advanced analytical chemistry lecture 17chemistry.unt.edu/~tgolden/courses/lecture 17 potential...
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Advanced Analytical Chemistry
Lecture 17
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The electrochemical cell contains the electrodes and a solvent and supporting
electrolyte.
How do you choose the best solution for the experiment?
Must consider the electrochemical as well as chemical reactions.
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Ideal Solvent System Properties
Electrochemically Inert (stable over a wide range of potentials)
High electrical conductivity
Good solvent power
Chemical Inertness
Available in pure form
Low cost
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Electrochemically Inertness
For a solvent system to be stable over a wide range of potentials – it cannot
undergo any electrochemical reaction during oxidation or reduction of the
analyte.
Solvent (solution) Breakdown Potential
Potential of an electrochemical reaction of the solvent system.
Examples: Water, Tetrabutylammonium hexafluorophosphate in ACN (+3.4 to -2.9 V vs
SCE), Nitrobenzene (hard to oxidize easy to reduce)
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Electrochemically Inertness
In addition the protic character of the solvent is important because the
electrochemical intermediates (radical anions) frequently react with
protons.
While this cannot be avoided in aqueous solvents (and the intermediate is fast)
it is important in organic electrochemistry.
Typical SOFC Test Station Chem 5570
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Electrochemically Inertness
Protic solvents are strong hydrogen-bond donors, exchange protons rapidly,
and have hydrogen bound to more electroactive atoms (i.e. F, O, N)Examples: water, methanol, formamide, ammonia, hydrogen fluoride
Aprotic solvents generally have hydrogen bound only to carbon, are poor
hydrogen-bond donors, weakly acid, and have slow proton exchange.Examples: ethers, acetone, nitrobenzene, acetonitrile, dimethyl sulfoxide dimethylformamide
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High electrical conductivity
Solvent system must support the passage of an electric current – have low
electrical resistance and high dielectric constant (>10).
A large dielectric constant promotes dissociation of an ionic solute and
thus lowers solution resistance.
Examples: Water 80.1
for organic electrochemistry – DMF 36.7, ACN 37.5, DMSO 46.7
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Good solvent power
Able to dissolve a wide range of substances at
acceptable concentrations (~0.1M for electrolytes and
0.1-1 mM for analytes)
Example of salts which dissolve well in organic solvents:
hexafluorophosphates, tetrafluoroborates, ethyltributylammonium
tetrafluoroborate, tetrabutylammonium hexafluorophosphate.
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Good solvent power
Viscosity is also important.
Low viscosity is better so that mass transport by diffusional control can last
up to 40-50 sec for techniques like chronopotentiometry.
Viscosity is affected by temperature where low temperatures could
essentially stop mass transport.
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Possible Roles of the Supporting Electrolyte (inert)
Regulates cell resistance and mass transport by electrical migration
Control or “buffer” the level of hydronium ion activity in solution
Associate with the electroactive analyte (i.e. complexing of ions with metals, EDTA)
Form ion-pair or micellar aggregates with the electroactive analyte (ionic
liquids)
Determines the structure of the double layer
Impose positive or negative voltage limits because of redox properties
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Control of Cell Resistance
Primary function of the supporting electrolyte is to provide a conducting
medium which will minimize iR drop and ohmic heating of the solution (i.e.
when i is large).
To have mass transport controlled only by diffusion need to use 50-100 times
the concentration of supporting electrolyte to electroactive analyte
(suppresses electrical migration).
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Advantages:
eliminates the contribution of migration to mass transfer for the electroactive species
decreases resistance of the cell
establishes the solution composition (i.e. pH, ionic strength, ligand concentration)
reduces electrical power dissipated in the cell leading to simpler instruments
eliminates sample matrix effects
helps keep double layer thin with respect to diffusion layer
established uniform ionic strength throughout the solution, while ions are being produced or consumed at the electrodes
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Disadvantages:
large concentrations can bring impurities and interferences
can give own faradic response
can react with products of electrode process
can adsorb on electrode surface and alter kinetics
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The number of electrons that cross an interface is related stoichiometrically to the chemical reaction (amount of reactant and product).
