Chapter 22, 23, 24 & 25
Electroanalytical Chemistry
Why electroanalytical chemistry?
Electroanalytical methods have certain advantages over other analytical methods.
Electrochemical analysis allows for the determination of different oxidation states of an
element in a solution, not just the total concentration of the element. Electroanalytical
techniques are capable of producing exceptionally low detection limits and an abundance
of characterization information including chemical kinetics information. The other
important advantage of this method is its low cost.
History:
Polarography was first discovered by a czechoslovavian chemist by the name of
Heyrovsky in 1920. He won the Nobel prize for it in 1959. He proposed that the current
recording generated by a oxidation or reduction in a cell as the A.P. is continuously
increased:
Ox + ne  Red
Oxygen Probe:
A.P. – 650 mV Ag|AgCl
Reactions:
O2 + 2H2O + 4e-  4OH4Ag + 4Cl-  4AgCl + 4e-
Hydrogen Probe:
A.P. +650 mV
Reactions:
H2O2  O2 +2H+ +2e2Ag+ + 2e-  2Ag
The following comprehensive chart below shows the applications of electrochemistry and
its interactions with other branches of science and technology.
Electrochemical Cells
-consists of 2 conductors called electrodes that are connected externally by means
of a metal conductor with 2 electrolyte solutions in contact in order to permit ion
movement
-figure 22.1 below shows an example of a typical cell:
-the two halves of the cell are joined by a salt bridge. This is a tube that contains
saturated potassium chloride, or other electrolyte, which allows electrical contact between
the cells without physical contact.
22A-1. Conduction in a Cell:
-charge is conducted by three distinct processes:
1) in the electrodes and the external conductor, electrons
serve as carriers
2)within the solution the flow of electricity involves the
migration of both the cations and the anions
3)oxidation and reduction occurs at the 2 electrode surfaces
22A-2 Solution Structure- The Double Layer
It is very important to realize that electrochemical measurements involve
heterogeneous systems because an electrode can only donate or accept electrons from a
species that is present in a layer of solution that is immediately adjacent to the electrode.
22A-3 Faradaic and Nonfaradiac Currents
Two types of processes can conduct currents across and electrode solution interface. One
kind involves a direct transfer of electrons via and oxidation reaction at one electrode and
reduction reaction at the other. Processes of this type are called faradaic processes
because they are govern by Faraday’s Law, which states that the amount of chemical
reaction at an electrode is proportional to the current; the resulting currents are called
faradaic currents.
22A-4 Mass Transfer in Cells with the Passage of Current
-
three mechanisms bring about these mass transfer: convection,
migration, and diffusion. Convection involves mechanical motion of
the solution as a result of stirring or the flow of the solution past the
surface of the electrode. Migration is the movement of ions through
the solution brought about by electrostatic attraction between the ions
and the charged electrode. Diffusion is the motion of species brought
about by a concentration gradient.
22A-5 Galvanic and Electrolytic Cells:
-the net reaction in a cell is the sum of the two half reactions
-the potential is the measure of the tendency of the cell to
move towards equilibrium
-Galvanic cells react in a way that produces electrical energy
-Electrolytic cells consume energy
-Chemically reversible cell exists when reversing the
direction of current reverses the reaction at the electrodes
22A-6 Anodes and Cathodes:
-Cathode- the electrode where reduction occurs in an
electrochemical cell
-Anode- the electrode where oxidation occurs in an
electrochemical cell
-reactions at cathodes
-electrons supplied by external circuit via an inert
electrode (platinum or gold)
-some examples are:
Cu2+ + 2e- <=> Cu(s)
Fe3+ + e- <=> Fe2+
2H+ + 2e- <=> H2(g)
AgCl(s) + e- <=> Ag(s) + Cl-reactions at anodes
-some examples are:
Cu(s) <=>
Cu2+ + 2e-
Fe2+ <=> Fe3+ + eH2(g) <=> 2H+ + 2eAg(s) + Cl- <=> AgCl(s) + e-
22A-7 Cells without Liquid Junctions:
-liquid junction - the interface between 2 different
electrolytic solutions
