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Chemical equilibrium: electrochemistry 자연과학대학 화학과 박영동 교수.

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Presentation on theme: "Chemical equilibrium: electrochemistry 자연과학대학 화학과 박영동 교수."— Presentation transcript:

1 Chemical equilibrium: electrochemistry 자연과학대학 화학과 박영동 교수

2 Examples In distilled water Hg 2 (IO 3 ) 2 ⇄ Hg IO 3  K sp = 1.3 × K sp = [Hg 2 2+ ][IO 3  ] 2 = x(2x) 2 → x = [Hg 2 2+ ] = 6.9 × 10  7 In M KNO 3 [Hg 2 2+ ] = 1.0 × 10  6 → 이온세기가 커지면 Hg 2 2+ 와 IO 3- 간의 인력은 순수한 물에서보다 줄어듦 → 서로 합치려는 경향이 줌 → Hg 2 (IO 3 ) 2 의 용해도가 !

3 Typical aqueous solution q1q1 q2q2

4 Examples


6 The activities of ions in solution

7 Fig The picture underlying the Debye-Hückel theory is of a tendency for anions to be found around cations, and of cations to be found around anions (one such local clustering region is shown by the circle). The ions are in ceaseless motion, and the diagram represents a snapshot of their motion. The solutions to which the theory applies are far less concentrated than shown here.

8 where A = for an aqueous solution at 25°C and I is the dimensionless ionic strength of the solution: Mean activity coefficients Debye-Hückel limiting law. extended Debye-Hückel law

9 ionic strength of the solution:

10 Fig An experimental test of the Debye-Hückel limiting law. Fig.5.35 The extended Debye-Hückel law gives agreement with experiment over a wider range of molalities.

11 The variation of the activity coefficient with ionic strength according to the extended Debye–Hückel theory. (a) The limiting law for a 1,1- electrolyte. (b) The extended law with B = 0.5. (c) The extended law, extended further by the addition of a term CI; in this case with C = 0.2. The last form of the law reproduces the observed behaviour reasonably well.

12 Calculate the ionic strength and the mean activity coefficient of 5.0 × mol kg -1 KCl(aq) at 25°C. Self-test 5.8 Calculate the ionic strength and the mean activity coefficient of 1.00 mmol kg -1 CaCl 2 (aq) at 25°C. where A = for an aqueous solution at 25°C and I is the dimensionless ionic strength of the solution: Answer: [3.00 mmol kg -1, 0.880] I = ½ (b + + b_)/b o = b/ b o where b is the molality of the solution (and b + = b_ = b). log γ ± = × (5.0 × ) 1/2 = Hence, γ ± = The experimental value is

13 migration of ions V = IR R = ρL/A κ = 1/ρ : conductivity Λ m = κ/c : molar conductivity Λ m = Λ m o – K c 1/2 Λ m o : limiting molar conductivity Λ m o : = λ + + λ −

14 Ionic conductivities Table 9.1 Ionic conductivities, λ/(mS m 2 mol −1 )

15 Ionic mobilities in water Ionic mobilities in water at 298 K, u/(10 −8 m 2 s −1 V −1 )

16 Grotthus mechanism

17 Electrochemical cells two electrodes share a common electrolyte. Figure 9.6 When the electrolytes in the electrode compartments of a cell are different, they need to be joined so that ions can travel from one compartment to another. One device for joining the two compartments is a salt bridge.

18 a galvanic cell Figure 9.7 The flow of electrons in the external circuit is from the anode of a galvanic cell, where they have been lost in the oxidation reaction, to the cathode, where they are used in the reduction reaction. Electrical neutrality is preserved in the electrolytes by the flow of cations and anions in opposite directions through the salt bridge.

19 an electrolytic cell Figure 9.8 The flow of electrons and ions in an electrolytic cell. An external supply forces electrons into the cathode, where they are used to bring about a reduction, and withdraws them from the anode, which results in an oxidation reaction at that electrode. Cations migrate towards the negatively charged cathode and anions migrate towards the positively charged anode. An electrolytic cell usually consists of a single compartment, but a number of industrial versions have two compartments.

20 a (H + ) = 1.00 H 2 (g) e-e- Pt gauze The Standard Hydrogen Electrode E o (H + /H 2 ) half-cell = V f{H 2 (g)} = 1.00

21 a silver–silver-chloride electrode Figure 9.10 The schematic structure of a silver–silver- chloride electrode (as an example of an insoluble-salt electrode). The electrode consists of metallic silver coated with a layer of silver chloride in contact with a solution containing Cl − ions.

22 a redox electrode Figure 9.11 The schematic structure of a redox electrode. The platinum metal acts as a source or sink for electrons required for the interconversion of (in this case) Fe 2+ and Fe 3+ ions in the surrounding solution.

23 A Daniell cell Figure 9.12 A Daniell cell consists of copper in contact with copper(II) sulfate solution and zinc in contact with zinc sulfate solution; the two compartments are in contact through the porous pot that contains the zinc sulfate solution. The copper electrode is the cathode and the zinc electrode is the anode.

24 measurement of cell potential Figure 9.13 The potential of a cell is measured by balancing the cell against an external potential that opposes the reaction in the cell. When there is no current flow, the external potential difference is equal to the cell potential.

25 The Nernst Equation At 25.00°C,

26 Calculating an equilibrium constant

27 A glass electrode Figure 9.14 A glass electrode has a potential that varies with the hydrogen ion concentration in the medium in which it is immersed. It consists of a thin glass membrane containing an electrolyte and a silver chloride electrode. The electrode is used in conjunction with a calomel (Hg 2 Cl 2 ) electrode that makes contact with the test solution through a salt bridge; the electrodes are normally combined into a single unit.

28 Silver/Silver Chloride Electrode mEm 1/2 E+C

29 Temperature dependence of the standard potential of a cell Figure 9.15 The variation of the standard potential of a cell with temperature depends on the standard entropy of the cell reaction.

30 Silver/Silver Chloride Electrode Ag/AgCl Electrode T (°C) E 0 (V vs SH E) E 0 (V) = x10 -4 t x10 -6 t x (t+273) for 25 < t < 275 °C

31 Thermodynamic properties and or

32 Thermodynamic properties E 0 = V F = C mol -1

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