Electrochemical System

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Electrochemical System Topic Outcomes Week Topic Topic Outcomes 8-10 Electrochemical systems It is expected that students are able to: Electrochemical system and its thermodynamics Chemical reactions

1 Electrochemical System Topic Outcomes Week Topic Topic Outcomes 8-10 Electrochemical systems It is expected that students are able to: Electrochemical system and its thermodynamics Chemical reactions in electrochemical cells and the Nernst equation The relationship between the cell emf and the equilibrium constant, Gibbs energies and reaction entropies Thermodynamics of Galvanic cells and Fuel cell Define the electrochemical system. Explain about electrochemical cell and differentiate galvanic cell and electrolyte cell Apply Nernst equation of thermodynamic principles to electrochemical cells. Determine G o, K o, S o, H o, Cp o activity coefficient and ph of cell s chemical reaction. Understand the cell emf in IUPAC convention. Estimate the liquid junction potential from emf measurement..

Quick Review Redox: Reduction-Oxidation, an electron transfer process Oxidation: loss one or more e - Reduction: gain one or more e - Oxidation number: the charge of the atom or molecule would have

2 Quick Review Redox: Reduction-Oxidation, an electron transfer process Oxidation: loss one or more e - Reduction: gain one or more e - Oxidation number: the charge of the atom or molecule would have Pure element and neutral compound 0 Mono- and poly-atomic ions sum of oxidation number is ionic charge Check your understanding CaBr 2 ; Ca = +2, Br = -1, Oxidation number (pure element), =(+2)+2(-1)=0 HCO 3 - O = -2 H = +1 C =? K 2 Cr 2 O 7 O = -2 K = +1 Cr =?

4 Quick Review (Cont d) Electrochemistry: study of the interchange between chemical change and electrical work Electrochemical cells: systems utilizing a redox reaction to produce or use electrical energy

bridge: Maintains electrical neutrality, allows current to flow, prevent solutions mixing

5 Becomes smaller Becomes larger Zn anode Cu cathode Cell and Cell Notation Anode: Electrode site of oxidation Cathode: Electrode site of reduction Electrons flow from anode to cathode Salt bridge: Maintains electrical neutrality, allows current to flow, prevent solutions mixing Positive ion migrates to cathode Negative ion migrates to anode Cell notation: Anode Salt Bridge Cathode

6 The difference in electrical potential between the anode and cathode is called: cell voltage electromotive force (emf) cell potential E + 0 Cell 0 0 = Eoxidation E (Unit: Volts) reduction Note: Volt (V) = Joule (J) Coulomb (C) Cell Potential

8 Standard Electrode Potentials Standard reduction potential (E 0 ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. Reduction Reaction 2e - + 2H + (1 M) H 2 (1 atm) E 0 = 0 V Reference electrode Standard hydrogen electrode (SHE)

34 (If copper is reduced (at cathode), voltmeter

9 (Cathode) (Anode) Note: Copper reduction potential, E o = (If copper is reduced (at cathode), voltmeter can give a reading- positive) Note: Zn reduction potential is negative. Change Zn as anode

E 0 is for the reaction as written The half-cell reactions are reversible; the sign of E 0 changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does

10 E 0 is for the reaction as written The half-cell reactions are reversible; the sign of E 0 changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0 The more positive E 0 the greater the tendency for the substance to be reduced Spontaneity of Redox Reaction E = E + 0 Cell 0 oxidation E 0 reduction E o reduction=-e o oxidation + E CELL : spontaneous reaction E CELL = 0 :equilibrium -E CELL : non-spontaneous reaction More positive E CELL means stronger oxidizing agent or more likely to be reduced

11 Balance the Half-Cell Reactions Write separate equations (half-reactions) for oxidation and reduction For each half-reaction Balance elements involved in e - transfer Balance number e - lost and gained To balance e - multiply each half-reaction by whole numbers Add half-reactions/cancel like terms (e - ) Check that all atoms and charges balance Check your understanding?? e e e e Zn(s) Zn 2+ SO M Zn 2+ solution Cu 2+ SO M Cu 2+ solution Cu(s)?? Anode Cathode??

