Electrochemistry objectives

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1 Electrochemistry objectives 1) Understand how a voltaic and electrolytic cell work 2) Be able to tell which substance is being oxidized and reduced and where it is occuring the anode or cathode 3) Students will be able to read line notation 4) Be able to calculate the change in free energy in a cell reaction to tell if a reaction is spontaneous or not 5) Be able to use Nernst equation to solve problems 6) Be able to determine how long currents must be applied to electroplate a particular amount 7) Be able to order oxidizing ability based on standard reduction potentials

2 A Summary of Redox Terminology Fig. 21.1

3 An Overview of Electrochemical Cells Ø There are two types of electrochemical cells based upon the general thermodynamic nature of the reaction: Ø 1) A voltaic cell (or galvanic cell) uses a spontaneous reaction to generate electrical energy. The reacting system does work on the surroundings. All batteries contain voltaic cells. Ø 2) An electrolytic cell uses electrical energy to drive a nonspontaneous reaction (ΔG > 0), the surroundings do work on the reacting system.

4 Galvanic (voltaic) cells Ø = device in which chemical energy is changed to electrical energy Ø The oxidizing agent and reducing agent are separated, requiring that electron transfer occur through a wire Ø The current produced in the wire can then be directed elsewhere to provide useful work

5 General Characteristics of Voltaic and Electrolytic Cells

6 Electrochemical cells have several common features: Ø 1) They have two electrodes: Ø Anode The oxidation half-reaction takes place at the anode. Ø Cathode The reduction half-reaction takes place at the cathode The electrodes are dipped into an electrolyte, a solution that contains a mixture of ions and will conduct electricity.

7 MnO 4 - is an extremely strong oxidizer, why won t this produce a current?

8 What does a salt bridge (or porous disk) do? Allows ions to flow between the two solutions to keep the net charge zero Must contain an ionic substance that does not react with either half-cell projectfolder/animations/cuzncell.html

9 A Voltaic Cell based on the Zinc-Copper Reaction Active and inactive electrodes.

10 Notation for a Voltaic Cell Ø There is a shorthand notation for describing the components of a voltaic cell. For example, the notation for the Zn/Cu 2+ cell is: Ø Zn (s) Zn 2+ (aq) Cu2+ (aq) Cu (s) Ø Examples: Draw the diagram for the voltaic cell represented by: Fe (s) Fe 2+ (aq) Cu 2+ (aq) Cu (s) Write the half reactions and overall reaction for this cell.

11 Electrochemical cll Fe (s) Fe 2+ (aq) Cu2+ (aq) Cu (s)

12 Complete descriptions of galvanic cells include: Ø Direction of the electron flow (look at ½ rxn) Ø Designation of anode & cathode (anodes oxidize ) Ø Nature of each electrode and the ions present in each compartment Ø A chemically inert conductor or ions (salt bridge or porous disk Ø Cell potential and balanced cell reaction (+ for E o cell = E o (cathode) E o (anode) )

13 Cell Voltage Ø Electromotive force: emf Ø driving force of voltaic cell Ø max. potential difference between electrodes Ø depends on: nature of reaction, conc., and temperature of cell Ø Symbol: E Ø Units: volts Ø 1 volt = 1 J/coul or J = coul x volts Ø Standard emf - E o [tables 25 o C, 1 atm, 1 M

14 Standard cell potentials 25C, 1 M, 1 atm (unless specifically noted)

15 Effects of Conc. On Voltage Ø Qualitative Ø As cell operates: Reactants Products q reactant conc. q product conc. q EMF SPONTANEITY AND REDOX REACTIONS + EMF = SPONTANEOUS - EMF = NONSPONTANEOUS 0 EMF = EQUILIBRIUM

16 Standard Voltage Ø E cell = E cathode + (E anode ) q E cathode, q E anode, q Sign of E anode is opposite sign on table. What is the standard cell potential for? Fe (s) Fe 2+ (aq) Cu2+ (aq) Cu (s)

17 Emf and Free Energy Ø ΔG measures spontaneity (-) q emf measures of spontaneity (+) q Must be relationship ΔG = - nf E n is # moles of electrons transferred F is Faraday s constant: electrical charge of 1 mole of electrons F = 96,500 coul/mole electrons

18 The Effect of Concentration on Cell Potential, emf. Ø The relationship between cell potential and concentration is based upon the free energy and its relation to concentration. ΔG = ΔG o + RT ln Q Since ΔG is related to E cell, we substitute in their values, and divide each side by - nf, and we get the Nernst equation: E cell =E o cell -RT ln Q nf

19 E cell = E o cell -(RT/nF )(ln Q) We substitute R and F, and operate the cell at 25 o C (298 K), so we get: E cell = E cell V ln Q n E cell = E cell V log 10 Q n (at 25 o C)

20 Outcome: Be able to compare or calculate Δ G, Emf, or K using these equations.

24 Electrolysis Ø Breaking down with electricity Ø Electricity causes the chemical change Comparison of Voltaic and Electrolytic Cells Cell Type Δ G E cell Name Process Sign Voltaic < 0 > 0 Anode Oxidation - Voltaic < 0 > 0 Cathode Reduction + Electrolytic > 0 < 0 Anode Oxidation + Electrolytic > 0 < 0 Cathode Reduction -

25 A Summary Diagram for the Stoichiometry of Electrolysis Outcome: Find any value on chart given any other value. 1 Faraday = 96,500 coul / mol e - s 1 coulomb = ampere X seconds

Zn+2 (aq) + Cu (s) Oxidation: An atom, ion, or molecule releases electrons and is oxidized. The oxidation number of the atom oxidized increases.

Zn+2 (aq) + Cu (s) Oxidation: An atom, ion, or molecule releases electrons and is oxidized. The oxidation number of the atom oxidized increases. Oxidation-Reduction Page 1 The transfer of an electron from one compound to another results in the oxidation of the electron donor and the reduction of the electron acceptor. Loss of electrons (oxidation)