1. Electrochemistry or REDOX Unit 9

2. I. The Vocabulary of Electrochemistry A] Electrochemsitry is…… The field of chemistry studying reactions resulting from the transfer of electrons The transfer of electrons occurs through oxidation and reduction reactions

3. B. Oxidation State This is the charge assigned to an ionThis is the charge assigned to an ion –Metals = positives –Nonmetals = negatives –Polyatomic ions = TABLE E!! Within the coupound, NaCl Each element has its own oxidation number: Na +1 Cl -1

4. C. Reduction Reduction = gain of electron(s); decrease in oxidation number Also may be the decrease in oxygen or increase in hydrogen for a compound LeO says GeR OiL RiG

5. D. Oxidation Oxidation =Oxidation = –loss of electron(s) by a species; –increase in oxidation number; –Also may be the increase in oxygen within a compound. LeO says GeR OiL RiG

6. Oxidation Numbers Hints: 1.All compounds have an overall charge of ZERO! 2.Polyatomics have their assigned overall charge 3.Calculate the charge of transition metals LAST! 4.Elements = Zero charge! 5.Total charge on each side of the equation should be equal! 6.Changes in the oxidation state signal a redox reaction!! 6.Changes in the oxidation state signal a redox reaction!! Examples:

7. OXIDATION AND REDUCTION REACTIONS ALWAYS OCCUR TOGETHER!!! Oxidation and Reduction

8. E. Agents (secret agents…) Oxidizing AGENT is Reduced and takes in electrons…Oxidizing AGENT is Reduced and takes in electrons… –therefore, allowing another species to be oxidized! Reducing AGENT is Oxidized and releases electrons… –Therefore allowing another species in solution to be reduced!

9. Identify as Redox Reactions See the link below and the Regents questions at the end.See the link below and the Regents questions at the end. http://www.kentchemistry.com/li nks/Redox/redoxrxns.htmhttp://www.kentchemistry.com/li nks/Redox/redoxrxns.htmhttp://www.kentchemistry.com/li nks/Redox/redoxrxns.htmhttp://www.kentchemistry.com/li nks/Redox/redoxrxns.htm

10. II. Balancing REDOX Equations 1.Assign oxidation numbers for all atoms in the equation. 2.Identify which atoms are changing oxidation states from reactants side to products side. 3.Pull these atoms that change charges and rewrite as “half- reactions”

11. Rules… 4.Balance the ATOMS first in the half- reactions using coefficients. 5.Balancs the charges of the half- reactions. Add electrons to the more positive side! 6.Look at both half-reactions and make sure the number of electrons transferred are equal. Multiple both by a factor to have the same number of electrons total, if necessary.

12. Rules… 7.Add the two half-reactions together to create a new equation. 8.Cancel out species that are repeated on both the reactants and products side. Electrons should all cancel now!

13. Ex. Balance the following redox equation: Cu (s) + AgNO 3(aq)  Cu(NO 3 ) 2(aq) + Ag (s)

14. III. Balancing Acidic Equations (Optional…) Follow rules #1-4 from above, then the rules continue here…Follow rules #1-4 from above, then the rules continue here… 5a. Balance extra oxygens by adding H 2 O to the other side. Counterbalance the additional hydrogens by adding H+ if necessary. 6a. Balance charges on each side of the half-reactions by adding electrons to the more positive side.

15. Acidic rules… 7a. Make sure the same number of electrons transfer between reactants and products. If not, multiply by a factor to balance. 8a. Add half-reactions together, canceling any extra species remaining on both the reactants and products side of the new equation.

16. Ex. Balance the following equation using the acidic rules… Cu (s) + HNO 3(g)  Cu(NO 3 ) 2(g) + NO (g) + H 2 O (g)

17. IV. Electrochemical Cells A.Voltaic Cells redox reactions that generate electricity Spontaneous redox reactions that generate electricity a.k.a. Galvanic cellsa.k.a. Galvanic cells Consist of 2 half-reactionsConsist of 2 half-reactions More active metal will oxidize Less active metal will reduce

18. Applications of voltaic cells… A BATTERY!

19. Voltaic Cell Diagram [wet cell or Daniel cell]

20. Parts of a Voltaic cell 1.Anode Where oxidation occurs! Negative terminal e- are leaving this terminal consists of the MORE reactive METAL!

21. 2.Cathode Where Reduction occurs! Positive Terminal Electrons are used here Consists of the LESS reactive METAL!

22. 3.Half-Cells Consist of a metal in a solution of its metal ions Need 2 different half-cells: one with a more reactive metal and a second with a less reactive metal

23. 4. Salt Bridge A solution of nonreactive negative and positive ions flowing between the half-cells half-cells Purposes of the salt bridge: 1. Connect the 2 half cells 2. Complete the circuit 3. Provide ions to balance the charges within the solutions of the half cells http://www.kentchemistry.com/links/Redox /GalvanicTutorial.htm http://www.kentchemistry.com/links/Redox /GalvanicTutorial.htm

24. Complete voltaic cell

25. Voltaic cell NOTES… More reactive metal =More reactive metal = Less Reactive metal =Less Reactive metal = Anode metal will…Anode metal will… Cathode metal will…Cathode metal will… Electrons flow through the metals from…Electrons flow through the metals from… Ions flow through the salt bridge…Ions flow through the salt bridge… ANODE CATHODE DECREASE in mass INCREASE in mass Anode to Cathode (-) to Anode, (+) to Cathode

26. Voltaic Cell Diagram [wet cell or Daniel cell]

27. B. Electrolytic Cells Nonspontaneous reactions!Nonspontaneous reactions! Need a power source to help substances oxidize/reduceNeed a power source to help substances oxidize/reduce Used for electroplatingUsed for electroplating

28. Notes on Electrolytic cells The ANODE… still Loses mass Site of Oxidation ** ** Produces ions of the metal that is being plated onto an object The Cathode… still Gains mass Site of Reduction **Object being plated with a second metal

29. Power source connects (-) terminal to the object that is plated with another metalPower source connects (-) terminal to the object that is plated with another metal The (+) terminal goes to the metal replenishing ions

30. Electrochemical Cells summary Oxidation ALWAYS occurs at the ANODEOxidation ALWAYS occurs at the ANODE Reduction ALWAYS occurs at the CATHODEReduction ALWAYS occurs at the CATHODE Sign of the terminals for anode and cathode FLIP between spontaneous and nonspontaneous cellsSign of the terminals for anode and cathode FLIP between spontaneous and nonspontaneous cells