chem 5336 (introduction)
TRANSCRIPT
Introduction to Electroanalytical Chemistry
Potentiometry, Voltammetry, Amperometry, Biosensors
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
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
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
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.
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.
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
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
• 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
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
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
Polarization - theoretical
Ideally Polarized ElectrodeIdeal Non-Polarized Electrode
No oxidation or reduction
reduction
oxidation
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
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
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
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