2. 2 Types of Electrochemical Cells 1)Voltaic Cells  Spontaneous reaction  Reaction itself creates electric current  Main concept for batteries 2)Electrolytic Cells

3. Electrolytic CellsElectrolytic Cells  Reaction is NONSPONTANEOUS  Electric current drives redox reaction  External source for electrical current—serves as “electron pump”  Electrolysis  Process of using an electric current to drive chemical reaction Ex. Electroplating

4. Electrolytic Cells (cont.)Electrolytic Cells (cont.)  Electrodes submerged into a salt solution, NO salt bridge  Anode  Picks up electrons from solution, negative ions give up electrons  Oxidation still occurs here  Cathode  Receives electrons from anode/wire  Positive ions accept electrons  Reduction still occurs here

7. Electrolyic Cells (cont.)Electrolyic Cells (cont.)  Polarities of electrodes (cathode/anode) change  Electrons pulled away from ANODE so it takes POSITIVE charge  External electric source gives excess electrons to CATHODE so it takes NEGATIVE charge  Oxidation and reduction still occur in same place  External electric source controlling electrons

8. Electric Current (I)Electric Current (I)  Amount of charge transferred through an area per second  I = q/T  I = current (A, amperes)  q = charge (C, coulombs)  T = time (seconds, s)

9. Example 1:Example 1:  How much charge goes through a wire containing 0.987A of current for 12.3 minutes?

10. Faraday’s LawFaraday’s Law  Relates amount of electric current in an electrochemical cell to the mass of a chemical substance  “amount of chemical consumed/made at electrode in electrolytic cell DIRECTLY proportional to amount of electric current through cell” (Spencer, p. 517)

11. Faraday’s constant (F)Faraday’s constant (F)  States the amount of charge (C) transferred by one mole of electrons  F = 96,500 C/mol of electrons  Stoichiometric relationship between mole of electrons and mole of chemical substance

12. Example 2:Example 2:  How many moles of electrons went through the area in Example 3?

13. Example 3:Example 3:  A current of 3.25 amperes is used to electrolyze a solution of copper (II) sulfate. How many hours will it take to deposit 12.71 grams of metallic copper?  A) 6.60; C) 1.65; E) 0.400;  B) 3.30; D) 63.5; F) 0.200 

14. Example 4:Example 4:  A steady current of 4.25 amperes is passed through an electrolytic cell until 3.55 grams of zinc are deposited. The cell contains a 2.00 M solution of Zn(NO 3 ) 2 and has a zinc anode and a platinum cathode. How long (in minutes ) does the current need to flow to obtain this deposit?

15. Example 5:Example 5:  A chromium (III) nitrate cell was hooked up in series with the above cell (1.00 Mwith a chromium anode instead of the zinc anode), and the same amount of current passed through as in #10, what weight (in grams) of chromium would plate out?

16. Homework  p. 671 #1-4