Yonsei University Fundamentals of Electrochemistry

References

Electrochemical Methods : Fundamentals and Applications, 2nd edition. by A. J. Bard and L.R. Faulkner

• Makes use of electrochemistry for the purpose of analysis

• A voltage (potentiometry) or current () signal originating from an electrochemical cell is related to the activity or concentration of a particular species in the cell.

• Excellent detection limit (10-8 ~ 10-3 M): 1959, Nobel Prize ()

• Inexpensive technique.

• Easily miniaturized : implantable and/or portable (biosensor, biochip) Electrochemical cells

Galvanic cell

Digital High input impedance voltmeter A B e- Anode reaction Zn Zn2+ + 2e- : oxidation e- Cathode reaction Cu2+ + 2e- Cu (-)KCl (+) : reduction Zn Cu Salt bridge Zn + Cu2+ Zn2+ + Cu

K+ - - 2e Cl 2e- Cell potential : a measure of difference in electron Zn2+ energy between the two 2- SO4 Zn Cu Open-circuit potential (zero-current potential) Zn2+ Cu2+ 2+ 2+ Zn 2- 2- Cu : can be calculated from thermodynamic data, ie. SO4 SO 4 standard cell potentials of the half-cell reactions. Anode Cathode

Fig. 27.1 Electrochemical cell consisting of a zinc in 0.1 M ZnSO4, a copper electrode in 0.1 M

CuSO4, and a salt bridge. . (From Heineman book) Standard Electrode Potential

Table 22.1 Standard Electrode Potentials

Reaction E0 at 25 ℃, V - - Cl2(g) + 2e  2Cl +1.359 + - O2(g) + 4H +4e  2H2O +1.229 - - Br2(aq) + 2e  2Br +1.087 - - Br2(l) + 2e  2Br +1.065 Ag+ + e-  Ag(s) +-.799 Reduction 자발적 Fe3+ + e-  Fe2+ +0.771 - - - I3 + 2e  3I +0.536 Cu2+ + 2e-  Cu(s) +0.337 - - Hg2Cl2(s) + 2e  2Hg(l) + 2Cl +0.268 AgCl(s) + e-  Ag(s) + Cl- +0.222 3- - 2- A quantitative description of the relative driving force Ag(S2O3)2 + e  Ag(s) + 2S2O3 +0.010 for a half-cell reaction. + - SHE 2H + 2e  H2(g) 0.000 A relative quantity vs standard hydrogen electron AgI(s) + e-  Ag(s) + I- -0.151 assigned to zero volt. E0(SHE)=0 - 2- PbSO4(s) + 2e  Pb(S) + SO4 -0.350 Oxidation Cd2+ + 2e-  Cd(s) -0.403 Fig. 22.5 Definition of the standard electrode 자발적 Zn2+ + 2e-  Zn(s) -0.763 potential for M2+(aq) + 2e- M(s). Nernst Equation (activities of all species = 1)

Le Chatelier’s principle: increasing reactant concentrations drives the reaction to the right

The net driving force of the reaction is expressed by the Nernst equation

The Nernst equation tells us

the potential of a cell whose reagents are not all unit activity

Nernst Equation for a Half-Reaction

R: gas constant = 8.314 J/Kmol - aA + ne bB T: temperature (K)

RT Ab o B ------(14.13) E  E  ln a nF AA

∆G = ∆Go + RT lnQ (Q; reaction quotient) -nFE = -nFEo + RT lnQ (양변을 nF 로 나누어 준다) E = Eo –(RT/nF) lnQ Eo and Equilibrium Constant

• A galvanic cell produces electricity because the cell reaction is not at equilibrium • The potentiometer allows negligible current to flow The concentration in each half-cell remains unchanged

0.412 V At the outset

[Cu2+] 증가 [Ag+] 감소

0 V At equilibrium Liquid Junction Potential (Elj)

Elj (diffusion potential) Develops at the interface between two liquids as a result of differences in the rates with which ions move from one liquid to the other. Type 1 Type 2 Type 3

0.01 M 0.1 M 0.1 M 0.1 M 0.1 M 0.05 M HCl HCl HCl KCl HCl KNO3 H+ H+ H+ Cl- - + Cl K K+ - NO3 + - - + - + Fig. 27.4 Types of liquid junction. Arrows show the direction of net transfer for each ion, and their lengths indicate relative mobilities. The polarity of the junction potential is indicated in each case by the circled signs.

