1. PN. ROZAINI ABDULLAH PROGRAMME INDUSTRIAL CHEMICAL PROCESS PRT 140 Physical Chemistry Equilibrium Electrochemistry

2. Subtopics Half-Reactions and Electrodes Varieties of Cells The Electromotive Force Standard Potentials Applications of Standard Potentials Impact on Biochemistry/Technology: Energy Conversion in Biological Cells

3. Equilibrium Electrochemistry Electrochemical system : heterogeneous system in which there is a difference of electrical potential between 2 or more phases.

4. Equilibrium Electrochemistry An electrochemical cell consists:  two electrodes (or metallic conductors)  an electrolyte (an ionic conductor – may be a solution, a liquid or a solid). An electrode & its electrolyte comprise an electrode compartment.  two electrodes may share the same compartment  if the electrolytes are different, the two compartments may be joined by a salt bridge [a tube containing a concentrated electrolyte solution (potassium chloride in agar jelly)] that completes the electrical circuit & enables the cell to function.

5. A galvanic cell is an electrochemical cell that produces electricity as a result of the spontaneous reaction occurring inside it. [ Eg: Daniell Cell]

6. A electrolytic cell is an electrochemical cells in which a non-spontaneous reaction is driven by an external source of current. Ionic conductor (at least 1 phase) to allow electronic charge to be transferred Terminals (electrods) made from same metal External Source Electrical current

7. HALF-REACTIONS AND ELECTRODES TermDefinition/Explanation Oxidation (Anode, - ve) The removal of electrons from a species Reduction (cathode, +ve) The addition of electrons to a species Redox reactionReaction in which there is a transfer of electrons from 1 species to another and hence a change in oxidation no of an element. Reductant (reducing agent) The electron donor Oxidant (oxidizing agent) The electron acceptor Half-reactionsConceptual reactions showing the gain of electrons Redox coupleThe reduce and oxidized species in a half reaction

8. Generally, we write it as Ox/Red and the corresponding reduction half reaction is Ox + v e -  Red Anode Cathode

9. Varieties Of Cells The simplest type of cell: single electrolyte common to both electrodes. When spontaneous reaction takes place in a galvanic cell, electrons are deposited in one electrode (the side of oxidation, the anode) Collected from another (the site of reduction, the cathode) So, there is a net flow of current, which can be used to do work. Note that the + sign of the cathode can be interpreted as indicating the electrode at which electrons enter the cell The sign – of the anode as where the electrons leave the cell.

10. Varieties Of Cells The Daniell cell - The electrodes are immersed in different electrolytes. - Normally porous ceramics barrier separates a compartment containing electrolytes (prevent extensive mixing of the solutions by convection currents but allow ions to pass from solution to the other) - The redox couple at one electrode is Cu 2+ /Cu and at the other is Zn 2+ /Zn. Anode (Negative) Cathode (positive) Cu 2+ (aq) + Zn  Cu + Zn 2+ (aq)

11. Electrolyte concentration cell: the electrode compartments are identical except for the concentration of the electrolytes. Electrode concentration cell: the electrodes themselves have different concentration, either because they are gas electrodes operating at different pressures or because they are amalgams (sol in mercury) with different concentration.

12. Example 1 Express the following reactions in terms of reduction half- reactions. a) The dissolution of silver chloride in water: (Note: it is not a redox reaction.) b) The formation of H 2 O from H 2 and O 2 in acidic solution.

13. Solution 1 Oxidant (reduction) = electron acceptor Reductant (oxidation) = electron donor Reduction half reaction: Redox couple: AgCl/Ag, Cl - and Ag + /Ag respectively. Ag + (aq) + e -  Ag (s)

14. Solution 2 The formation of H 2 O from H 2 and O 2 in acidic solution. Oxidant (reduction) = electron acceptor Reductant (oxidation) = electron donor 4H + (aq) + 4e -  2H 2 (g) Reduction half reaction: O 2 (g) + 4H + (aq) +4e -  H 2 O(l) Redox reaction: H + /H 2 and O 2, H+/H 2 O H 2 (g) + O 2 (g)  2H 2 O (l)

15. Liquid Junction Potentials Liquid junction potentials (E lj ):  When 2 different electrolyte solutions in contact, there is an additional source of potential difference across the interface of the two electrolytes.  E.g. In the Daniel cell (i) two different electrolyte solutions are in contact, (ii) different concentration of hydrochloric acid. In different conc of HCl, at the junction, the mobile H + ions diffuse into more dilute solution. The bulkier Cl - ions follow but initially do so more slowly, which result in a potential difference at the junction. The potential then settle downs to a value such that, after a that brief period, the ions diffuse at the same rate.  The contribution of the liquid junction to the potential can be reduced (to about 1 to 2 mV) by joining the electrolyte compartments through a salt bridge.

