrode counter electrode

working electrode

indicator electrode

detector electrode

Potentiometry: Measures equilibrium E

Amperometry: Control E, measures I as function of time

Coulometry: Control E, measure total Q over a period of time

control measurement

Electrochemical Cells and Analytical Methods

Electrochemical Cells

 Galvanic cell: a cell that produces electrical energy

 Electrolytic cell: a cell that consumes electrical

energy

 Chemically-reversible cell: a cell in which reversing

the direction of the current reverses the reactions at

the two electrodes

Conduction in an Electrochemical Cell

 Electrons serve as carriers (e.g moving from Zn

through the conductor to the Cu)

 In the solution, electricity involves the movement of

cations and anions

– In the salt bridge both chloride and potassium

ions move

 At the electrode surface: an oxidation or a

reduction occurs

– Cathode: the electrode at which reduction

occurs

– Anode: the electrode at which oxidation occurs

Oxidation occurs when a chemical species loses an electron.

LEO = lose electron is oxidation

Reduction is when a species gains an electron.

GER = gain an electron is reduction

For example, the chemical reaction

can be decomposed into two half reactions:

“Leo the Lion Says Ger”

Faradaic and Non-Faradaic Currents

Figure 22-2

Mass Transfer occurs by:

Convection Migration Diffusion

 Faradaic (governed by Faraday’s law): direct transfer of

electrons, i.e oxidation at one and reduction at the other

electrode

 Non-Faradaic: increasing charge of the double layer

Electrical charge, q, is measured in coulombs (C) The

charge associated with chemical species is related to the

number of moles through the Faraday constant,

F=96,485.3 (~96,500) C/mole

Electrical current, I, is measured in Amperes (A) Current is

the amount of charge that passes in a unit time interval

(seconds)

Ohm's law relates current to potential (E) through the

resistance (R) of a circuit by E=IR The potential is

measured in Volts (V) and the resistance in Ohms ()

Power (P) is measured in Watts (W = J/s) and is related to

the current and potential by P= IE

The work is measured in Joules (J) and is related to the

The relationship between the standard Gibb's free energy

(EMF), E°(V), is given by

G°=-n F E°

superscript on E0refers to ‘standard state.’

Fundamentals

Fundamentals: The Nernst Equation

● The Nernst equation gives the cell potential E (in volts):

● Q (the activity quotient) is the ratio of products over reactants

as in equilibrium calculations For the generic reaction:

● Q is given by:

● The A’s are activities For low-concentration solutions (low

ionic strengths):

F = faraday (constant)

n = # moles electrons in process

E 0 = standard potential for cell

Electrode Potentials

 The reactions in an electrochemical cell can be

thought of as two half-cell reactions, each with its

own characteristic electrode potential

– These measure the driving force for the

reaction

– By convention, always written as reductions

 Standard electrode potential (E0): the measure

of individual potential of an electrode at standard

ambient conditions (298K, solutes at a

concentration of 1 M, and gas pressure at 1 bar).

Some Standard Electrode Potentials

Reaction E 0 at 298K (Volts)

Cl 2 (g) + 2e -  2 Cl2 +1.359

O 2 (g) + 4H + + 4e -  2 H2O +1.229

Ag + + e -  Ag(s) +0.799

Cu 2+ + 2e -  Cu(s) +0.337

Hg 2 Cl 2 + 2e -  2Hg(l) + 2 Cl2 +0.268

2H + + 2e -  H2(g) 0.000

AgI(s) + e -  Ag(s) + I2 -0.151

Cd 2+ + 2e -  Cd(s) -0.403

Zn 2+ + 2e -  Zn(s) -0.763

See appendix 3 in Skoog et al for a more complete list

The Standard Hydrogen Electrode (SHE)

 A universal reference, but is really a hypothetical

electrode (not used in practice)

– Uses a platinum electrode, which at its surface

oxidizes 2H+ to H2 gas.

– Very sensitive to temperature, pressure, and

H+ion activity

 Because the SHE is difficult to make, the

saturated calomel electrode (SCE) is used

instead

– Calomel = mercury (I) chloride

Q: What is the electrode potential for the

electrode Ag/AgI(s)/I-(0.01 M) ?