The number of electrons is measured in terms of total charge, Q.
Charge is in units of coulombs (C),
1 C = 6.24 x 1018 electrons
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Faraday’s Law
Passage of 96,485.4 C equals consumption of 1 mole of reactant or production of 1 mole of product in a one electron reaction.
Current, i, = rate of flow of coulombs (or e-’s), where 1 ampere (A) = 1 C/sec.
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A plot of the current as a function of the potential, one obtains a current-potential (i vs. E) curve.
Interpretation of i vs. E curves makes up the bulk of electrochemical research.
Two techniques: linear sweep voltammetry (LSV) and cyclic voltammetry (CV) are effective techniques for mechanistic probing of redox systems.
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For an Electrochemical cell:
Disconnecting the counter electrode and using an internal high resistance (high impendence circuit), there is no current flow between the WE and Ref.
The potential measured under these conditions is called the open-circuit potential (OCP) (zero-current potential or rest potential).
This OCP is a measure of the interfacial energy between the electrode interface and solution.
Sometimes OCP can be calculated from thermodynamic data and is a measure of the equilibrium or thermodynamic potential.
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There is a half-reaction for the Ag/AgBr electrode.
AgBr + e- Ag + Br- Eo = 0.0713 V vs NHE
but no half-reaction for the Pt electrode.
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If a power supply is connected to the cell and scanned in the negative or positive direction, then an i-V curve is generated.
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e-’s flow from
solution to electrode.
Br2 + 2e- 2Br-
e-’s flow from electrode to
protons in solution.
2H+ + 2e- H2
Near the thermodynamic
potential of –0.0711 V.
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Since the electrode (Ag/AgBr) potential remains constant all potential change is due to reactions at the Pt/solution interface.
So OCP lies somewhere between the background limits.
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If the Pt is replaced with a different WE, i.e., Hg, the cell is:
Hg/H+, Br- (1M)/AgBr/Ag
The i-V curve changes (shifts):
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Hg2Br2 + 2e- 2 Hg + 2 Br-
Hydrogen evolution
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Hydrogen evolution does not occur near the thermodynamic potential as before but is far negative.
Thermodynamics here are not changed since it is the same half reaction, but the rate is lower.
So an increase electron energy (more negative potential) is needed to make the reaction occur. This additional needed potential is called the overpotential.
Rate constant for a heterogeneous electron-transfer reaction is a function of applied potential
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If we take the same cell and add Cd2+,
Hg/H+,Br (1M), Cd2+ (10-3 M)/AgBr/Ag
There is an additional curve at ~-0.4V for reduction of Cd(Hg) amalgam.
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Typical potential sweep experiments vary the potential linearly with time.
The sweep rate, u, ranges from 10 mV/s to 1000 V/s for conventional electrodes to 106 V/s for UMEs.
Experiment is performed in an unstirred solution.
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In linear sweep voltammetry (LSV) a fixed potential range is employed much like potential step measurements. However in LSV the voltage is scanned from a lower limit to an upper limit as shown below.
The current is recorded as a function of potential.
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Begin at OCP and scan either positive or negative to find the redox couple.
Scan past the formal potential of the redox couple until the [O] is effectively zero at the electrode surface.
Past this point the system is under diffusion control (essentially chronoamperometric conditions apply and current drops with a t-1/2
dependence.
This gives a peak-like profile.
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Two important parameters for the experiment are the peak potential, Ep, and the peak current, ip.
The peak current in a reversible system is defined by the Randles-Sevcikequation.
ip = (2.69x105) n3/2 A D1/2 Co u1/2 (at 25oC)
where ip is peak current (A)
n - # of e-’s transferred
A – electrode area (cm2)
D – diffusion coefficient (cm2/s)
Co – concentration of species in the bulk (mol/cm3)
u – scan rate (V/s)
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Cyclic voltammetry (CV) is a reversal technique and very popular for initial studies of new electrochemical systems.
For the experiment, start at potential, Ei, where no electrolysis occurs and scan through potential where species in solution is either oxidized or reduced.
Then reverse the experiment.
