ELECTROCHEMISTRY CHEM 4700 CHAPTER 1

1. ELECTROCHEMISTRYCHEM 4700CHAPTER 1 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

2. CHAPTER 1 FUNDAMENTAL CONCEPTS

3. ELECTROCHEMISTRY - Electrochemistry is the study of the relations between chemical reactions and electricity - Electrochemical processes involve the transfer of electrons from one substance to another

4. ELECTROANALYTICAL TECHNIQUE - Deals with the relationship between electricity and chemistry - Deals with the measurement of electrical quantities (current, potential, charge) and their relationship to chemical parameters Applications Fabrication of flow detectors Quality control Environmental monitoring Electronics Electrochemical sensors

5. REDOX CHEMISTRY - Electron transfer occurs in redox reactions Oxidation - Loss of electrons Reduction - Gain of electrons Oxidizing agent (oxidant) is the species reduced Reducing agent (reductant) is the species oxidized

6. REDOX CHEMISTRY Oxidizing Agent - The species that gains electrons - The species that is reduced - Causes oxidation Aox + ne-↔ Ared Cu2+(aq) + 2e-↔ Cu(s)

7. REDOX CHEMISTRY Reducing Agent - The species that loses electrons - The species that is oxidized - Causes reduction Bred↔ Box + ne- Fe(s) ↔ Fe2+(aq)+ 2e-

8. REDOX CHEMISTRY Half Reactions - Just the oxidation or the reduction is given - The transferred electrons are shown Oxidation Half-Reaction - Electrons are on the product side of the equation Reduction Half-Reaction - Electrons are on the reactant side of the equation

9. REDOX CHEMISTRY Half Reactions - Oxidation half reaction Bred↔ Box + ne- Fe(s) ↔ Fe2+(aq)+ 2e- - Reduction half reaction Aox + ne- ↔ Ared Cu2+(aq) + 2e- ↔ Cu(s)

10. REDOX CHEMISTRY The Overall Reaction - Both an oxidation and a reduction must occur in a redox reaction - The oxidizing agent accepts electrons from the reducing agent Aox + Bred↔ Ared + Box Cu2+(aq) + Fe(s) ↔Cu(s) + Fe2+(aq) - Reducing agent - Oxidized species - Electron loss - Oxidizing agent - Reduced species - Electron gain

11. CHARGE Charge (q) of an electron = - 1.602 x 10-19 C Charge (q) of a proton = + 1.602 x 10-19 C C = coulombs Charge of one mole of electrons = (1.602 x 10-19 C)(6.022 x 1023/mol) = 9.649 x 104 C/mol = Faraday constant (F) C = n x F

12. CURRENT - The quantity of charge flowing past a point in an electric circuit per second Units Ampere (A) = coulomb per second (C/s) Charge (C) = current (A) x time (s)

13. VOLTAGE Potential Difference (E) - Work done by or on electrons when they move from one point to another Units: volts (V or J/C) Work (J) = E x C

14. ELECTRODE - Conducts electrons into or out of a redox reaction system - The electrode surface serves as a junction between an ionic conductor and an electronic conductor Examples platinum wire carbon (glassy or graphite) indium tin oxide (ITO) Gold Silver

15. ELECTROANALYTICAL MEASUREMENT • Electroactive Species • Donate or accept electrons at an electrode • - Can be made to oxidize or reduce • Chemical Measurements • - Involve homogeneous bulk solutions • Electrochemical Measurements • - Occur at the electrode solution interface

16. ELECTROANALYTICAL MEASUREMENT Two main types Potentiometric and Potentiostatic - The type of technique reflects the type of electrical signal used for quantitation - Techniques require at least two electrodes and an electrolyte (containing solution)

17. ELECTROANALYTICAL MEASUREMENT Potentiometric Technique - Based on a static (zero-current) situations - Based on measurement of the potential established across a membrane - Used for direct monitoring of ionic species (Ca2+, Cl-, K+, H+)

18. ELECTROANALYTICAL MEASUREMENT Potentiostatic Technique - Controlled-potential technique - Based on dynamic (non-zero-current) situations - Deals with the study of charge transfer processes at the electrode-solution interface - Resultant current is measured based on an electron transfer reaction derived from an electrode surface

19. ELECTROANALYTICAL MEASUREMENT Potentiostatic Technique - Chemical species are forced to gain or lose electrons (forced to oxidize or reduce) - Referred to as ‘electron pressure’ Advantages - High sensitivity and selectivity - Very low detection limits - Wide range of electrode types - Wide range of linearity - Portable and low cost instrumentation