-cell can contain more than one
-a small Junction Potential arises at these interfaces
-sometimes it is possible to prepare cells that share a
common electrolyte to avoid this problem
22A-8 Schematic Representation of Cells:
-anode and information on the solution it is contacting on left
-single vertical line indicates phase boundary where
potential might develop
-two vertical lines indicates liquid junction
-concentration or activity put in parentheses
-example:
Zn/ZnSO4(azn2+ = 1.00)//CuSO4(acu2+ + 1.00)/Cu
Potentials in Electroanalytical Cell:
-we will deal mainly with activities rather than concentration
therefore:
ax = fx [X]
where fx= activity coefficient of solute X and [X] is molar
concentration
-equilibrium constant for a reaction (a + b <=> c + d + e)
K =ac x ad x ae/pb x aa
where a is the activity and p is the pressure in atmospheres
-a pure solid at unity gives this equation:
K = ac x ad/pb
-the second quantity Q:
Q= (ac)i(ad)i/(pb)i
-the change in free energy for the cell reaction:
G= RT ln Q -RT ln K
-cell potential:
G =-nFEcell
where n = # of moles and F = faraday(96,485C/mole of electrons)
Ecell:
Ecell = -RT/nF lnQ + RT/nF ln K
= E0cell - RT/nF ln (ac)i(ad)i/(pb)i
-standard electrode potential:
E0cell = RT/nF ln K
Electrode potentials:
Ecell = E(cathode) - E(anode)
-potentials are measured by difference, they are relative not
absolute
-SHE =Standard hydrogen electrode-1 atm 0volts at all temps.
-NHE = normal hydrogen electrode
-IUPAC says electrode potential refers to reduction 1/2 reactions
-in dilute solutions molar concentration can be used for
computations, rather than activities
-shifts in equilibria cause a shift in cell potential
-to compensate for activity effects and side reactions we can
use (Ef) formal potential instead of standard electrode potential
Calculation of Cell Potential from Electrode Potentials:
-calculated potentials are sometimes called thermodynamic
potentials
Ecell = E(cathode) - E(anode)
-negative Ecell indicates non-spontaneity of reaction
-Liquid Junction Potentials can be calculated from the knowledge
of the mobility of the two ions involved, but it is rare that the
system in question is simple enough for this computation
Currents in Electrochemical Cells:
-Ohms law is usually obeyed:
E=IR
where E is the potential difference in volts responsible for the
movement of the ions, I is the current in amps, and R is the
resistance in ohms of the electrolyte to the current
Effect of current on cell potential:
-can result in
1) reduced potential of galvonic cell
2) increased potential needed to develop current in an
electrolytic cell:
-this is due to ohmic resistance and polarization effects such
as: charge transfer over-voltage, crystallization over-voltage
-to account for an ohmic potential, or IR drop:
Ecell = E(cathode) - E(anode) – IR
Polarization:
-concentration polarization-when the mass transfer limits
the rate of the reaction and therefor the current
-reaction polarization-when the rate of formation of the
intermediate limits the current
-when a physical process limits the reaction it is said to be
adsorption, desorption, crystallization etc. polarization
-charge transfer polarization-when rate of electron transfer
from electrode to oxidized species or reduced species to
electrode limits the current
-the degree of polarization is measured by the overvoltage n
n = E - Eeq
where E is electrode potential and Eeq is thermodynamic or
equilibrium potential and E<Eeq
-It is necessary that the reactant be brought to the surface
of the solution from the bulk at a rate of:
I = dQ/dt = nFdc/dt
where dQ/dt is the rate of flow of the electrons in the
electrode, n is the # of electrons in the 1/2 reaction, and
F is the faraday
-the rate of concentration change is:
dc/dt = AJ
where A=surface area of electrode(m2) and J= the
concentration flux in mol/s.m2
I = nFAJ
Ecell = E(cathode) - E(anode) - IR + n(cathode) + n(anode)
-the rate of diffusion:
dc/dt = k(c -c0)
where k=proportionality constant, c=reactant concentration,
and c0=equilibrium concentration.
Summary:
To summarize, concentration polarization is observed when diffusion, migration and
convection are insufficient to transport the reactant to or from an electrode surface at a
rate demanded by the theoretical current.