Table 27.4 Liquid Junction Potentials of 0.1 M Concentrations of Electrolytes

Elj Elj Elj Elj Junction observed Junction observed Junction observed Junction observed (mV) (mV) (mV) (mV)

HCl : KCl 26.78 HCl : NH4Cl 28.40 KCl : NH4Cl 2.16 NaCl : NH4Cl -4.21

HCl : NaCl 33.09 KCl : LiCl 8.79 NaCl : LiCl 2.62 LiCl : NH4Cl -6.93 KCl : LiCl 34.86 KCl : NaCl 6.42 Cell Potential

Ecell = Eright –Eleft + Elj -EiR

Eright = Ecathod Eleft = Eanode Elj = liquid junction potential

EiR = Ohmic loss (iR drop) : The potential of an EC cell when electrolysis occurring is diminished by resistance of the cell to current. Electrochemical cells

Electrolytic cell

(-) Power (+) - e Supply

e-

(-)KCl (+) Anode reaction 2+ - Zn Cu : oxidation Cu Cu + 2e Salt bridge Cathode reaction Zn2+ + 2e- Zn : reduction K+ Cl- - 2e 2e- Cu + Zn2+ Cu2+ + Zn

2- 2+ SO4 Cu Zn Cu Zn2+ Cu2+ 2+ 2+ Zn 2- 2- Cu SO4 SO4 Cathode Anode

Fig. 27.1 Electrochemical cell consisting of a zinc electrode in 0.1 M ZnSO4, a copper electrode in 0.1 M

CuSO4, and a salt bridge. Electrolytic cell. Reduction and oxidation process

We observe or control the potential of the with respect to the reference. → Controlling the energy of the electrons within the working electrode.

Reduction A + e- → A- Oxidation A - e- → A+ Electrode Solution Electrode Solution e

Vacant MO Vacant - (LUMO) - MO

Energy level Potential Potential of electrons e Energy level of electrons + Occupied MO + Occupied (HOMO) MO

The critical potentials at which these processes occur are related to the Standard Potentials, E0, for the specific chemical substances in the system. Determination of HOMO-LUMO Band Gap Energy

- Use of (CV) or differential pulse voltammetry (DPV) - UV absorption spectrum - Photoluminescence spectrum Determination of HOMO-LUMO Band Gap Energy

3.18 eV

3.59 eV Definition of words (What is current?)

Faraday’s law

relate the amount of electrical charge passed through an electrochemical cell to the quantity of material that has undergone electrolysis Q = n F N (1.3.1) F : faraday constant (96,485 C/mol) N : the number of moles electrolyzed n : the number of electrons involved in the electrode or redox reaction dQ dN i = = nF (1.3.3) rearrangement dt dt of 1.3.3 dN i Rate (mol/s) = = dt nF Potentiometry

Measurement of cell voltages

Chemical information (activity) Potentiometry

Measurement of the difference in potential between the two electrodes of a galvanic cell under the condition of zero current are described by the term potentiometry Equilibrium Method Accurate measurements of (a) activities or concentration (b) free-energy change and equilibrium constants of many solution reactions

pH/mV meter Reference Indicator Electrode Electrode Ecell = Eind –Eref + Elj

The indicator electrode is chosen Sample or so that its half-cell potential responds to standard the activity of a particular species in solution whose activity or concentration

Magnetic is to be measured Stirring bar Fig. 28.1 Schematic diagram of apparatus for potentiometry. Magnetic Stirrer Electrode and Potentiometry

Reference Electrodes : provides a constant potential

Votmeter : Ag + Cl- → AgCl(s) + e-

Salt bridge Anode : E-° = 0.222 V KCl Indicator Electrode (Pt) : Fe3+ + e- → Fe2+ Ag

Cathode : E+° = 0.771 V

2+ 3+ E+ = 0.771 - 0.0592 log[Fe ]/[Fe ] AgCl Saturated Pt - KCl solution E- = 0.222 – 0.0592 log[Cl ] variable 2+ 3+ Fe3+, Fe2+ E = E+ -E- = (0.771-0.0592 log[Fe ]/[Fe ]) Solid KCl Measured - Anode Cathode – (0.222 – 0.0592 log[Cl ]) - - 3+ - 2+ voltage Ag + Cl ↔ AgCl +e Fe +e ↔ Fe Constant ; depend on KCl solubility 3+ - 2+ Ag(s)/AgCl(s)/Cl-(aq)//Fe +e ↔ Fe (aq)/Pt(s) saturated solution Fig. 15.1 A galvanic cell that can be used to measure the quotient [Fe2+]/[Fe3+] in the right half-cell. The Pt wire is the indicator electrode, and the entire left half-cell plus salt bridge (enclosed by the dash line) can be considered to be a reference electrode. Reference Electrode: Ag/AgCl