20. ELECTROANALYTICAL MEASUREMENT Electrochemical Cell - Made up of the electrodes and the contacting sample solution Three types of electrodes in a cell - Working (Indicator) Electrode: responds to target analyte - Couter (Auxiliary) Electrode: current carrying electrode - Reference Electrode: provides known and constant potential and does not depend on analyte properties

21. CHEMICAL CHANGE Spontaneous Process - Takes place with no apparent cause Nonspontaneous Process - Requires something to be applied in order for it to occur (usually in the form of energy)

22. ELECTROLYSIS - Voltage is applied to drive a redox reaction that would not otherwise occur Examples - Production of aluminum metal from Al3+ - Production of Cl2 from Cl-

23. ELECTROLYTIC CELL - Nonspontaneous reaction - Requires electrical energy to occur - Consumes electricity from an external source

24. GALVANIC CELL - Spontaneous reaction - Produces electrical energy - Can be reversed electrolytically for reversible cells Example Rechargeable batteries Conditions for Non-reversibility - If one or more of the species decomposes - If a gas is produced and escapes

25. GALVANIC CELL - Also known as voltaic cell - A spontaneous redox reaction generates electricity - One reagent is oxidized and the other is reduced - The two reagents must be separated (cannot be in contact) - Electrons flow through a wire (external circuit)

26. GALVANIC CELL Oxidation Half reaction - Loss of electrons - Occurs at anode (negative electrode) - The left half-cell by convention Reduction Half Reaction - Gain of electrons - Occurs at cathode (positive electrode) - The right half-cell by convention

27. GALVANIC CELL Salt Bridge - Connects the two half-cells (anode and cathode) - Filled with gel containing saturated aqueous salt solution (KCl) - Ions migrate through to maintain electroneutrality - Prevents charge buildup that may cease the reaction process Preparation - Heat 3 g of agar and 30 g of KCl in 100 mL H2O - Heat until a clear solution is obtained - Pour into a U-tube and allow to gel - Store in a saturated aqueous KCl

28. GALVANIC CELL For the overall reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) e- Voltmeter e- - + Cu electrode Zn electrode Cl- K+ Zn2+ Salt bridge (KCl) Cu2+ Anode Oxidation Zn(s) → Zn2+(aq) + 2e- Cathode Reduction Cu2+(aq) + 2e- → Cu(s)

29. GALVANIC CELL Voltage or Potential Difference (E) - Is the voltage measured - Measured by a voltmeter (potentiometer) connected to electrodes Greater Voltage - More favorable net cell reaction - More work done by flowing electrons

30. GALVANIC CELL Line Notation Phase boundary: represented by one vertical line Salt bridge: represented by two vertical lines Fe(s) FeCl2(aq) CuSO4(aq) Cu(s)

31. STANDARD POTENTIALS Electrode Potentials - A measure of how willing a species is to gain or lose electrons Positive Voltage (spontaneous process) - Electrons flow into the negative terminal of voltmeter (flow from negative electrode to positive electrode) Negative Voltage (nonspontaneous process) - Electrons flow into the positive terminal of voltmeter (flow from positive electrode to negative electrode) Conventionally - Negative terminal is on the left of galvanic cells

32. STANDARD POTENTIALS Standard Reduction Potential (Eo) - Used to predict the voltage when different cells are connected - Potential of a cell as cathode compared to standard hydrogen electrode - Species are solids or liquids - Activities = 1 - We will use concentrations for simplicity Concentrations = 1 M Pressures = 1 bar

33. STANDARD POTENTIALS Standard Hydrogen Electrode (SHE) - Used to measure Eo for half-reactions - Connected to negative terminal - Pt electrode - Acidic solution in which [H+] = 1 M - H2 gas (1 bar) is bubbled past the electrode H+(aq, 1 M) + e- ↔ 1/2H2 (g, 1 bar) Conventionally, Eo = 0 for SHE

34. STANDARD POTENTIALS The Eo for Ag+ + e- ↔ Ag(s) is 0.799 V Implies that if a sample of silver metal is placed in a 1 M Ag+ solution, a value of 0.799 V will be measured with S. H. E. as reference Pt(s) H2(g, 1 bar) H+(aq, 1 M ) Ag+ (aq, 1 M) Ag(s) SHE Ag+ (aq, 1 M) Ag(s)

