Transcript Introduction to Electroanalytical Chemistry

Introduction to Electroanalytical
Nov 16, 2004
Chemistry
Lecture Date: April 27h, 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”, 2nd
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)
Michael Faraday
(1791-1867)
Walter Nernst
(1864-1941)
 Jaroslav Heyrovsky:
the
invention of polarography:
(Nobel Prize 1959)
Jaroslav Heyrovsky
(1890-1967)
Main Branches of Electroanalytical Chemistry
Interfacial
methods
Voltammetry
(I = f(E))

Conductometry
(G = 1/R)
Dynamic
methods
(I > 0)
Static methods
(I = 0)
Potentiometry
(E)
Bulk methods
Based on Figure 22-9 in Skoog, Holler and
Crouch, 6th ed.
Controlled
potential
Amperometric
titrations
(I = f(E))
Constant
current
Electrogravimetry
(m)
Coulometric
titrations
(Q = It)
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
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


Zinc (Zn) wants to ionize more than copper (Cu).
We can use this behavior to construct a cell:
Voltmeter
e-
eSalt bridge
(KCl)
Cu electrode
Zn electrode
0.010M CuSO4
solution
0.010M ZnSO4
solution
Zn  Zn2+ (aq) + 2ea Zn 2+ = 0.010
Anode
Cu2+ (aq) + 2e-  Cu(s)
a Cu 2+ = 0.010
Cathode
Electrochemical Cells and Analytical Methods
Potentiometry: Measures equilibrium E
Amperometry: Control E, measures I as function of time
Coulometry: Control E, measure total Q over a period of time
e-
working electrode
indicator electrode
detector electrode
control
measurement
e-
reference electrode
counter electrode
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
“Leo the Lion Says Ger”
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:
Faradaic and Non-Faradaic Currents
Mass Transfer occurs by:
Convection
Migration
Diffusion
Figure 22-2
 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
Fundamentals
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 ().
Fundamentals
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
potential and the amount of charge by work=q E.
The relationship between the standard Gibb's free energy
change, G° (J/mole), and the standard electromotive force
(EMF), E° (V), is given by
G°=-n F E°
where n is the number of electrons transferred and
superscript on E0 refers to ‘standard state.’
Fundamentals: The Nernst Equation
● The Nernst equation gives the cell potential E (in volts):
F = faraday (constant)
n = # moles electrons in process
E0 = standard potential for cell
● 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):
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
E0 at 298K (Volts)
Cl2(g) + 2e-  2 Cl2
+1.359
O2 (g) + 4H+ + 4e-  2 H2O
+1.229
Ag+ + e-  Ag(s)
+0.799
Cu2+ + 2e-  Cu(s)
+0.337
Hg2Cl2 + 2e-  2Hg(l) + 2 Cl2
+0.268
2H+ + 2e-  H2 (g)
0.000
AgI(s) + e-  Ag(s) + I2
-0.151
Cd2+ + 2e-  Cd(s)
-0.403
Zn2+ + 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
Electrode Potentials
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.
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
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
pH Measurements
● A combination pH electrode combines the indicator and
reference into a single unit.
● The potential of this cell is:
● where Eij and Eoj are 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.
pH Measurements
aH + ,glass,in RT
aH + , glass,inaH + ,sol'n,out
aH + ,sol'n,out
RT
RT
=Emem = log
log
log
F
aH + ,sol'n,in F
aH + ,glass,out
F
aH + ,sol'n,inaH + ,glass,out
If the two solutions are identical
a H+ ,glass,in = aH + ,glass,out
aH + ,sol'n,out
RT
Emem = log
F
aH + ,sol'n,in
if the internal solution has a fixed composition, then
RT
RT
Emem = logaH + ,sol'n,out +
logaH + ,sol'n,in = k +  0.05916pH
F
F
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
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 Electrodes: Errors, Accuracy and Precision

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
The Combination pH Electrode
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
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
QUESTION: If Ecell = -0.115 V at a pH of 4.00,
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
liquid
dispensors
reagents
The Sirius
GLpKa system:
combination pH
electrode
sample tray
Conductometry
 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