12 Check your understanding A half-cell reduction reaction is Fe(OH) 2 (s) + 2e - Fe (s) + 2OH - (aq) E o = V Is to combine with another half-cell, which the reduction reaction is given as: a) Al 3+ (aq) + 3e - Al (s) E o = V b) AgBr (s) + e - Ag (s) + Br - (aq) E o = V Determine E o cell, and whether Fe is oxidized or reduced. Nernst Equation G = G 0 + RTlnQ G 0 = -nfe 0 cell the emf for an electrochemical cell to be calculated if the activity and E are known. -nfe cell =-nfe 0 cell +RTlnQ E cell =E 0 cell - RTln Q nf E cell = E 0 cell ln Q n E cell =E 0 cell log Q n Note: At equilibrium, G=0, thus E cell =0

G o = -nfe o cell Faraday, F: charge on 1 mole e - F = 96485 C/mole n= moles of balanced e - ΔG 0 = -RT ln K -RT ln K = -nfe 0 cell

13 G o = -nfe o cell Faraday, F: charge on 1 mole e - F = C/mole n= moles of balanced e - ΔG 0 = -RT ln K -RT ln K = -nfe 0 cell ΔG 0 = -nfe 0 cell 0 E Cell = RT nf ln K RT F J 8.31 molk = C mole ( 298K ) = E Cell = ln K n = logk n

14 Entropy Change The reaction entropy related to ΔG R pprrtenftgs = Δ = Δ o o o Measurement of the T dependence of E can be used to determine ΔS R To obtain the entropy change for the cell reaction The reaction entropy related to ΔG R PTEnFTnFEH + = Δ o o o Enthalpy Change H o = G o + T S o

15 Example (from text book) The half-cell reduction reactions, Fe 3+ (aq) + e - Fe 2+ (aq) Fe 2+ (aq) + 2e - Fe (s) E o = V E o = V Calculate E o for half-cell reduction reaction Fe 3+ (aq) + 3e - Fe (s) Example Calculate ΔG o r and the equilibrium constant at K for two half-cell reduction reactions: Cr 2 O 7 2- (aq) + 14H + (aq) + 6e - 2Cr 3+ (aq) + 7H 2 O(l) Eº = V 2H + (aq) + 2e - H 2 (g) Eº = 0.00 V

16 Example The standard potential E for a given cell in which a two electron transfer takes place is 1.10 V at K and ( E / T) = V K 1. Calculate ΔG rxn, ΔS rxn, ΔH rxn. Assume that n = 2. Check your understanding Calculate the equilibrium constant at K for the following half-cell reactions a) Cathode: NO H + + 3e - NO + 2H 2 O Eº = V Anode: 2H 2 O O 2 + 4H + + 4e - Eº = V b) Anode: Cd(OH) 2 + 2e - Cd + 2OH - Eº = 0.809V Cathode: O 2 + 2H 2 O + 4e - 4OH - Eº = V

17 Example Calculate the equilibrium constant at 25 C for the reaction occurring in the Daniell cell, if the standard EMF is 1.10 V. Calculate the equilibrium constant for the reaction H 2 + 2Fe 3+ 2H + + 2Fe 2+ Galvanic cell vs. Electrolytic cell

19 Galvanic cell vs. Electrolytic cell Electrochemical cell (Galvanic Cell) Converts chemical energy into electrical energy Redox reaction is spontaneous and is responsible for the production of electrical energy The two half-cells are set up in different containers, being connected through the salt bridge or porous partition Anode is negative and cathode is the positive electrode. The reaction at the anode is oxidation and that at the cathode is reduction. The electrons are supplied by the species getting oxidized. They move from anode to the cathode in the external circuit. Electrolytic cell Converts electrical energy into chemical energy Redox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction Both the electrodes are placed in a same container in the solution of molten electrolyte Here, the anode is positive and cathode is the negative electrode. The reaction at the anode is oxidation and that at the cathode is reduction. The external battery supplies the electrons. They enter through the cathode and come out through the anode.

17.1 Redox Chemistry Revisited

Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6