35. STANDARD POTENTIALS Silver does not react spontaneously with hydrogen 2H+(aq) + 2e- → H2(g) Eo = 0.000 V Ag+(aq) + e- → Ag(s) Eo = +0.799 V Reverse the second equation (sign changes) Ag(s) → Ag+(aq) + e- Eo = -0.799 V Multiply the second equation by 2 (Eo is intensive so remains) 2Ag(s) → 2Ag+(aq) + 2e- Eo = -0.799 V Combine (electrons cancel) 2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = -0.799 V

36. STANDARD POTENTIALS Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Cu2+(aq) + 2e- → Cu(s) Eo = +0.339 V Zn2+(aq) + 2e- → Zn(s) Eo = -0.762 V Reverse the second equation (sign changes) Zn(s) → Zn2+(aq) + 2e- Eo = +0.762 V Combine (electrons cancel) Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V Eo is positive so reaction is spontaneous Reverse reaction is nonspontaneous

37. STANDARD POTENTIALS - Half-reaction is more favorable for more positive Eo Formal Potential - The potential for a cell containing a specified concentration of reagent other than 1 M

38. STANDARD POTENTIALS Eo (V) 2.890 1.507 1.280 1.229 0.799 0.339 0.000 -0.402 -0.440 -0.763 -1.659 -2.936 -3.040 Half Reaction F2 + 2e-↔ 2F- MnO4- + 5e-↔ Mn2+ Ce4+ + e- ↔ Ce3+ (in HCl) O2 + 4H+ + 4e- ↔ 2H2O Ag+ + e- ↔ Ag(s) Cu2+ + 2e- ↔ Cu(s) 2H+ + 2e- ↔ H2(g) Cd2+ + 2e- ↔ Cd(s) Fe2+ + 2e- ↔ Fe(s) Zn2+ + 2e- ↔ Zn(s) Al3+ + 3e- ↔ Al(s) K+ + e- ↔ K(s) Li+ + e- ↔ Li(s) Oxidizing agents Reducing agents Increasing reducing power Increasing oxidizing power

39. NERNST EQUATION - For systems controlled by the laws of thermodynamics O + ne- ↔ R E = electrode potential Eo = standard potential for the redox reaction R = gas constant = 8.314 J/K-mol T = absolute temperature in Kelvin F = Faraday’s constant = 9.649 x 104 C/mol n = number of electrons transferred

40. NERNST EQUATION For the half reaction aA + ne-↔ bB The half-cell potential (at 25 oC), E, is given by

41. NERNST EQUATION For the overall reaction aA + bB ↔ cC + dD The potential at 25 oC is given by

42. NERNST EQUATION - E = Eo when [O] = [R] = 1M - Concentration for gases are expressed as pressures in bars or atm - Concentrations for pure solids, liquids, and solvents are omitted - Reduction is more favorable on the negative side of Eo

43. NERNST EQUATION - When a half reaction is multiplied by a factor Eo remains the same - For a complete reaction Ecell = E+ - E- and Eo = E+o - E-o E+ = potential at positive terminal E- = potential at negative terminal

44. NERNST EQUATION For the Cu – Fe cell at standard conditions Cu2+ + 2e- ↔ Cu(s) 0.339 V Fe2+ + 2e- ↔ Fe(s) -0.440 V Ecell = 0.779 V Galvanic Reaction Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq) Fe Fe2+ (1M) Cu2+ (1 M) Cu

45. NERNST EQUATION - Positive E implies spontaneous forward cell reaction - Negative E implies spontaneous reverse cell reaction If cell runs for long - Reactants are consumed - Products are formed - Equilibrium is reached - E becomes 0 - Reason why batteries run down

46. FARADAIC PROCESSES - Obtaining a current response that is related to the concentration of the target analyte - Accomplished by monitoring the transfer of electrons during a redox process of the analyte O + ne- ↔ R O = oxidized form of the analyte R = reduced form of the analyte O and R are a redox couple - Occurs in a potential region within which the electron transfer is thermodynamically or kinetically favorable

47. FARADAIC CURRENT - The current resulting from a change in oxidation state of an electroactive species - A measure of the rate of the redox reaction - Obeys Faraday’s law - The reaction of 1 mol of a substance involves a charge of nF coulombs

48. VOLTAMMOGRAM - Current versus potential plot - Current on vertical axis and excitation potential on horizontal axis - Electrode reactions involve several steps and can be complicated - The rate is determined by the slowest step and depends on the potential range

49. STEPS IN SIMPLE REACTIONS - Mass transport of the electroactive species to the electrode surface - Electron transfer across the solution-electrode interface - Transport of the product back to the bulk solution

50. STEPS IN COMPLEX REACTIONS - In addition to the three steps on previous slide - Chemical processes occur - Surface processes occur - These additional processes may precede or follow the electron transfer