1. Nov 16, 2004 Introduction to Electroanalytical Chemistry Lecture Date: April 8 h, 2013

2. Electroanalytical Methods  Matched against a wide range of spectroscopic and chromatographic techniques, the techniques of electroanalytical chemistry find an important role for several reasons: –Electroanalytical methods are often specific for a particular oxidation state of an element –Electrochemical instrumentation is relatively inexpensive and can be miniaturized –Electroanalytical methods provide information about activities (rather than concentration)

3. History of Electroanalytical Methods  Michael Faraday: the law of electrolysis –“…the amount of a substance deposited from an electrolyte by the action of a current is proportional to the chemical equivalent weight of the substance.”  Walter Nernst: the Nernst equation (Nobel Prize 1920)  Jaroslav Heyrovsky: polarography (Nobel Prize 1959) Walter Nernst (1864-1941) Michael Faraday (1791-1867) Jaroslav Heyrovsky (1890-1967)

4. Main Branches of Electroanalytical Chemistry  Key to measured quantity: I = current, E = potential (in volts, V), R = resistance, G = conductance, q or Q = quantity of charge (in coulombs, C), t = time, vol = volume of a standard solution, m = mass of an electrodispensed species Interfacial methods Bulk methods Static methods (I = 0) Dynamic methods (I > 0) Potentiometry (E) Conductometry (G = 1/R) Controlled potential Constant current Voltammetry (I = f(E)) Amperometric titrations (I = f(E)) Based on Figure 22-9 in Skoog, Holler and Crouch, 6th ed. Electro- gravimetry (m) Coulometric titrations (q = It)

5. Main Branches of Electroanalytical Chemistry  Potentiometry: measure the potential of electrochemical cells without drawing substantial current –Examples: pH measurements, ion-selective electrodes, titrations (e.g. KF endpoint determination)  Coulometry: measures the quantity of charge required to drive an electrolytic oxidation/reduction to completion –Examples: titrations (KF titrant generation), “chloridometers” (AgCl)  Voltammetry: measures current as a function of applied potential under conditions that keep a working electrode polarized –Examples: cyclic voltammetry, many biosensors

6. Electrochemical Cells Measurement Salt bridge (aqueous KCl) Zn electrode Cu electrode e- Zn  Zn 2+ (aq) + 2e - a Zn 2+ = 0.010 Anode Cu 2+ (aq) + 2e-  Cu(s) a Cu 2+ = 0.010 Cathode 0.010M ZnSO 4 solution 0.010M CuSO 4 solution  Zinc (Zn) wants to ionize more than copper (Cu).  We can use this behavior to construct a cell:

7. e- reference electrode counter electrode working electrode indicator electrode detector electrode Potentiometry: Measures equilibrium E (the potential, or voltage) Amperometry: Control E, measures I as function of time Coulometry: Control E, measure total Q over a period of time control measurement Electrochemical Cells and Analytical Methods

8. Electrochemical Cells  Galvanic cell: a cell that produces electrical energy  Electrolytic cell: a cell that consumes electrical energy  Chemically-reversible cell: a cell in which reversing the direction of the current reverses the reactions at the two electrodes

9. Conduction in an Electrochemical Cell  Electrons serve as carriers (e.g. moving from Zn through the conductor to the Cu)  In the solution, electricity involves the movement of cations and anions –In the salt bridge both chloride and potassium ions move  At the electrode surface: an oxidation or a reduction occurs –Cathode: the electrode at which reduction occurs –Anode: the electrode at which oxidation occurs

10. Oxidation occurs when a chemical species loses an electron. LEO = lose electron is oxidation Reduction is when a species gains an electron. GER = gain an electron is reduction For example, the chemical reaction can be decomposed into two half reactions: “Leo the Lion Says Ger”

11. Faradaic and Non-Faradaic Currents The electrical double layer Mass Transfer occurs by: Convection Migration Diffusion  Faradaic (governed by Faraday’s law): direct transfer of electrons, i.e. oxidation at one and reduction at the other electrode  Non-Faradaic: increasing charge of the double layer

12. Fundamentals Electrical charge (q) is measured in coulombs (C). The charge associated with chemical species is related to the number of moles through the Faraday constant: F = 96,485.3 (~96,500) C/mole. Electrical current, I, is measured in Amperes (A). Current is the amount of charge that passes in a unit time interval (seconds). Ohm's law relates current to potential (E) through the resistance (R) of a circuit by E=IR. The potential is measured in Volts (V) and the resistance in Ohms (  ).