The following diagram summarizes the procedures for the voltammetric analysis of
various types of water.
The above diagram illustrates the importance of derivative polarogram in clarifying a
regular polarogram recording.
Some very important types of Polarography are shown below:
LCEC:
The illustration below depicts the comparison of the detection ranges of a number of
different techniques including Polarometry.
The following table shows the application of the Electrochemical Stripping Analysis:
WHAT IS POTENTIOMETRY?
Introduction
Potentiometry is the field of electroanalytical chemistry in which potential is measured
under the conditions of no current flow. The measured potential may then be used to
determine the analytical quantity of interest, generally the concentration of some
component of the analyte solution. The potential that develops in the electrochemical cell
is the result of the free energy change that would occur if the chemical phenomena were
to proceed until the equilibrium condition has been satisfied.
Grxn = -nFErxn
This concept is typically introduced in quantitative analysis courses in relation to
electrochemical cells that contain an anode and a cathode. For these electrochemical
cells, the potential difference between the cathode electrode potential and the anode
electrode potential is the potential of the electrochemical cell.
Ecell = Ecathode – Eanode
If the reaction is conducted under standard state conditions, this equation allows the
calculation of the standard cell potential. When the reaction conditions are not standard
state, however, one must utilize the Nernst equation to determine the cell potential.
Ecell = Eo – RT/nF(ln(Keq)
Physical phenomena which do not involve explicit redox reactions, but whose initial
conditions have a non-zero free energy, also will generate a potential. An example of this
would be ion concentration gradients across a semi-permeable membrane. This can also
be a potentiometric phenomena, and is the basis of measurements that use ion-selective
electrodes.
Emem = (constant) – RT/ziF(ln(ai)
WHAT IS COULOMETRY?
Introduction
Coulometry is an analytical method for measuring an unknown concentration of an
analyte in solution by completely converting the analyte from one oxidation state to
another. Coulometry is an absolute measurement similar to gravimetry or titration and
requires no chemical standards or calibration. It is therefore valuable for making absolute
concentration determinations of standards.
Coulometry uses a constant current source to deliver a measured amount of charge. One
mole of electrons is equal to 96,485 coulombs of charge, and is called a faraday.
Schematic of a coulometric cell
Coulometric Titration
Due to concentration polarization it is very difficult to completely oxidize or reduce a
chemical species at an electrode. Coulometry is therefore usually done with an
intermediate reagent that quantitatively reacts with the analyte. The intermediate reagent
is electrochemically generated from an excess of a precursor so that concentration
polarization does not occur. An example is the electrochemical oxidation of I- (the
precursor) to I2 (the intermediate reagent). I2 can then be used to chemically oxidize
organic species such as ascorbic acid.
The point at which all of the analyte has been converted to the new oxidation state is
called the endpoint and is determined by some type of indicator that is also present in the
solution. For the coulometric titration of ascorbic acid, starch is used as the indicator. At
the endpoint, I2 remains in solution and binds with the starch to form a dark purple
complex.
The analyte concentration is calculated from the reaction stoichiometry and the amount of
charge that was required to produce enough reagent to react with all of the analyte.
WHAT IS VOLTAMMETRY?
Voltammetry refers to the measurementof current that results from the application of
potential. Unlike potentiometry measurements, which employ only two electrodes,
voltammetric measurements utilize a three electrode electrochemical cell. The use of the
three electrodes (working, auxillary, and reference) along with the potentiostat
instrument allow accurate application of potential functions and the measurement of the
resultant current. The different voltammetric techniques that are used are distinguished
from each other primarily by the potential function that is applied to the working
electrode to drive the reaction, and by the material used as the working electrode.
What is Cyclic Voltammetry (CV)?
Cyclic voltammetry (CV) is an electrolytic method that uses microelectrodes and an
unstirred solution so that the measured current is limited by analyte diffusion at the
electrode surface. The electrode potential is ramped linearly to a more negative potential,
and then ramped in reverse back to the starting voltage. The forward scan produces a
current peak for any analytes that can be reduced through the range of the potential scan.
