1. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 2 Oxidation-Reduction Reactions and Electrochemistry Chapter 17

3. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 3 Oxidation-Reduction Reactions Also Known As Redox or Electron Transfer reactions One or more elements change oxidation number –All single displacement, and combustion, –Some synthesis and decomposition Always have both oxidation and reduction –Split reaction into oxidation half-reaction and a reduction half-reaction –Half-reactions include electrons Oxidizing agent is reactant molecule that causes oxidation –Contains element reduced Reducing agent is reactant molecule that causes reduction –Contains the element oxidized

4. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 4 Oxidation & Reduction Oxidation is the process that occurs when –Oxidation number of an element increases –Element loses electrons –Compound adds oxygen –Compound loses hydrogen –Half-reaction has electrons as products Reduction is the process that occurs when –Oxidation number of an element decreases –Element gains electrons –Compound loses oxygen –Compound gains hydrogen –Half-reactions have electrons as reactants

5. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 5 Predicting Products of Metal + Nonmetal Reactions metal + nonmetal  ionic compound –ionic compounds always solids unless dissolved in water in the ionic compound the metal is now a cation in the ionic compound the nonmetal is now an anion to predict direct synthesis of metal + nonmetal ¬determine the charges on the cation and anion from their position on the Periodic Table ­balance the charges to get the formula of the compound ®balance the equation

6. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 6 Assigning Oxidation States ¬Uncombined elements = 0 ­Monatomic ions have an oxidation state equal to their charge –Including ions within compounds –e.g. Na = +1 in Na 2 CO 3 ®O = -2 as oxide, O = -1 as peroxide (O 2 -2 ) ¯H = +1; except in MH x, H = -1 –(where M = metal)

7. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 7 Assigning Oxidation States °Most electronegative nonmetal in a compound has a negative oxidation state assigned; base it on position on the Periodic Table –Group Number - 8 ±For a neutral compound, the sum of the oxidation states = 0 ²For a polyatomic ion, the sum of the oxidation states = the charge on the ion

8. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 8 Balancing Oxidation-Reduction Reactions by the Half-Reaction Method Balancing by Inspection can often be very difficult with redox reactions One method we use is to write separate reactions for the oxidation and reduction processes –Called half-reactions We multiply the half-reactions so the number of electrons lost in the oxidation and gained in the reduction are equal Then we simply add the half-reactions together and cancel the electrons

9. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 9 ¬Write the oxidation and reduction half-reactions ­For Each Half-Reaction –Balance all elements except hydrogen and oxygen –Balance O using H 2 O –Balance H using H + –Balance the charge using electrons ®If necessary, multiply the half-reactions to equalize the number of electrons transferred ¯Add the half-reactions and cancel identical species °Check to ensure elements and charges are balanced Balancing Oxidation-Reduction Reactions by the Half-Reaction Method in Acidic Solution

10. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 10 Example ¬Assign Oxidation States and Determine the elements that are Oxidized and Reduced. Write the Half-Reactions Balance the Following MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq) +7-2+2+3+2 Mn is reduced from +7 to +2Fe is oxidized from +2 to +3 Ox:Fe +2  Fe +3 Red:MnO 4 -1  Mn +2

11. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 11 ­Balance the Half-Reactions –First balance all the elements besides H and O –In acid solution, add H 2 O to side that is O deficient; add H +1 to side that is H deficient, then balance O first, then H –Add number of electrons transferred to product side for oxidation, and to reactant side for reduction = balance charge Example Balance the Following MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq) Ox:Fe +2  Fe +3 + 1 e -1 Red:MnO 4 -1 + 8 H +1 + 5 e -1  Mn +2 + 4 H 2 O Fe and Mn already balanced

12. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 12 ®Multiply the Half-Reactions by factors so the electrons lost in oxidation equals the electrons gained in reduction Ox:Fe +2  Fe +3 + 1 e -1 Red:MnO 4 -1 + 8 H +1 + 5 e -1  Mn +2 + 4 H 2 O } x 5 Ox:5 Fe +2  5 Fe +3 + 5 e -1 Red:MnO 4 -1 + 8 H +1 + 5 e -1  Mn +2 + 4 H 2 O Example Balance the Following MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq)

13. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 13 ¯Add the Half-Reactions, cancel identical species that appear on both sides Total: 5 Fe +2 + MnO 4 -1 + 8 H +1  5 Fe +3 + Mn +2 + 4 H 2 O Ox:5 Fe +2  5 Fe +3 + 5 e -1 Red:MnO 4 -1 + 8 H +1 + 5 e -1  Mn +2 + 4 H 2 O Total: 5 Fe +2 + MnO 4 -1 + 8 H +1 + 5 e -1  5 Fe +3 + Mn +2 + 4 H 2 O + 5 e -1 Example Balance the Following MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq)

14. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 14 Electrochemical Cells Electrochemistry is the study of the interchange of chemical and electrical energy The conversion between chemical energy and electrical energy is carried out in an electrochemical cell –Keep the oxidizing and reducing agents separated Spontaneous redox reactions take place in a galvanic cell –Also known as voltaic cells Non-spontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy Using electrical energy to decompose a compound is called electrolysis

15. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 15 Electrochemical Cells oxidation and reduction reactions kept separate –half-cells requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons –through external circuit ion exchange between the two halves of the system –electrolyte

16. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 16 Electrodes Anode –electrode where oxidation occurs –anions attracted to it –connected to positive end of battery in electrolytic cell –loses weight in electrolytic cell it contains substance oxidized Cathode –electrode where reduction occurs –cations attracted to it –connected to negative end of battery in electrolytic cell –gains weight in electrolytic cell it contains substance reduced electrode where plating takes place in electroplating Electrode surface area dictates the number of electrons that will flow, therefore the current

17. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 17 Galvanic Cells Oxidation occurs at the anode Spontaneous reaction means the anode metal more reactive –metal (or ion species) with most negative E 0 red Cathode reaction involves material with high oxidation state being reduced Charge balance must be maintained –as electrons are lost from anode, replaced by anions –as cations are reduced from cathode, replaced by cations –in “wet” cell need a salt bridge to “complete circuit” –“dry” cells can “come back to life” after charge balance re- established

18. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 18 Galvanic Cells produce an electric current –batteries –size indicates amount of current primary cells are batteries that run down and can be used only the one time –acidic dry cell secondary cells are batteries that can be recharged –alkaline dry cell, Pb storage, NiCad, watch batteries corrosion cathodic protection –sacrificial anode anode anode electrolyte cathode electrolyte cathode

19. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 19 Lead Storage Battery 6 cells connected together electrolyte = 6 M H 2 SO 4 anode = Pb Pb(s) + H 2 SO 4 (aq)  PbSO 4 (s) + 2 H + + 2 e - cathode = Pb coated with PbO 2 PbO 2 is reduced PbO 2 + H 2 SO 4 + 2 H + + 2 e -  PbSO 4 (s) + 2 H 2 O(l) cell voltage = 2.04 v rechargeable, heavy

20. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 20 LeClanche’ Acidic Dry Cell electrolyte in paste form –ZnCl 2 + NH 4 Cl or MgBr 2 anode = Zn (or Mg) Zn(s)  Zn +2 (aq) + 2 e - cathode = graphite rod MnO 2 is reduced 2 MnO 2 + 2 NH 4 + + 2 e -  Mn 2 O 3 + 2 NH 3 + 2 H 2 O cell voltage = 1.5 v expensive, non-rechargable, heavy, easily corroded

21. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 21 Alkaline Dry Cell same basic cell as acidic dry cell, except electrolyte is alkaline –KOH paste anode = Zn (or Mg) Zn(s)  Zn +2 (aq) + 2 e - cathode = brass rod MnO 2 is reduced 2 MnO 2 + 2 NH 4 + + 2 e -  Mn 2 O 3 + 2 NH 3 + 2 H 2 O cell voltage = 1.54 v longer shelf life than acidic dry cells and rechargable, little corrosion of zinc

22. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 22 NiCad Battery electrolyte is concentrated KOH solution anode = Cd Cd(s) + 2 OH - (aq)  Cd(OH) 2 (s) + 2 e - cathode = Ni coated with NiO 2 NiO 2 is reduced NiO 2 (s) + 2 H 2 O(l) + 2 e -  Ni(OH) 2 (s) + 2OH - cell voltage = 1.30 v rechargeable, long life, light

23. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 23 Corrosion Corrosion = spontaneous process in which metals are oxidized by the environment Most metals are found as ores in nature –Generally compounds of metal and oxygen or sulfur –Exceptions being the noble metals Ag, Au and Pt Some metals, even though very reactive, do not corrode quickly because their surface gets coated with a tightly held layer of the metal oxide or sulfide Oxide coat –Metals that corrode quickly form oxides that do not adhere well to the surface

24. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 24 Preventing Corrosion Coat the metal surface to keep it from exposure to the environment –With paint –With a metal that forms a strong oxide coat Chrome plating, tin plating Alloy with metals that do not corrode easily Attach a piece of metal that corrodes easier than the metal you need to protect –Cathodic Protection –Sacrificial Anode Galvanizing = coating metal with zinc, which acts as both a coating and sacrificial anode

25. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 25 Electrolysis Electrolysis = using electrical energy to break a compound into its elements –Separate H 2 O into H 2 and O 2 Cause a non-spontaneous reaction to occur Often used to separate metals from ores Requires direct current