13. Power (P) is measured in Watts (W = J/s) and is related to the current and potential by P= IE. The work is measured in Joules (J) and is related to the potential and the amount of charge by work=q E. The relationship between the standard Gibb's free energy change,  G° (J/mole), and the standard electromotive force (EMF), E° (V), is given by  G°=-n F E° where n is the number of electrons transferred and superscript on E 0 refers to ‘standard state.’ Fundamentals

14. Fundamentals: The Nernst Equation ● The Nernst equation gives the cell potential E (in volts): ● Q (the activity quotient) is the ratio of products over reactants as in equilibrium calculations. For the generic reaction: ● Q is given by: ● Where the A’s are activities. For low-concentration solutions (low ionic strengths): F = faraday (constant) n = # moles electrons in process E 0 = standard potential for cell

15. Activity Coefficients a x =  x [C] IDEAL [ ] < 10 -3 NON-IDEAL [ ] > 10 -3 in dilute solutions,  = 1  < 1 ● When the concentrations of ions in a solution are greater than approximately 0.001 M, a shielding effect occurs around ions. The electrolyte behaves in a non-ideal manner. ● Cations tend to be surrounded by nearby anions and anions tend to be surrounded by nearby cations ● We use the activity (a) and the activity coefficent (  ) to explain this effect:

16. Electrode Potentials  The reactions in an electrochemical cell can be thought of as two half-cell reactions, each with its own characteristic electrode potential –These measure the driving force for the reaction –By convention, always written as reductions  Standard electrode potential (E 0 ): the measure of individual potential of an electrode at standard ambient conditions (298K, solutes at a concentration of 1 M, and gas pressure at 1 bar).

17. Some Standard Electrode Potentials ReactionE 0 at 298K (Volts) Cl 2 (g) + 2e -  2 Cl 2 +1.359 O 2 (g) + 4H + + 4e -  2 H 2 O +1.229 Ag + + e -  Ag(s) +0.799 Cu 2+ + 2e -  Cu(s) +0.337 Hg 2 Cl 2 + 2e -  2Hg(l) + 2 Cl 2 +0.268 2H + + 2e -  H 2 (g) 0.000 AgI(s) + e -  Ag(s) + I 2 -0.151 Cd 2+ + 2e -  Cd(s) -0.403 Zn 2+ + 2e -  Zn(s) -0.763 See appendix 3 in Skoog et al. for a more complete list

18. The Standard Hydrogen Electrode (SHE)  A universal reference, but is really a hypothetical electrode (not used in practice) –Uses a platinum electrode, which at its surface oxidizes 2H + to H 2 gas. –Very sensitive to temperature, pressure, and H + ion activity  Because the SHE is difficult to make, the saturated calomel electrode (SCE) is used instead. –Calomel = mercury (I) chloride

19. Q: What is the electrode potential for the electrode Ag/AgI(s)/I - (0.01 M) ? The overall reaction for this electrode is This reaction cannot be found in tables of reduction potentials. But the reaction is comprised of two components Electrode Potentials

20. We can initially ignore the fact that the electrode contains AgI and find E for the silver ion reduction. Electrode Potentials

21. The Glass pH Electrode ● One of the most common potentiometric measurements is pH (a so-called “p-Ion” measurement). ● The common glass pH electrode makes use of junction potentials to determine the hydronium ion concentration in a sample solution. ● A typical glass pH electrode is configured as shown here: Image from http://www.isa.org/Images/InTech/Oct2002/20021048.gif Glass pH electrode Basic layout

22. The glass pH electrode is used with a Ag/AgCl reference electrode. For most modern pH electrodes the reference electrode is incorporated with the pH indicator electrode. A small frit or hole connects the reference electrode and the sample solutions The Glass pH Electrode

23. pH Measurements ● A combination pH electrode combines the indicator and reference into a single unit. ● The potential of this cell is: ● where E ij and E oj are the junction potentials at the inner and outer layers of the glass membrane. ● Junction potential: occurs at the interface of two electrolytes, caused by unequal diffusion rates of cation and anions across the boundary (e.g the frit in a salt bridge)

24. More About pH Measurements ● The surface of the glass is hydrated, which allows exchange of hydronium ions for the cation in the glass (sodium or lithium). ● There are four interface regions, the external solution and hydrated glass, hydrated glass and dry glass on the outside, dry glass and hydrated glass on the inside, and hydrated glass and the internal solution. ● If the glass is uniform, the two hydrated glass/dry glass interfaces should be identical and should have the same junction potential. ● Since the glass interface junction potentials then cancel each other, the junction potential is then the difference between the internal and external solutions.