The current will increase as the potential reaches the reduction potential of the analyte,
but then falls off as the concentration of the analyte is depleted close to the electrode
surface. As the applied potential is reversed, it will reach a potential that will reoxidize
the product formed in the first reduction reaction, and produce a current of reverse
polarity from the forward scan. This oxidation peak will usually have a similar shape to
the reduction peak. The peak current, ip, is described by the Randles-Sevcik equation:
ip = (2.69x105) n3/2 A C D1/2 v1/2
where n is the number of moles of electrons transferred in the reaction, A is the area of
the electrode, C is the analyte concentration (in moles/cm3), D is the diffusion coefficient,
and v is the scan rate of the applied potential.
The potential difference between the reduction and oxidation peaks is theoretically 59
mV for a reversible reaction. In practice, the difference is typically 70-100 mV. Larger
differences, or nonsymmetric reduction and oxidation peaks are an indication of a
nonreversible reaction. These parameters of cyclic voltammograms make CV most
suitable for characterization and mechanistic studies of redox reactions at electrodes.
What is Polarography?
Introduction
Polarography is an voltammetric measurement whose response is determined by
combined diffusion/convection mass transport. Polarography is a specific type of
measurement that falls into the general category of linear-sweep voltammetry where
the electrode potential is altered in a linear fashion from the initial potential to the final
potential. As a linear sweep method controlled by convection/diffusion mass transport,
the current vs. potential response of a polarographic experiment has the typical
sigmoidal shape. What makes polarography different from other linear sweep
voltammetry measurements is that polarography makes use of the dropping mercury
electrode (DME). Close-up of the DME.
Picture
A plot of the current vs. potential in a polarography experiment shows the current
of a
oscillations corresponding to the drops of Hg falling from the capillary. If one
DME
connected the maximum current of each drop, a sigmoidal shape would result. The
limiting current (the plateau on the sigmoid), called the diffusion current because
diffusion is the principal contribution to the flux of electroactive material at this point
of the Hg drop life, is related to analyte concentration by the Ilkovic equation:
id = 708nD1/2m2/3t1/6c
Where D is the diffusion coefficient of the analyte in the medium (cm2/s), n is the
number of electrons transferred per mole of analyte, m is the mass flow rate of Hg
through the capillary (mg/sec), and t is the drop lifetime is seconds, and c is analyte
concentration in mol/cm3.
There are a number of limitations to the polarography experiment for quantitative
analytical measurements. Because the current is continuously measured during the
growth of the Hg drop, there is a substantial contribution from capacitive current. As the
Hg flows from the capillary end, there is initially a large increase in the surface area. As a
consequence, the initial current is dominated by capacitive effects as charging of the
rapidly increasing interface occurs. Toward the end of the drop life, there is little change
in the surface area which diminishes the contribution of capacitance changes to the total
current. At the same time, any redox process which occurs will result in faradaic current
that decays approximately as the square root of time (due to the increasing dimensions of
the Nernst diffusion layer). The exponential decay of the capacitive current is much more
rapid than the decay of the faradaic current; hence, the faradaic current is proportionally
larger at the end of the drop life. Unfortunately, this process is complicated by the
continuously changing potential that is applied to the working electrode (the Hg drop)
throughout the experiment. Because the potential is changing during the drop lifetime
(assuming typical experimental parameters of a 2mV/sec scan rate and a 4 sec drop time,
the potential can change by 8 mV from the beginning to the end of the drop), the charging
of the interface (capacitive current) has a continuous contribution to the total current,
even at the end of the drop when the surface area is not rapidly changing. As such, the
typical signal to noise of a polarographic experiment allows detection limits of only
approximately 10-5 or 10-6 M. Better discrimination against the capacitive current can be
obtained using the pulse polarographic techniques.
Qualitative information can also be determined from the half-wave potential of the
polarogram (the current vs. potential plot in a polarographic experiment). The value of
the half-wave potential is related to the standard potential for the redox reaction being
studied.
References:
http://www.anachem.umu.se/jumpstation.htm
http://userwww.service.emory.edu/~kmurray/mslist.html
http://www.anachem.umu.se/jumpstation.htm
http://www.acs.org
http://www.chemcenter/org
http://www.sciencemag.org
http://www.scimedia.com
http://www.kerouac.pharm.uky.edu/asrg/wave/wavehp.html