1. Chapter 7 Electrochemistry §7.7 Thermodynamics of reversible cell

2. 7.7.1. Measurement of Electromotive forces (emf's) Can voltameter be used to measure electromotive force? V ERiRi RoRo U Discussion What is electromotive forces? High-impedance input

3. 1) Poggendorff’s compensation method i = 0, thermodynamic reversibility. Principle of potentiometer E W : working cell E x : test cell E s : standard cell A EsEs ExEx G EwEw ExEx EsEs K A B C1C1 C2C2

4. 2) Weston standard cell Drawing from Edward Weston's US Patent 494827 depicting the standard cell The Weston cell, is a wet-chemical cell that produces a highly stable voltage suitable as a laboratory standard for calibration of voltmeters. Invented by Edward Weston in 1893, it was adopted as the International Standard for EMF between 1911 and 1990.

5. Hg Hg 2 SO 4 Saturated CdSO 4 solution Cd(Hg) x + -- Cork sealed with paraffin or wax Commercial Weston Standard cell 2) Weston standard cell

6. E(T) /V = 1.01845 – 4.05  10 -5 (T/K –293.15) – 9.5  10 -7 (T/K –293.15) 2 + 1  10 -8 (T/K –293.15) 3 Weston standard cell Temperature-dependence of emf The original design was a saturated cadmium cell producing a convenient 1.018638 Volt reference and had the advantage of having a lower temperature coefficient than the previously used Clark cell

7. 2. Nernst equation and standard EMF of cell 1889, Nernst empirical equation cC + dD = gG + hH Walther H. Nernst 1920 Noble Prize Germany 1864/06/25~1941/11/18 Studies on thermodynamics physical meaning of E 

8. For a general electrochemical reaction: cC + dD = gG + hH Van’t Horff equation Theoretical deduction of Nernst Equation:

9. 7.7.3. Standard electromotive forces E Ө equals E when the activity of any chemical species is unit. For cell: Pb(s)-PbO(s)|OH – (c)|HgO(s)-Hg(l) Write out the cell reaction and Nernst equation.

10. For: Pt(s), H 2 (g, p  )|HCl(m) |AgCl(s)-Ag(s) Write out the cell reaction and Nernst equation.

11. EE E / V Experimental determination of standard electromotive force Cf. Levine, p. 430

12. 7.7.4. Temperature-dependence of emf's Temperature coefficient: For Weston Standard Cell: E/V = 1.018646 - 4.05  10 -5 (T/ ℃ -20) - 9.5  10 -7 (T/ ℃ -20) 2 + 1  10 -8 (T/ ℃ -20) 3 By differentiating the equation -  r G m = nFE with respect to temperature, we obtain

13. By measuring E and (  E/  T) p, thermodynamic quantities of the cell reaction can be determined. Because E and (  E/  T) p can be easily measured with high accuracy, historically, the thermodynamic data usually measured using electrochemical method other than thermal method.

14. 7.7.5. Thermodynamic quantities of ions The customary convention is to take the standard free energy of formation of H + (aq) at any temperatures to be zero. How to solve this deadlock?

15. Ion / kJ·mol -1 Ion / kJ·mol -1 H+H+ 0.000OH  -157.3 Li + -298.3Cl  -276.5 Na + -261.87Br  -131.2 K+K+ -282.3SO 4 2  -742.0 Ag + 77.1CO 3 2  -528.1 Standard free energies of formation of aqueous ions at 298.3 K H+H+ Cl  Br  II K+K+ Na + Mg 2+ Ca 2+ By definition

16. Exercise-1 At 298 K, for cell Ag(s)-AgCl(s)|KCl(m)|Hg 2 Cl 2 (s)-Hg(l), E = 0.0455V, (  E/  T) p = 3.38  10 -4 V·K -1. Write the cell reaction and calculate  r G m,  r S m,  r H m, and Q re. At 198 K, for cell Pt(s), H 2 (g, p  )|KOH(aq)|HgO(s)-Hg(l) E  = 0.926 V, product of water K w =10 -14. Given  f G m of HgO(s) is –58.5 kJ· mol -1, calculate  f G m of OH . Exercise-2

17. Self reading: Ira N. Levine, Physical Chemistry, 5 th Ed., McGraw-Hill, 2002. pp. 294-310 Section 10.10 standard-state thermodynamic properties of solution components pp. 426 Section 14.6 thermodynamics of galvanic cells Section 14.7 standard electrode potentials Section 14.8 concentration cells Section 14.9 liquid-junction potential