25. pH Measurements  For a real electrode, the two surfaces will not be identical and the constant E ref needs to be determined experimentally (termed the asymmetry potential). The constant  is termed the electromotive efficiency.

26. pH Measurements

27.  Errors in pH measurements with glass electrodes arise from the following effects: –Calibration problems (e.g. drift, or error in the calibration) –Junction potential –High [Na+] interacting with electrode –High acid concentration –Equilibration time –Temperature control  Typical electrodes have the following performance: –Accuracy = +/- 0.02 pH units –Precision = +/- 0.002 pH units pH Electrodes: Errors, Accuracy and Precision

28. Modern pH electrodes are usually of the "combination" type, meaning that a single cylinder contains both the reference electrode, and a glass membrane electrode. Schematically, the total cell may be expressed as SCE//test solution ([H 3 O + ]=a1)/glass membrane/[H 3 O + ]=a2, Cl-/AgCl(s)/Ag The Combination pH Electrode

29. A Modern Combination pH Electrode

30. Automatic pKa and log P Determination pKa (ionization constant) and log P (octanol/water partition) are important physical parameters that play critical roles in determining how compounds behave in physiological environments and how they interact with enzymes, receptors and cell membranes The Sirius GLpKa system: combination pH electrode sample tray reagents liquid dispensors

31. The Ion Selective Electrode (ISE) ● An ISE generally consists of the ion-selective membrane, an internal reference electrode, an external reference electrode, and a voltmeter. ● Example: an ISE for fluoride (F - )

32. Conductometry  Conductometry: Detection of electrical conductivity –Key analytical applications: conductometric detection in ion-exchange chromatography (IEC or IC) and capillary electrophoresis (CE)  Also used to detect titration endpoints

33. Reading Material ● Optional: ● A. J. Bard and L. R. Faulkner, “Electrochemical Methods”, 2 nd Ed., Wiley, 2001. ● F. Scholz, Ed. “Electroanalytical methods”, 2 nd Ed., Springer, 2010.

34. 0.05916pH log E then composition, fixed a hassolution internal theIf log a If the two solutions are identical log,',,', mem,',,',,,,,H,,,',,',,,,,,',,',,, + +=+-= -= = -=--= ++ + + + ++ ++ + + + + ka F RT a F a a F E a aa aa F a a F a a F E innsolHoutnsolH innsolH outnsolH mem outglassHinglass outglassHinnsolH outnsolHinglassH outglassH outnsolH innsolH inglassH mem  pH Measurements

35. Electrochemical pH Measurements Concluded Consider a typical problem related to the use of the combination pH electrode. Recall that E cell = L - 0.0592 V pH QUESTION: If E cell = -0.115 V at a pH of 4.00, what is the pH of a solution for which E cell is -0.352 V?

36. ANSWER: First, find L from the measurement of the standard: -0.115 V = L -0.0592 x pH -0.115 V = L -0.0592 x 4.00 Therefore, L = 0.122 V Second, use this value of L to find pH: -0.352 V = 0.122 V - 0.0592 V x pH pH = (0.122 V -(-0.352 V))/0.0592 pH = 7.84

37. QUESTION: What does the pH meter read if the pH is 7.00 in a 1 M salt solution having 1 M Na + ions present? ANSWER: [H + ] obs = 1 x 10 -7 + 1 x 10 -12 Conclusion -- the pH meter reads the true pH under these conditions.

38. Q: Why does the pH change the interfacial potential of the glass/aqueous interface? A: The motion of the sodium ions leave behind a negatively charged glass layer that is neutralized to a lesser or greater extent according to the pH. More explanation about how a pH meter really works: The sodium ions must move through the dry part of the membrane and this process is slow. For this reason, the membrane is made very thin. Also, a nonperturbing (low-current) voltmeter is used to read the cell voltage so that only a few sodium ions must move through the dry glass in a given time period. pH Measurements