Reference Electrode : Ag | AgCl | Cl− + Salt Bridge → Constant potential

V

Anode Cathode

Pt : indicator electrode

Ag/AgCl/Cl- plus Salt bridge reference electrode Fe2+, Fe3+

Fig. Another view of Figure 15.1. The contents of the colored box in Figure 15.1 are now considered to be a reference electrode dipped into the analyte solution. Ion Selective Electrode (ISE)

Ca2+-binding ligand (ionophore) soluble in membrane + -

0.01 M Ca2+ 0.1 M Ca2+ + - (0.01+)M Ca2+ (0.1-)M Ca2+ 0.02 M Cl- 0.2 M Cl- + - 0.02 M Cl- 0.2 M Cl- + - Ca2+ Ca2+ + - + - Aqueous Aqueous solution Membrane solution Low conc. + Membrane - High conc. (a) (b)

Fig. Mechanism of ion-selective electrode. (a) Initial conditions prior to Ca2+ migration across the membrane. (b) After  moles of Ca2+ per liter have crossed the membrane, giving the left side a charge of +2 mol/L and the right side a charge of -2 mol/L. Ion Selective Electrode (ISE)

Concentration difference → free energy difference : membrane potential

A1 ∆G = -RT ln A2 ∆G = -nFE

A1 -RT ln = -nFE A2 A1 0.05916 A1 E = RT/nF ln = log (volts at 25℃) memb A2 n A2

Ememb: membrane potential 1: Sample solution, 2: Internal solution

Ca2+ → n = 2 10-fold change in [Ca2+] → 59.16/2 = 29.58 mV change Ion Selective Electrode (ISE)

Ecell = Eref·ext -Eref·int + Ememb + Elj RT E = E -E + RT ln 1 + E + ln (a ) sample cell ref·ext ref·int lj ZF i ZF (ai)int Constant mV meter RT E = K + ln (a ) sample cell ZF i 0.05916 E = K + log (a ) sample cell n i Internal reference External (n = Z) at 25 ℃ electrode reference electrode

Internal Solution Liquid junction

(ai) internal potential Potential(V) ai sample = constant log [ai] Selectivity Coefficients in ISE

Since membrains respond to a certain degree to ions other than the analyte (i.e. interferents) A more general expression: RT E = k + ln(a + k a z/x) cell ZF i ij j

ai = activity of analyte ion. i

aj = activity of interferent ion. j x = charge of interferent ion

kij = selectivity constant

kij = response to j/response to i For Na+, K+ = 1/2800

Small values of kij are characteristics of electrode with good selectivity for the analyte, i Ionophores (neutral carriers) for ISE

antibiotic Ionophores (neutral carriers) for ISE Selectivity of Li – ion selective electrode

0 Li+

-1 K+ Cs+ -2 Na+ + n+ Log kLi , M Rb+ + -3 NH4 H+ 2+ -4 Sr Ba2+ Cs2+ -5 Mg2+

pH Electrode (Glass-membrane electrode)

Glasses of certain compositions respond to pH due to a membrane potential generated by an ion-exchange mechanism with H+

(+) Lead to pH meter (-)

Aqueous solution Fig. 15.9 Diagram of a glass combination Ag wire saturated with electrode having a silver-silver AgCl and KCl chloride reference electrode. The glass electrode is immersed in Porous plug to allow slow a solution of unknown pH so that drainage of electrolyte out the porous plug on the lower right of electrode is below the surface of the liquid. The two silver electrodes measure the voltage across the glass 0.1 M HCl membrane. Glass Saturated with membrane AgCl

Glass : 22% Na2O, 6% CaO, and 72% SiO2 (Corning 015 glass membrane) Gas-Sensing Electrode : (compound electrode)

Ag/AgCl external reference electrode pH electrode

pH-Sensitive glass membrane

H+ ISE Thin-solution layer + - CO2 + H2O ↔ HCO3 Thin- solution layer Gas-permeable membrane

