by Steven S. Zumdahl & Donald J. DeCoste University of Illinois Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed.

Chapter 18 Oxidation-Reduction Reactions & Electrochemistry

Oxidation-Reduction Reactions Section 18-1

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 4 Oxidation-Reduction Reactions Also known as redox or electron transfer reactions One or more elements change oxidation number –All single displacement and combustion reactions –Some synthesis and decomposition reactions

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 5 Oxidation-Reduction Reactions (cont.) Always have both oxidation and reduction –Split reaction into oxidation half-reaction and a reduction half-reaction –Half-reactions include electrons

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 6 Section 18-2 Oxidation States

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 7 Oxidation States Also called oxidation numbers Lets us keep track of electrons in redox reactions by assigning charges to various atoms in a compound. Sometimes oxidation states are apparent. Sometimes you just need to remember the rules!!

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 8 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)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 9 Assigning Oxidation States Most electronegative nonmetals in compounds are assigned negative oxidation states based on their positions on the periodic table. –Group number - 8 For a neutral compound, the sum of the oxidation states = 0 For a polyatomic ion, the sum of the oxidation states = the charge on the ion

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 10 Specific Rules Element(s)Oxidation #Exceptions Group I+1None Group II+2None FNone H (with metals and B)None H (with non-metals)+1None O-2-1 in peroxides or with F Halogenswith oxygen or halogens higher in column

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 11 Section 18-3 Oxidation-Reduction Reactions Between Non- metals

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 12 A bad mnemonic… LEO the lion goes GER

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 13 Oxidation-Reduction Reactions (cont.) Oxidizing agent: reactant molecule that causes oxidation –Contains element reduced Reducing agent: reactant molecule that causes reduction –Contains the element oxidized

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 14 Oxidation Oxidation is the process that occurs when: 1. Oxidation number of an element increases 2. Element loses electrons 3. Compound adds oxygen 4. Compound loses hydrogen 5. Half-reaction has electrons as products

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 15 Reduction Reduction is the process that occurs when: 1. Oxidation number of an element decreases 2. Element gains electrons 3. Compound loses oxygen 4. Compound gains hydrogen 5. Half-reactions have electrons as reactants

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 16 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.

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 17 Predicting Products of Metal + Nonmetal Reactions (cont.) To predict direct synthesis of metal + nonmetal: 1. Determine the charges on the cation and anion from their positions on the periodic table 2. Balance the charges to get the formula of the compound 3. Balance the equation

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 18 Section 18-4 Balancing Oxidation-Reduction Reactions by the Half Reaction Method

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 19 Balancing Oxidation-Reduction Reactions: the Half-Reaction Method Balancing by inspection can often be very difficult with redox reactions. One method is to write separate reactions for the oxidation and reduction processes. –Called half-reactions Multiply the half-reactions so the number of electrons lost in the oxidation and gained in the reduction is equal. Add the half-reactions together and cancel the electrons.

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 20 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 Balancing Oxidation-Reduction Reactions: the Half-Reaction Method in Acidic Solution

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 21 Balancing Oxidation-Reduction Reactions: the Half-Reaction Method in Acidic Solution (cont.) 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.

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 22 Example #1 Balance the following MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 23 Example #1 (cont.) Assign oxidation states and determine the elements that are oxidized and reduced. Write the half-reactions: MnO 4 -1 (aq) + Fe +2 (aq)  Fe +3 (aq) + Mn +2 (aq) Mn is reduced from +7 to +2Fe is oxidized from +2 to +3 Ox:Fe +2  Fe +3 Red:MnO 4 -1  Mn +2

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 24 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 + 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 #1 (cont.) Red:MnO H e -1  Mn H 2 O Ox:Fe +2  Fe e -1 Fe and Mn already balanced

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 25 Multiply the half-reactions by factors so the electrons lost in oxidation equal the electrons gained in reduction: Ox:Fe +2  Fe e - Red:MnO H e -  Mn H 2 O } x 5 Ox:5 Fe +2  5 Fe e - Red:MnO H e -  Mn H 2 O Example #1 (cont.)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 26 Add the half-reactions and cancel identical species that appear on both sides Total: 5 Fe +2 + MnO H +  5 Fe +3 + Mn H 2 O Ox:5 Fe +2  5 Fe e - Red:MnO H e -  Mn H 2 O Total: 5 Fe +2 + MnO H e -  5 Fe +3 + Mn H 2 O + 5 e - Example #1 (cont.)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 27 Electrochemical Cells Electrochemistry: 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 so electron transfer occurs through wires Spontaneous redox reactions take place in a galvanic cell (voltaic cell).

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 28 Electrolysis Non-spontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy. Electrolysis: use of electrical energy to decompose a compound

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 29 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 30 Electrodes Anode –Electrode where oxidation occurs –Anions attracted to it Cathode –Electrode where reduction occurs –Cations attracted to it

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 31 Electrodes (cont.)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 32 Galvanic Cells Spontaneous reaction that generates electricity Cathode reaction involves material with high oxidation state being reduced

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 33 Galvanic Cells (cont.) 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 34 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 Galvanic Cells (cont.)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 35 Galvanic Cells (cont.) Corrosion Cathodic protection –Sacrificial anode anode anode electrolyte cathode electrolyte cathode

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 36 Lead Storage Battery 6 cells connected together PbO 2 + H 2 SO H e -  PbSO 4 (s) + 2 H 2 O(l) Cell voltage = 2.04 v Rechargeable, heavy

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 37 LeClanche’s Acidic Dry Cell Electrolyte in paste form –ZnCl 2 + NH 4 Cl or MgBr 2 2 MnO NH e -  Mn 2 O NH H 2 O Cell voltage = 1.5 v Expensive, non-rechargable, heavy, easily corroded

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 38 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 39 Alkaline Dry Cell (cont.) MnO 2 is reduced 2 MnO NH e -  Mn 2 O NH H 2 O Cell voltage = 1.54 v Longer shelf life than acidic dry cells and rechargeable; little corrosion of zinc

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 40 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 41 NiCad Battery (cont.) 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 42 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 43 Corrosion (cont.) 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

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 44 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, e.g. chrome plating, tin plating Alloy with metals that do not corrode easily Cathodic protection Galvanizing: coating metal with zinc, which acts as both a coating and sacrificial anode

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 45 Preventing Corrosion (cont.)

Copyright © Houghton Mifflin Company. All rights reserved. 18 | 46 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