Chem 5336 (Introduction)
Transcript of Chem 5336 (Introduction)
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Introduction to Electroanalytical Chemistry
Potentiometry, Voltammetry, Amperometry, Biosensors
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Applications• Study Redox Chemistry
– electron transfer reactions, oxidation, reduction, organics & inorganics, proteins
– Adsorption of species at interfaces• Electrochemical analysis
– Measure the Potential of reaction or processE = const + k log C (potentiometry)
– Measure the Rate of a redox reaction; Current (I) = k C (voltammetry)
• Electrochemical SynthesisOrganics, inorganics, materials, polymers
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Electrochemical Cells• Galvanic Cells and Electrolytic Cells• Galvanic Cells – power output; batteries • Potentiometric cells (I=0) read Chapter 2
– measure potential for analyte to react– current = 0 (reaction is not allowed to occur)– Equil. Voltage is measured (Eeq)
• Electrolytic cells, power applied, output meas.– The Nernst Equation
• For a reversible process: Ox + ne- → Red• E = Eo – (2.303RT/nF) Log (ared/aox)• a (activity), related directly to concentration
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Voltammetry is a dynamic method
Related to rate of reaction at an electrode
O + ne = R, Eo in Volts
I = kA[O] k = const. A = areaFaradaic current, caused by electron transfer
Also a non-faradaic current forms part of background current
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Electrical Double layer at Electrode
• Heterogeneous system: electrode/solution interface• The Electrical Double Layer, e’s in electrode; ions in
solution – important for voltammetry:– Compact inner layer: do to d1, E decreases linearly.
– Diffuse layer: d1 to d2, E decreases exponentially.
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Electrolysis: Faradaic and Non-Faradaic Currents
• Two types of processes at electrode/solution interface that produce current– Direct transfer of electrons, oxidation or reduction
• Faradaic Processes. Chemical reaction rate at electrode proportional to the Faradaic current.
– Nonfaradaic current: due to change in double layer when E is changed; not useful for analysis
• Mass Transport: continuously brings reactant from the bulk of solution to electrode surface to be oxidized or reduced (Faradaic)– Convection: stirring or flowing solution– Migration: electrostatic attraction of ion to electrode– Diffusion: due to concentration gradient.
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Typical 3-electrode Voltammetry cell
Counterelectrode
Reference electrode
Working electrode
End of Working electrode
O
R
O
Re-
Bulk solution
Mass transport
Reduction at electrodeCauses current flow inExternal circuit
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Analytical Electrolytic Cells
• Use external potential (voltage) to drive reaction
• Applied potential controls electron energy• As Eo gets more negative, need more
energetic electrons in order to cause reduction. For a reversible reaction: Eapplied is more negative than Eo, reduction
will occur if Eapplied is more positive than Eo, oxidation
will occurO + ne- = R Eo,V electrode reaction
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• Current Flows in electrolytic cells– Due to Oxidation or reduction– Electrons transferred– Measured current (proportional to reaction
rate, concentration)
• Where does the reaction take place?– On electrode surface, soln. interface – NOT in bulk solution
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Analytical Applications of Electrolytic Cells
• Amperometry– Set Eapplied so that desired reaction occurs– Stir solution– Measure Current
• Voltammetry– Quiet or stirred solution– Vary (“scan”) Eapplied
– Measure Current• Indicates reaction rate• Reaction at electrode surface produces concentration gradient
with bulk solution• Mass transport brings unreacted species to electrode surface
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E, V
time
Input: E-t waveform
potentiostat
Electrochemical cell
counter
working electrode
N2
inlet
reference
insulator electrodematerial
Cell for voltammetry, measures I vs. Ewire
Output, I vs. E, quiet solution
reduction
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Polarization - theoretical
Ideally Polarized ElectrodeIdeal Non-Polarized Electrode
No oxidation or reduction
reduction
oxidation
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Possible STEPS in electron transfer processes
Rate limiting step may be mass transfer
Rate limiting step may be chemical reaction
Adsorption, desorption or crystallization polarization
Charge-transfer may be rate limiting
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Overvoltage or Overpotential η
• η = E – Eeq; can be zero or finite
– E < Eeq η < 0
– Amt. of potential in excess of Eeq needed to make
a non-reversible reaction happen, for example
reduction
Eeq
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NERNST Equation: Fundamental Equation for reversible electron transfer at electrodes
O + ne- = R, Eo in Volts•E.g., Fe3+ + e- = Fe2+
If in a cell, I = 0, then E = Eeq
All equilibrium electrochemical reactions obey the Nernst Equation
Reversibility means that O and R are at equilibrium at all times, not all Electrochemical reactions are reversible
E = Eo - [RT/nF] ln (aR/aO) ; a = activity
aR = fRCR ao = foCo f = activity coefficient, depends on ionic strength
Then E = Eo - [RT/nF] ln (fR/fO) - [RT/nF] ln (CR/CO)
F = Faraday const., 96,500 coul/e, R = gas const.T = absolute temperature
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Ionic strength I = Σ zi2mi,
Z = charge on ion, m = concentration of ion
Debye Huckel theory says log fR = 0.5 zi2 I1/2
So fR/fOwill be constant at constant I.
And so, below are more usable forms of Nernst Eqn.
E = Eo - const. - [RT/nF] ln (CR/CO)
OrE = Eo’
- [RT/nF] ln (CR/CO); Eo’ = formal potential of O/R
At 25 oC using base 10 logs
E = Eo’ - [0.0592/n] log (CR/CO); equil. systems