Chapter 7 Electrochemistry

1. Chapter 7 Electrochemistry §7.7 Thermodynamics of reversible cell

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

3. Es A Ew Ex A B C1 C2 Ex G K Es 1) Poggendorff’s compensation method i = 0, thermodynamic reversibility. EW: working cell Ex: test cell Es: standard cell Principle of potentiometer

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. Cork sealed with paraffin or wax Saturated CdSO4 solution Hg2SO4 Cd(Hg)x Hg -- + 2) Weston standard cell Commercial Weston Standard cell

6. Weston standard cell 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 Temperature-dependence of emf 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

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

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

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

10. For: Pt(s), H2 (g, p)|HCl(m) |AgCl(s)-Ag(s) Write out the cell reaction and Nernst equation.

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

12. 7.7.4. Temperature-dependence of emf's 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 Temperature coefficient: By differentiating the equation - rGm = 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 How to solve this deadlock? The customary convention is to take the standard free energy of formation of H+(aq) at any temperatures to be zero.

15. K+ Na+ Cl Mg2+ Ca2+ H+ Br I By definition Standard free energies of formation of aqueous ions at 298.3 K

16. Exercise-1 At 298 K, for cell Ag(s)-AgCl(s)|KCl(m)|Hg2Cl2(s)-Hg(l), E = 0.0455V, (E/T)p = 3.38  10-4 V·K-1. Write the cell reaction and calculate rGm, rSm, rHm, and Qre. Exercise-2 At 198 K, for cell Pt(s), H2(g, p)|KOH(aq)|HgO(s)-Hg(l) E = 0.926 V, product of water Kw=10-14. Given fGm of HgO(s) is –58.5 kJ· mol-1, calculate fGm of OH.

17. Self reading: Ira N. Levine, Physical Chemistry, 5th 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