EChemIntro_CHM5336

7/30/2019 EChemIntro_CHM5336

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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 process

E = const + k log C ( potentiometry )

Measure the Rate of a redox reaction; Current(I) = k C ( voltammetry ) Electrochemical Synthesis Organics, 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 ( E eq )

Electrolytic cells , power applied, output meas.

The Nernst Equation For a reversible process: Ox + ne- Red E = E o (2.303RT/nF) Log (a red /a ox) a (activity), related directly to concentration


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Voltammetry is a dynamicmethod

Related to rate of reaction at an electrode

O + ne = R, E o in Volts

I = kA[O] k = const. A = areaFaradaic current, caused by electron transfer

Also a non-faradaic current formspart of background current


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Electrical Double layer at Electrode

Heterogeneous system: electrode/solution

interface The Electrical Double Layer, es in electrode;ions in solution important for voltammetry: Compact inner layer: d o to d 1, E decreases linearly.

Diffuse layer: d 1 to d 2, E decreases exponentially.


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Electrolysis: Faradaic and Non-FaradaicCurrents

Two types of processes at electrode/solutioninterface that produce current Direct transfer of electrons, oxidation or reduction

Faradaic Processes . Chemical reaction rate atelectrode 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 thebulk 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-electrodeVoltammetry cell

Counter electrode

Reference electrode

Working electrode

End of Working electrode

O

R

O

R

e -

Bulk solution

Mass transport

Reduction at electrodeCauses current flow inExternal circuit


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Analytical Electrolytic Cells

Use external potential (voltage) to drivereaction

Applied potential controls electron energy As E o gets more negative, need more

energetic electrons in order to causereduction. For a reversible reaction: E applied is more negative than E o, reduction

will occur if Eapplied is more positive than E o, oxidation

will occur

O + 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 E applied so that desired reaction occurs Stir solution Measure Current

Voltammetry Quiet or stirred solution Vary (scan) E applied 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

N2inlet

reference

insulator electrodematerial

Cell for voltammetry, measures I vs. E wire

Output, I vs. E, quiet solution

reduction


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Polarization - theoretical

Ideally Polarized Electrode Ideal 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

= E Eeq; can be zero or finite E < E eq < 0 Amt. of potential in excess of E eq needed to makea non-reversible reaction happen, for example

reduction

E eq


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NERNST Equation : Fundamental Equationfor reversible electron transfer at electrodes

O + ne - = R, E o in VoltsE.g., Fe 3+ + e- = Fe 2+

If in a cell, I = 0, then E = E eq All equilibrium electrochemical reactions obey the

Nernst EquationReversibility means that O and R are at equilibrium at all times, not allElectrochemical reactions are reversible

E = Eo

- [RT/nF] ln (a R/a O) ; a = activitya R = f RCR a o = f oCo f = activity coefficient, depends on ionic strength

Then E = E o - [RT/nF] ln (f R/f O) - [RT/nF] ln (C R/C O)F = Faraday const., 96,500 coul/e, R = gas const.T = absolute temperature


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Ionic strength I = z i2m i, Z = charge on ion, m = concentration of ion

Debye Huckel theory says log f R = 0.5 z i2 I1/2

So f R/f Owill be constant at constant I.

And so, below are more usable forms of Nernst Eqn.

E = E o - const. - [RT/nF] ln (C R/C O)Or

E = Eo

- [RT/nF] ln (C R/C O); Eo

= formal potential of O/R

At 25 oC using base 10 logs

E = E o - [0.0592/n] log (C R/CO); equil. systems

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EChemIntro_CHM5336