CO2 Sample (a) (b) Gas-permeable membrane Fig. 28.8 Schematic representation of (a) gas-sensing electrode and (b) membrane and thin layer of electrolyte region of CO2 electrode. Gas-Sensing Electrode : (compound electrode)

Table 28.3 Gas-Sensing Electrodes

Gas Electrode Internal Solution Equilibrium Sensing Element

+ CO2 CO2 + H2O ↔ HCO3 + H Glass, pH + - NH3 NH3 + H2O ↔ NH4 + OH Glass, pH + - HCN HCN ↔ H + CN Ag2S, pCN + - HF HF ↔ H + F LaF3, pF + 2- H2S H2S ↔ 2H + S Ag2S, pS - + SO2 SO2 + H2O ↔ HSO3 + H Glass, pH Bio-catalytic Membrane Electrode (potentiometric biosensor)

+ NH4 - selective ISE Ion-selective or gas-permeable membrane

Enzyme (urease) + 2- Biocatalyst Urea s P 2NH4 + CO3 layer

Semi-permeable membrane

s

Fig. 28.9 Schematic diagram of biocatalytic electrode. Bio-catalytic Membrane Electrode (potentiometric biosensor)

Table 28.4 Representative Biocatalytic Membrane Electrode i-STAT Co. (Princeton, NJ) Microfluidic and Biosensor Chip Technology (Multibiosensor)

“Chem 7” test: Sodium, Potassium, Chloride, Ionized Calcium, pH and PCO + - 2 Na+, K , Cl , total CO2, by ion-selective electrode potentiometry. glucose, urea, creatinine Urea is first hydrolyzed to ammonium ions in a reaction catalyzed by the enzyme urease. The ammonium ions are measured by an Cartridge label ion-selective electrode. Sample entry well gasket Glucose is measured amperometrically. Fluid channel PO2 is measured amperometrically. Cartridge cover

Hematocrit is determined conductometrically. Sample entry well

HCO3, TCO2 , BE, sO2, Anion Gap and Hemoglobin. Tape gasket

Biosensor chips

Calibrant pouch

Puncturing barb Cartridge base Air bladder Voltammetry

Control of cell voltages

Measurement of current Electrochemical cells

Microelectrodes for working electrodes

drop diameter : 0.1 ~ 1.0 mm commercially available: can be operated A new drop forms and breaks as a dropping mercury electrode or hanging every 2 ~ 6 s mercury drop electrode Commercial

Potentiostat vs Galvanostat

A is used to apply a controlled potential to an electrochemical cell and measured the current response. A galvanostat is used to apply a controlled current.

EG&G PAR 273A BAS 100B/W Electrochemical Cells for Voltammetric Sensors

Control E – measure i voltammetry – scan E amperometry – hold E

Au Pt carbon 3 electrode BAS, Inc. Adapted from Sawyer, Heineman & Beebe, Chemistry Experiments for Instrumental Methods, Wiley I. Papautksy, University of Cincinnati Biosensor Design Fabrication Process

connection pad Screen printing technology for thick-film electrode RE WE2 A Outermembrane layer

WE2 WE1 Non-active enzyme paste layer A’ polyester substrate insulation WE1 non-enzyme paste Active enzymepastelayer enzyme paste RE outer membrane Ag/AgCl layer AgCl

A ----- Dielectriclayer A’ Carbon base electrode carbon Ag conductor Silverconductionlayer Polyester substrate Electrochemical cells

Reference electrode (ideal non-polarized electrode)

E - I curves for reference and working electrode reference electrode working electrode (ideal polarizable electrode) Current A Current

B

Electrode potential A Electrode potential B Schematic diagram of the electrochemical cell

Power Supply Reference electrode i V AgBr + e Ag + Br-

(E0 = 0.0713 V vs NHE) Ag

Pt AgBr Fig. 1.1.3 Schematic diagram of the electrochemical cell Pt/HBr(1 M)/AgBr/Ag attached to power supply and meters for 1 M HBr obtaining a current-potential(i-E) curve. Schematic Current-Potential Curve

Use of Pt electrode

Negative direction scan H+ reduction at Pt electrode + 2H + 2e → H2 (E0 = 0 V vs NHE or -0.07 V vs Ag/AgBr) Oxidation at reference electrode Ag + Br- → Ag+Br- + e- (Br- concentration no change→ no potential change)