The overall reaction for this electrode is

This reaction cannot be found in tables of reduction potentials

But the reaction is comprised of two components

Electrode Potentials

We can initially ignore the fact that the electrode contains

AgI and find E for the silver ion reduction

Electrode Potentials

The Glass pH Electrode

● One of the most common

potentiometric measurements is pH

(a so-called “p-Ion” measurement)

● The common glass pH electrode

makes use of junction potentials to

determine the hydronium ion

concentration in a sample solution

● A typical glass pH electrode is

configured as shown here:

The glass pH electrode is used with a Ag/AgCl reference

electrode For most modern pH electrodes the reference

electrode is incorporated with the pH indicator electrode

A small frit or hole connects the reference electrode and the sample solutions

The Glass pH Electrode

pH Measurements

● A combination pH electrode combines the indicator and

reference into a single unit

● The potential of this cell is:

● where Eij and Eojare the junction potentials at the inner

and outer layers of the glass membrane

● Junction potential: occurs at the interface of two

electrolytes, caused by unequal diffusion rates of cation

and anions across the boundary (e.g the frit in a salt

bridge)

More About pH Measurements

● The surface of the glass is hydrated, which allows

exchange of hydronium ions for the cation in the glass

(sodium or lithium)

● There are four interface regions, the external solution and

hydrated glass, hydrated glass and dry glass on the

outside, dry glass and hydrated glass on the inside, and

hydrated glass and the internal solution

● If the glass is uniform, the two hydrated glass/dry glass

interfaces should be identical and should have the same

junction potential

● Since the glass interface junction potentials then cancel

each other, the junction potential is then the difference

between the internal and external solutions

0.05916pH log

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pH Measurements

pH Measurements

 For a real electrode, the two surfaces will not be identical

and the constant k needs to be determined experimentally

The constant k is termed the asymmetry potential The

constant  is termed the electromotive efficiency

pH Measurements

Q: Why does the pH change the interfacial

potential of the glass/aqueous interface?

A: The motion of the sodium ions leave behind a

negatively charged glass layer that is neutralized to a

lesser or greater extent according to the pH

More explanation about how a pH meter really

works: The sodium ions must move through the dry

part of the membrane and this process is slow For

this reason, the membrane is made very thin Also, a

nonperturbing (low-current) voltmeter is used to read

the cell voltage so that only a few sodium ions must

move through the dry glass in a given time period

pH Measurements

 Errors in pH measurements with glass electrodes arise from

the following effects:

– Calibration problems (e.g drift, or error in the calibration)

– Junction potential

– High [Na+] interacting with electrode

– High acid concentration

– Equilibration time

– Temperature control

 Typical electrodes have the following performance:

– Accuracy = +/- 0.02 pH units

– Precision = +/- 0.002 pH units

pH Electrodes: Errors, Accuracy and Precision

Modern pH electrodes are usually of the "combination"

type, meaning that a single cylinder contains both the

reference electrode, and a glass membrane electrode

Schematically, the total cell may be expressed as

SCE//test solution ([H3O+]=a1)/glass

membrane/[H3O+]=a2, Cl-/AgCl(s)/Ag

The Combination pH Electrode

A Modern Combination pH Electrode

Electrochemical pH Measurements Concluded

Consider a typical problem related to the

use of the combination pH electrode

Recall that

Ecell = L - 0.0592 V pH

what is the pH of a solution for which Ecell is

-0.352 V?

ANSWER: First, find L from the measurement of

the standard:

-0.115 V = L -0.0592 x pH -0.115 V = L -0.0592 x 4.00 Therefore, L = 0.122 V Second, use this value of L to find pH:

-0.352 V = 0.122 V - 0.0592 V x pH

pH = (0.122 V -(-0.352 V))/0.0592

pH = 7.84

QUESTION: What does the pH meter read if the

pH is 7.00 in a 1 M salt solution having 1 M Na+

ions present?

ANSWER:

[H+]obs = 1 x 10-7 + 1 x 10-12

Conclusion the pH meter reads the true pH

under these conditions.

The Ion Selective Electrode (ISE)

● An ISE generally consists

of the ion-selective

membrane, an internal

reference electrode, an

external reference

electrode, and a

voltmeter

● Example: an ISE for

fluoride (F-)

Automatic pKa and log P Determination

pKa (ionization constant) and log P (octanol/water partition) are

important physical parameters that play critical roles in determining how

compounds behave in physiological environments and how they

interact with enzymes, receptors and cell membranes

The Sirius

GLpKa system:

combination pH

electrode

sample tray

reagents liquid

dispensors

 Conductometry: Detection of electrical

conductivity

– Key analytical applications: conductometric detection

in ion-exchange chromatography (IEC or IC) and

capillary electrophoresis (CE)

 Used to detect titration endpoints

Homework Problems (for Study Only)

 Chapter 22:

– 22-1

 Chapter 23:

– 23-11