16. Galvanic cell without liquid junction. Galvanic cell with liquid junction (salt bridge).

17. Notation 1) Phase boundaries are denoted by a vertical bar. Eg: Pt(s) І H 2 (g) І HCl(aq) І AqCl(s) І Ag(s) 2) A liquid junction is denoted by 3) Interface is denoted by a double vertical line ||. Fig 1:  Zn (s)|ZnSO4 (aq) CuSO4 (aq) |Cu (s) Fig 2:  Zn (s)|ZnSO4 (aq)||CuSO4 (aq) |Cu (s) Fig 2 Fig 1

18. The Cell Potential (Electromotive Force) The cell reaction IS THE REACTION IN THE CELL WRITEN ON THE ASSUMPTION THAT: 1 st : write the right hand half-reaction as a reduction (Assumption: spontaneous reaction is 1 in which reduction is taking place in the right hand compartment). 2 nd : subtract from it the left-hand reduction half-reaction. (By implication, the electrode is the site of oxidation) In the cell: Zn(s)|ZnSO4(aq)||CuSO4(aq)|Cu(s) Right-hand electrode: Cu 2+ (aq)+2e -  Cu(s) Left-hand electrode: Zn 2+ (aq)+2e -  Zn(s) Overall cell reaction: Cu 2+ (aq)+ Zn(s)  Cu(s) +Zn 2+ (aq)

19. The Nernst Equation A cell in which the overall cell reaction has not reached chemical equilibrium can do electrical work as the reaction drives electrons through an external circuit.  The work that a given transfer of electrons can accomplish depends on the potential difference between the two electrodes.  This potential differences is called the cell potential and is measured in volts, V (1 V = 1 JC -1 s).  A cell in which the overall reaction is at equilibrium can do no work, & then the cell potential is zero.

20. When expressed in terms of a cell potential, the spontaneous direction of change can be expressed in terms of the cell emf.  The reaction is spontaneous when E>0.  The reverse reaction is spontaneous when E<0.  When the cell reaction is at equilibrium, the cell potential is zero. Note: The potential difference is called the electromotive force (emf), E

21. The Nernst equation relates the cell’s emf (E) to the activities a i of the substances in the cell’s chemical reaction & to the standard emf of the cell (E Ѳ )(the cell’s chemical reaction).  where F = Faraday constant, F=eN A v = the stoichiometric coefficient of the electron in the half-reactions. Q = the reaction (or activity) quotient,

22. A practical form of the Nernst equation is  because at 25 0 c,

23. Cells At Equilibrium Suppose the reaction has reached equilibrium; then Q = K (K= the equilibrium constant of the cell reaction). A chemical reaction at equilibrium cannot do work, & hence it generates zero potential difference between the electrodes of a galvanic cell. Setting E=0, Q=K:  the Nernst equation: Equilibrium constant and standard cell potential Equilibrium constant and standard cell potential

24. Example 2 Three different galvanic cells have standard electromotive force (E Ѳ ) of 0.01, 0.1 and 1.0V, respectively, at 25 0 C. Calculate the equilibrium constants (K) of the reactions that occur in these cells assuming the charge number (v) for each reaction is unity.

25. Solutions For E Ѳ = 0.01V, = 1.476 For E Ѳ = 0.1V, K = 49.0 For E Ѳ = 1.0V, K = 8.02 x 10 16

26. Standard Potential (Standard Electromotive Force) A galvanic cell: combination of 2 electrodes, each of which can be considered to make a characteristic contribution to the overall cell potential. Although it is not possible to measure the contribution of a single electrode, we can define the potential of one of the electrodes as zero an then assign values to others on that basis. The specially selected electrode is the standard hydrogen electrode (SHE): at all temps Pt (s) | H 2 (g) | H + (aq) E θ = 0

27. To achieve a standard conditions, the activity of the hydrogen ions must be 1 (that is, pH=0) and the pressure of H 2 gas = 1 bar. The standard potential, E θ, of another couple is then assigned by constructing a cell in which it is the right-hand electrode and the SHE is the left-hand electrode.