Positive direction scan Br- oxidation at Pt electrode - - 2Br  Br2 + 2e Fig. 1.1.4 Schematic current-potential curve for the cell Pt/K+, Br- (E0 = +1.09 V vs NHE or +1.02 V vs Ag/AgBr) (1 M)/AgBr/Ag, showing the limiting proton reduction and bromide Reduction at reference electrode oxidation processes. The cell potential is given for the Pt electrode with respect to the Ag electrode, so it is equivalent to Ept(V vs AgBr). + - - - Ag Br + e → Ag + Br Since EAg/AgBr = 0.07 V vs NHE, the potential axis could be converted

to Ept(V vs NHE) by adding 0.07 V to each value of potential. Schematic Current-Potential Curve

Use of Hg electrode

Negative direction scan The heterogeneous rate constant for

H2 evolution at Hg electrode is much lower than at Pt. The heterogeneous rate constant is a function of applied potential. The additional potential (beyond the thermo- dynamic requirement) needed to drive a reaction at a certain rate is called overpotential :  = Eapp -Eeq Fig. 1.1.5 Schematic current-potential curve for the Hg Positive direction scan electrode in the cell Hg/H+, Br-(1 M)/AgBr/Ag, showing Mercury oxidation the limiting process : proton reduction with a large Hg Br + 2e → 2Hg + 2Br- overpotential and mercury oxidation. The potential 2 2 axis is defined through the process outlined in the (E0 = +0.14 V vs NHE or +0.07 V vs Ag/AgBr) caption to Figure 1.1.4 Overpotential

+ - H3O + e → ½ H2 + H2O The reaction begins -0.35 V at Pt -0.8 V at Ag

Energy of electron + Added to H3O + H + H H O H O H Decreased H

Free energy activation energy Activation energy for electron transfer Over Energy of electron in metal Free energy electrode after overpotential potential is applied

Original energy of electron in metal electrode

+ Fig. 17.5 Schematic energy profile for electron transfer from a metal to H3O (a) with no applied potential; (b) after a potential is applied to the electrode. The overpotential increases the energy of the electrons in the electrode. Overpotential

Table 17.1 Overpotential(V) for gas evolution at various current densities at 25 ℃

10 A/m2 100 A/m2 1000 A/m2 10000 A/m2 Electrode H2 O2 H2 O2 H2 O2 H2 O2 Platinized Pt 0.0154 0.398 0.0300 0.521 0.0405 0.638 0.0483 0.766 Smooth Pt 0.024 0.721 0.068 0.85 0.288 1.28 0.676 1.49 Cu 0.479 0.422 0.584 0.580 0.801 0.660 1.254 0.793 Ag 0.4751 0.580 0.7618 0.729 0.8749 0.984 1.0890 1.131 Au 0.241 0.673 0.390 0.963 0.588 1.244 0.798 1.63 Graphite 0.5995 0.7788 0.9774 1.2200 Sn 0.8561 1.0767 1.2230 1.2306 Pb 0.52 1.090 1.179 1.262 Zn 0.716 0.746 1.064 1.229 Cd 0.981 1.134 1.216 1.254 Hg 0.9 1.0 1.1 1.1 Fe 0.4036 0.5571 0.8184 1.2915 Ni 0.563 0.353 0.747 0.519 1.048 0.726 1.241 0.853

Higher overpotential is required for higher current density (From D. C. Harris Book) Potential Ranges

1 M H2SO4 (Pt) Pt pH 7 buffer (Pt) 1 M NaOH (Pt)

1 M H2SO4 (Hg) Hg 1 M KCl (Hg) 1 M NaOH (Hg) High overpotential 0.1 M Et4NOH (Hg) for H2 evolution 1 M HClO (C) C 4 0.1 M KCl (C) - H2O + e → 1/2H2+OH +3 +2 +1 0 -1 -2 -3

Positive potential limitation Negative potential limitation

-H2O oxidation -H2O reduction

-O2 evolution -H2 evolution

Fig.E.2. Potential ranges for three types of electrodes in various supporting electrolytes. (Adapted from A. J. Bard and L. R. Faukner, Electrochemical Methods, back cover. New York:Wiley 2001) Schematic Current-Potential Curve