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For the experiment, the parameters are: the initial potential, Ei
the initial sweep direction the sweep rate, u the maximum potential, Emax
the minimum potential, Emin
the final potential, Ef
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Nernstian SystemsFor a planar electrode at 25oC,
ip = (2.69 x 105) n3/2 A DO1/2 C* u1/2 (Randles-Sevcik equation)
A – cm2
DO – cm2/s
CO* - mol/cm3
u – V/s
ip – Amps
ip u1/2
Plot of ip versus u1/2 gives a straight line for reversible systems
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Nernstian Systems
ipa/ipc = 1
Eo' = (Epa + Epc)/2
Ep = Eo' – 0.029/n
DEp = Epa – Epc = 0.058/n
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Nernstian SystemsPeak current increases with scan rate, but potential where peak occurs does not
change.
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Nernstian Systems
The figure below shows a series of linear sweep voltammograms recorded at different scan rates for an electrolyte solution containing only Fe3+
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Other Systems
If the electron transfer processes were 'slow' (relative to the voltage scan rate), the reactions are referred to as quasi-reversible or irreversible electron transfer reactions.
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Irreversible System
ip = (2.99 x 105) n(ana)1/2 A DO
1/2 C* u1/2
Where a is the transfer coefficient and na is the # of e-’s in the
rate-determining step.
In this case the peak potential is no longer independent of scan
rate.
Ep = Eo' + RT/(anaF) [-0.78 + ln (ko/DO1/2) – ½ ln ((anaF)/(RT)
u)]
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Irreversible System
a is the transfer coefficient and is a kinetic term from the Butler-
Volmer model.
a typically lies between 0 and unity,
i.e. the transition state lies intermediate in its electrical behavior
between the reactants and products.
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Evaluating the CV curve over a wide scan rate range can help elucidate the kinetics:
DEp
Reversible: DEp = Epa – Epc = 0.058/n and is independent of scan rate.
Quasi- or irreversible:
DEp is dependent on scan rate.
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Background current
Peak current is measured after subtraction of background current.
Current obtained on a solution containing all ingredients except the electroactive species of interest.
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Background current
Made up of:
• residual current (oxidation or reduction of electroactive
components) (i.e. dissolved oxygen)
• charging current (non-faradaic current required to charge the
electrode to a given potential)
ic = C dE/dt = A Cdl u
Cdl – capacitance of the double layer
ic u and ip u1/2 – so ic increases more rapidly
with u.
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Nernstian Systems
For the cyclic voltammetry (CV) experiment, the linear
scan is switched directions at a certain time, t = l
(switching potential, El).
So the potential is given at any time by:
(0 < t < l) E = Ei – ut
(t > l) E = Ei – 2ut + ut
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Nernstian SystemsSwitching potential affects the shape of the overall appearance of the
curve.
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Nernstian Systems
DEp varies slightly with El
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Chem 5570
Nernstian Systems
If the Diffusion coefficients of the
reactants and products are
vastly different, then there will
be a peak shift for one of the
peaks.
Typically Dox = Dred and
chronoamperometry can be
used to measure the values.
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Nernstian Systems
pH effects
For a redox potential, with equal diffusion coefficients, there can be uptake or
loss of protons.
Ox + mH+ + ne- Red
E = Eo – RT/nF ln[Red]/([Ox][H+]m
Eoeff = Eo – 2.303 mRT/nF pH
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Nernstian SystemsScheme of Squares (Jacq)
Developed using the reduction of quinones
Assume that electron transfer is the rate limiting step and all protonations
are at equilibrium.
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Nernstian Systems
Scheme of Squares (Jacq)
The reduction route is dependent on the pKa associated with the intermediates,
the pH of the environment next to the electrode, and the formal potentials.
Horizontal movement represents electron transfer and vertical movement
represents proton transfer.
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Nernstian SystemsScheme of Squares (Jacq)
Example anthraquinone-2-sulfonate (AQMS)
Q + H+ QH+
QH+ + e- QH·
QH· + H+ QH ·+
QH ·+ + e- QH2
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Applications
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Applications
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Applications
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Applications
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Applications
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Applications
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Test 3 – Due next Tuesday November 21st
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Chem 5570