28. Determination Of Standard Cell Potential The principle use for standard potentials is to calculate the standard potential of a cell formed from any two electrodes. To do so, we subtract the standard potential of the left-hand electrode from the standard potential of the right-hand electrode: Note that all standard electrode potentials are reduction potentials. The reduction potentials is independent of the stoichiometric coefficients in the cell reaction. E 0 cell = E 0 R - E 0 L

29. Example 3 Consider the following galvanic a cell: What is: (a) the cell reaction? (b) the standard electromotive force? (c) The equilibrium constant? Given Table 9.1 standard potential for: Cu 2+ (aq)+2e - Cu(s) E 0 =0.34V Zn 2+ (aq)+2e - Zn(s) E 0 = - 0.76V

30. Solutions (a) & (b) Right-hand electrode: Cu 2+ (aq)+2e - Cu(s) E 0 =0.34V Left-hand electrode: Zn 2+ (aq)+2e - Zn(s) E 0 = - 0.76V Overall cell reaction:Cu 2+ (aq)+ Zn(s) Cu(s) +Zn 2+ (aq) E 0 = 0.34 – (-0.76) V E 0 = 1.10 V (c) The equilibrium constant: K = 1.80 x 10 37

31. Exercise 4 For the cell Cu | Cu + (aq) | Ag + (aq) | Ag | Cu’ Write the cell reaction Find E θ cell Calculate the equilibrium constant Given that Cu 2+ /Cu = 0.339 V Ag + /Ag = 0.7792 V

32. Application Of Standard Potentials The electrochemical series: the metallic elements (and hydrogen) arranged in the order of their reducing power as measured by their standard potentials in aqueous solution. A metal low in the series (with a lower standard potential) can reduce the ions of metals with higher standard potentials. For eg, to determine whether zinc can displace Mg from aqueous sol at 298K, we note that Zn lies above Mg in electrochemical series, so Zn cannot reduce Mg ions in sol. Zn can reduce H ions, H lies higher in the series.

33. Applications Of Standard Potentials The cell potential is used to measure the activity coefficient of electroactive ions The standard cell potential is used to infer the equilibrium constant of the cell reaction Species selective electrodes contribute the potential that is characteristics of certain ions in solution The temperature coefficient of the cell potential is used to determine the standard entropy and enthalpy of reaction.

34. Impact On Biochemistry/Technology The whole of life’s activities depends on the coulping of exergonic & endergonic reactions, for the oxidation of food drives other reactions forward. In biological cells, the energy released by the oxidation of foods is stored in adenosine triphosphate (ATP).  the essence of the action of ATP is its ability to lose its terminal phosphate group by hydrolysis & to form adenosine diphosphate (ADP). where Pi denotes an inorganic phosphate group e.g. H 2 SO 4.

35. Examples:  Glycolysis – the oxidation of glucose to CO 2 and H 2 O by O2 (the breakdown of foods is coupled to the formation of ATP in the cell).  Glycolysis is the main source of energy during anaerobic metabolism, a form of metabolism in which inhaled O 2 does not play a role.  The citric acid cycle & oxidative phosphorylation are the main mechanisms for the extraction of energy from carbohydrates during aerobic metabolism (in which inhaled O 2 does play a role).

36. Microbial Fuel Cells (MFC) Is a device that converts chemical energy to electrical energy by the catalytic reaction of microorganisms. A typical MFC consists of anode and cathode compartments separated by cation (+vely charged ion) specific membrane. In anode compartment, fuel (plant/wastewater)/substrates) is oxidized by microorganisms generating electrons and protons. Electrons are transferred to cathode compartment thru external circuit, while protons are transferred thru the membrane. Electron and protons are consumed in the cathode compartment, combining to form H 2.

37. Applications of MFC As power source Microbes would consume waste materials from the wastewater and produced supplementary power for a plant. MFC are very clean and efficient method of energy production. Microbes can consumed waste materials to generate energy.

38. Exercises Write the cell reaction and electrode half-reactions and calculate the standard potential of each of the following cells: a) Pt | Cl (g) | HCl (aq) | | K 2 Cr 2 O 4 (aq) | Ag 2 Cr 2 O 4 (s) |Ag b) Pt | Fe 3+ (aq), Fe 2+ (aq) | | Sn 4+ (aq), Sn 2+ (aq)| Pt c) Cu | Cu 2+ (aq) | | Mn 2+ (aq), H + (aq)| MnO 2 (s)|Pt

39. Exercises Calculate the equilibrium constants of the following reactions at 25oC from standard potential data: a) Sn (s) + CuSO 4 (aq)  Cu(s) + SnSO 4 (aq) b) Cu2+ (aq) + Cu (s)  2Cu + (aq)

40. Key Ideas

42. …OF THIS CHAPTER…” …OF THIS COURSE….” The End