Use of Hg electrode

Addition of a small amount of Cd2+

2- Hg - CdBr4 + 2e → Cd(Hg) + 4Br Cd2+ cadmium amalgam E0 = -0.35 V vs NHE

Fig. 1.1.6 Schematic current-potential curve for the Hg electrode in the cell Hg/H+, Br-(1 M), Cd2+(10-3 M)/AgBr/Ag, showing reduction wave for Cd2+. Predictions for possible Reduction or oxidation

Prediction for possible reduction or oxidation based on thermo- dynamic consideration (E0)

Possible - Possible - reduction reduction E0 reaction 0 reaction E (V vs NHE) -0.76 Zn2+ + 2e → Zn (V vs NHE) Oxidation Approximate potential 0 for zero current (Au) -0.25 Ni2+ + 2e → Ni -0.41 Cr3+ + e → Cr2+ +0.15 Sn2+ → Sn4+ + 2e 0 2H+ + 2e → H 2 Kinetically slow +0.15 4+ 2+ Reduction Sn + 2e → Sn +0.54 - + 2I + → I2 + 2e 0 2H + 2e → H2

+0.77 Fe2+ → Fe3+ + e (Hg ) Reduction +0.77 Fe3+ + e → Fe2+ Approximate potential (Pt) for zero current 0 + + E Approximate potential +1.23 2H2O → O2 + 4H + 4e for zero current (V vs NHE) +1.50 Au → Au3+ +3e (a) (b) (c)

Fig. 1.1.7 (a) Potentials for possible reductions at a platinum electrode, initially at ~1 V vs NHE in a solution of 0.01 M each of Fe3+, Sn4+, and Ni2+ in 1 M HCl. (b) Potentials for possible oxidation reactions at a gold electrode, initially at ~ +0.1 V vs. NHE in a solution pf 0.01 M each of Sn2+ and Fe2+ in 1 M HI. (c) Potentials for possible reductions at mercury electrode in 0.01 M Cr3+ and Zn2+ in 1 M HCl. Faradaic and Nonfaradaic Processes

Faradaic Process

Charge (e.g. electrons) are transferred across the metal-solution interface Electron transfer causes oxidation or reduction to occur. Such reactions are governed by Faraday’s Law. The amount of chemical reaction caused by the flow of current is proportional to the amount of electricity passed.

Non-faradaic Process

Adsorption and desorption The electrode-solution interface can change with changing potential or solution composition. Although charge does not cross the interface, external currents can flow (at least transiently) Electrode – Solution Interface (electrical double layer)

Water molecule

+ Three models of electrical - + double layer + A - A Helmholtz model (1879) - + + Bulk Diffuse double layer model solution (Guoy-Chapman) model - + A + Stern model - A Specifically IHP OHP adsorbed anion Compact layer Diffusion layer Linear decay

Exponential decay Potential (mV) Potential (mV)

Compact layer Diffusion layer X (Å) Simplified Model

Faradaic current Charging current

2+ Cd + - - + Cd - V C e- + - - + + +

M S - + Capacitor i Charging current - + = capacitance current - + V C (double layer capacitance) c - + d - very small charge - + - c/cm2, 10 ~ 100 μF/cm2

-qM = qs Cd = q/E 0t Voltage (Potential) Step

q = Cd · EC q dq E = ER + EC = iRs + ( i = 이므로 ) Cd dt Rs Cd dq E q = (1) dt Rs RsCd E Initially, capacitor is uncharged q = 0 at t = 0 i

-t / RsCd (2) q = E Cd [ 1 – e ]

Fig. 1.2.6 Potential step experiment for an RC circuit. E i = e-t / RsCd Rs Voltage (Potential) Step

For a potential step input, there is an exponentially decaying current having a time

constant,  = Rs ·Cd. • Double-layer charging current drops to 37 % of its initial value at t = , 5 % at t = 3.

E/Rs Resultant ( i) if Rs = 1 Ω, Cd = 20 μF i

0.37E/Rs  = 20 μs

 = R C t s d at 60 μs 95 % complete

E Applied (E) E

Fig. 1.2.7 Current transient (I vs. t) resulting from a potential step experiment . t Voltage Ramp (or potential sweep)

Linear Potential Sweep

E  (in V/s)

Current rises from zero to a steady-state Applied E(t) Value, Cd

t E = t (a) q i Resultant ( i) E = iRs + 에서 Cd dq q t = Rs + Cd dt Cd If q = 0 at t = 0, t (b) i = Cd [1-exp(-t /RsCd)] Fig. 1.2.10 Current –time behavior resulting from a linear potential sweep applied to an RC circuit . Voltage Ramp (or potential sweep)

Cyclic Potential Sweep

From CV, we can get Cd

전극 면적에 의존

Charging current Cd =  (scan rate) Fig. 1.2.11 Current –time and current-potential plots resulting from a cyclic linear potential sweep (or triangular wave) applied to an RC circuit. Factors affecting electrode reaction rate and current

Heterogeneous at the electrode/electrolyte interface

Rate depends on… Mass transfer to the electrode Surface effects (or reaction) : adsorption / desorption, crystallization Kinetic variables (electron transfer) current density (i/cm2) Considering i j -1 -2 electrode area Rate (mol·s ·cm ) = = nFA nF

Chemical rxn O` O O Electrode surf bulk O`ads Electrode Bulk solution ne- surface region

R`ads

R` Rsurf Rbulk Mechanisms of Mass Transport

The electrochemical reduction/oxidation of a molecule occurs at the electrode- solution interface. A molecule dissolved in solution in an electrochemical cell must be transported to the electrode surface for the electrochemical event to occur. The transport of molecules from regions in the solution phase to the electrode surface is an important aspect of many electrochemical techniques. The movement of material from one place to another is generally termed mass transport or mass transfer.

Hydrodynamic movement (convection) Transport of dissolved chemical species by physical movement of the solution or the electrode relative to each other. stirring the solution rotating the electrode flowing solution across the electrode by placing it in a tube through which solution is pumped. Mechanisms of Mass Transport

Diffusion A process of spatial movement of ions or molecules due to their kinetic motion (random Brownian motion) Diffusion acts to remove a concentration gradient by the mass transport of molecules from a region of high concentration to a region of low concentration. Diffusion increases the entropy of a system.

Concentration profile

2 C -x Cx,t = 0 exp (4πDt)1/2 4Dt

Cx,t : concentration of solute molecules at x, t (mol·cm-3)

C0 : concentration of solute molecules at initial impulse x : distance from center of impulse (cm) D : diffusion coefficient (cm2/s) Mechanisms of Mass Transport

Diffusion Distance

σ = 2Dt

Migration Movement of charged particles under the influence of electric field. (e.g) positively charged electrode attracts a negatively charged solute species, but repels a positively charged solute species. Potential excitation signal used in voltammetry

Linear scan Differential pulse Square wave Triangular Polarography linear Differential pulse Square-wave Cyclic -scan voltammetry Polarography voltammetry voltammetry E E E E

Time → Time → Time → Time →

Pulse Staircase Constant LCEC Amperometry Pulse voltammetry Staircase voltammetry Deposition in ASV E E E controlled E

Time → Time → Time → Cyclic Voltammtry(CV) : no stirring

CV: very popular technique for initial electrochemical studies of new systems Figure.excitation Cyclicvoltammetric Potential, V vs SCE Figure 25-20. 0 signal used toobtain scan +0.8 Forward +0.8 : 1.0 V/20 s=50mV/s Scan rate = V/s -0.2 20 Cyclic Voltammtry(CV) (reverse scan) scan Backward Time, s s Time, 40 ed voltammogram in 60 E i pa pc 80 : anodicpeakcurrent, i : cathodic peak potential, E : cathodicpeakpotential, Figure 25.20Cyclic voltammogram for a and 1.0 M in KNO and 1.0 Min solution that is6.0 mM inK pc : cathodic peak current pa : anodicpeakpotential 3 . 3 Fe(CN) 6 Cyclic Voltammtry(CV)

Nernst Equation for a Reversible System at 25 ℃ At point C, E=Eo’ 0.0591 Cs(Fe2+) Cs(Fe2+) E = Eo’ - log =1 Fe3+, Fe2+ n Cs(Fe3+) Cs(Fe3+)

Peak current: 최대 concentration gradient

Re-oxidation 과정

Figure (a) Cyclic voltammogram of 6 mM K3Fe(CN)6 in 1 M KNO3. Scan initiated at 0.8 V versus SCE in negative direction at 50 mV/s. Platinum electrode, A=2.54 mm2. (b) Concentration-distance(C-x) profiles a-k keyed to voltammogram. Cyclic Voltammtry(CV)

Peak current for a reversible system the working electrode (Randles-Sevcik equation)

5 3/2 1/2 1/2 ip = (2.69x10 ) n AD C*v

A : electrode area, D : diffusion coefficient C* : bulk concentration (mol ·cm-3) v : scan rate (V · s-1)

ip 1/2 ip ∞ v from slope D can be calculated v1/2 Cyclic Voltammtry(CV): Scan Rate Effect

The slope of the concentration distance profile at the electrode surface Co Co – Cos i = nFADo  x x=0, t = nFADo  Diffusion layer distance

Figure 30.3 (a) Effect of variation of scan rate on cyclic voltammograms and (b) plot of ip versus v1/2. Cyclic Voltammetry for Reversible System

For reversible redox couple: o E ’ = (Epa + Epc) / 2 : formal potential

실제로 switching potential(E) - Peak separation ∆Ep = Epa-Epc = 59 mV/n 에 따라 약간씩 변한다.

the number of electrons transferred in the electrode reaction(n) for a reversible couple can be determined from the separation between peak potential at 25 ℃. 3- 4- For Fe(CN)6 → Fe(CN)6 ∆Ep = 0.059 V (n=1) Cyclic Voltammtry(CV)

reversible 1 mM O2

irreversible

0.06 mM 2-nitropropane

“Reversible” means the reaction is fast enough to maintain equilibrium concentrations of the reactant and product at the electrode surface UltraMicroelectrode(~m diameter)

small size (implantable) small iR drop (useful in resistive and nonaqueous media small charging current (high fast scan possible, high sensitivity)

Figure 18.22 (a) scanning electron micrograph of a carbon fiber that has been etched to ~1m diameter by drawing through a flame. (b) Carbon fiber with thin, insulating coating formed by electrolytic copolymerization of phenol and 2-alkylphenol. Once coated, the tip of the fiber is the only electrochemically active surface. Cyclic Voltammetry with UME

Fast scan: linear diffusion dominates Slow scan: radial diffusion dominates

Ferrocene Oxidation

Scan direction 1 Iss = 4ronFCoDo Dt

3+ Ru(NH3)6 reduction

Iss = 4ronFCoDo

Scan direction Amperometry (i vs time) at fixed potential

Stirring Current, nA nA Current,

Time, s Figure 6.18 Amperometric response of Steady-state current a thick-film RuO electrode to step-wise because of convection 2 Current, nA nA Current, changes of the H2O2 concentration in a stirred PBS solution. Time, s Figure 3.13 Current measured during an amperometric experiment. A constant potential(-700 mV) as a function of time is applied to a plamar Au electrode in PBS saturated with air. Current, nA nA Current,

H2O2 concentration, mM Calibration curve obtained from the measurement in figure 6.18. Oxygen Electrode

Clark Electrode : Pt cathode held at -0.6 V vs Ag/AgCl + - cathode : O2 + 4H + 4e ↔ 2H2O anode : 2Ag + 2Cl- ↔ 2AgCl + 2e- Glucose Biosensor based on Clack Electrode

Biosensor (Amperometric) Limiting factor

Glucose + O2 → H2O2 + gluconic acid - - H2O2 + OH ↔ O2 + H2O + 2e

Pt electrode

Cellulose acetate membrane

: permeable to small molecules H2O2 GOx Polycarbonate film : permeable to glucose : impermeable to proteins and other constituents in blood st 1 generation : O2 electrode (stoichiometric limitation) nd 2 generation : H2O2 detection(interference) 0.5 vs Ag/AgCl CH OH Oxidant Mediators 2 CH2OH O O H O2 Glucose Biosensor O OH OH OH + H2O2 OH GOx OH OH OH Glucose Gluconolactone

Mred FAD Glucose

Glucose Oxidase 2e -

Mox FADH2 Gluconolactone

A.E.G. Cass, et. al., Anal. Chem., 1984, 56, 667-671 Electron Transport Mediator + -e-

Graphite electrode 1,1’-Dimethylferrocene 1,1’-Dimethylferricinium cation Enzyme-Linked Immunosorbent Assay (ELISA)

S SS P S P SSS P P P P P Color

Antigen

S Substrate (colorless)

P Product (colored) Antibody Antigen Secondary antibody

Yonsei University, Department of Chemistry Electroanalytical Chemistry Lab. Amperometric Immunoassay

Alkaline phosphatase 2- 2- R-OPO3 + H2O → ROH + HPO4 Phosphomonoester electroactive