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Chapter 20
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Chapter 20

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  • 1. Chapter 20Oxidation-Reduction Reactions
  • 2. Some Common ReactionsThe combustion of gasoline in an automobile engine requires oxygenBurning of wood in a fireplace requires oxygenThe reactions that break down food in your body and release energy use oxygenThe oxide of hydrogen is waterCharcoal oxidizes when it burn forming CO2Bleaching stains in fabric is still oxidation even though it does not burn.
  • 3. OxidationWhen methane burns in air, it oxidizes and forms oxides of carbon and hydrogen. CH4(g) + 2O2(g) CO2(g) + 2H2O(g)When elemental iron turns to rust, it slowly oxidizes to compounds such as iron (III) oxide. 4Fe(s) + 3O2(g) 2Fe2O3(s)
  • 4. ReductionOriginally reduction meant a loss of oxygen from a compound 2Fe2O3(s) + 3C(s) 4Fe(s) + 3CO2(g) iron oxide carbon iron carbon dioxideReduction of iron ore to metallic iron involves the removal of oxygen from iron (III) oxide.Involves heating the ore with carbon.
  • 5. QuestionWhat happens to magnesium and oxygen when they react to form magnesium oxide? +2 -2 2Mg + O2 2MgO magnesium oxygen magnesium oxide Magnesium loses electrons to form Mg2+ Oxygen gains electrons to form O2-
  • 6. Electron Shift in Redox ReactionsThe modern concept of oxidation and reduction have been extended to include many reactions that do not even involve oxygen.Oxygen is the most electronegative element (besides fluorine)When oxygen bonds with an atom of a different element (except fluorine), electrons from that atom shift toward oxygen.
  • 7. Redox ReactionsRedox reactions are currently understood to involve any shift of electron between reactants.Oxidation – a process that involves a complete or partial loss of electrons or a gain of oxygen.• Results in an increase in the oxidation number of an atomReduction – a process that involves a complete or partial gain of electrons or the loss of oxygen.• Results in a decrease in the oxidation number of an atom
  • 8. Redox Reactions Oxidation and reduction always occur simultaneously.The substance gaining oxygen or losing electrons is oxidizedThe substance losing oxygen or gaining electrons is reduced.
  • 9. Redox ReactionsDuring a reaction between a metal and a nonmetal, electrons are transferred from atoms of the metal to atoms of the nonmetal. Mg + S Mg2+S2- magnesium sulfur magnesium sulfide2 electrons are transferred from a magnesium atom to a sulfur atom.Magnesium atoms become more stable by the loss of electrons. Sulfur atoms become more stable by the gain of electrons
  • 10. Redox Reactions Mg + S Mg2+S2- magnesium sulfur magnesium sulfideOxidation: Mg Mg2+ + 2e- (loss of electrons)Reduction: S + 2e- S2- (gain of electrons)Magnesium atom is said to be oxidized to a magnesium ionSulfur atom is said to be reduced to a sulfide ion.
  • 11. Redox ReactionsWhen a metal combines with oxygen, it loses electronsWhen oxygen is removed from the oxide of a metal, the metal gains electrons.This knowledge is what led to the broader definition of oxidation and reduction as an exchange of electrons.
  • 12. Redox ReactionsReducing agent – the substance that loses the electrons Mg + S MgS reducing agent oxidizedOxidizing agent – the substance that accepts electrons Mg + S MgS oxidizing agent reduced
  • 13. Sample ProblemSilver nitrate reacts with copper to form copper nitrate and silver. From the equation below, determine what is oxidized and what is reduced. Identify the oxidizing agent and the reducing agent. 2AgNO3 + Cu Cu(NO3)2 + 2Ag Rewrite the equation in ionic form 2Ag+ + 2NO3- + Cu Cu2+ + 2NO3- + 2Ag
  • 14. Sample Problem 2Ag+ + 2NO3- + Cu Cu2+ + 2NO3- + 2AgOxidation: 2 e- are lost from copper when it becomes a Cu2+Reduction: 2e- are gained by two silver ions which become neutral silver atoms.2Ag+ + 2NO3- + Cu Cu2+ + 2NO3- + 2Agoxidizing reducingAgent agentreduced oxidized
  • 15. Sample Problem 2Na + S 2Na+ + S2- oxidized reduced reducing agent oxidizing agent4Al + 3O2 4Al3+ + 3O2-Oxidized reducedReducing agent oxidizing agent 2I- I2 + 2e- oxidation Zn2+ + 2e- Zn reduction
  • 16. Redox with Covalent Compounds Some reactions involve covalent compounds. In these compounds complete electron transfer does not occur. 2H2 (g) + O2 (g) 2H2O (l)In each reactant hydrogen molecule, the bonding electrons are shared equally between the hydrogen atoms.In water, however, the bonding electrons are pulled toward oxygen because it is much more electronegative than hydrogen.
  • 17. Redox with Covalent Compounds 2H2 (g) + O2 (g) 2H2O (l)There is a shift of bonding electrons away from hydrogen, even though there is not a complete transfer.Hydrogen is oxidized because it undergoes a partial loss of electrons.
  • 18. Redox with Covalent Compounds 2H2 (g) + O2 (g) 2H2O (l)In oxygen, the bonding electrons are share equally between oxygen atoms in the reactant oxygen molecule.When oxygen bonds to hydrogen in the water molecule, there is a shift of electrons toward oxygen.Oxygen is thus reduced because it undergoes a partial gain of electrons.
  • 19. Redox with Covalent CompoundsH H O O H O shift of bondinge- shared e- shared e- away from Hequally equally H and toward OH is reducing O is oxidizing agent agent
  • 20. Redox with Covalent CompoundsSome reactions involving covalent reactants or products, the partial electron shifts are less obvious.General guideline for covalent reactants or products:• for carbon compounds, the addition of oxygen or the removal of hydrogen is always oxidation
  • 21. Processes Leading to Oxidation & ReductionProcesses Leading to Oxidation & Reduction Oxidation ReductionComplete loss of electrons Complete gain of electrons(ionic reactions) (ionic reactions)Shift of electrons away from Shift of electrons toward anan atom in a covalent bond atom in a covalent bondGain of oxygen Loss of oxygenLoss of hydrogen by a Gain of hydrogen by acovalent compound covalent compoundIncrease in oxidation Decrease in oxidationnumber number
  • 22. CorrosionIron, often used in theform of the alloy steel,corrodes by beingoxidized to ions of ironby oxygen. 2Fe(s) + O2 (g) + 2H2O (l) 2Fe(OH)2 (s)4Fe(OH)2(s) + O2 (g) + 2H2O (l) 4Fe(OH)3 (s)• Equations describe the corrosion of iron to iron hydroxides in moist conditions.
  • 23. CorrosionWater in the environment accelerates the rate of corrosion.Corrosion occurs more rapidly in the presence of salts and acids.Salts and acids produce electrically conducting solutions that make electron transfer easier.
  • 24. Resistance to Corrosion Not all metals corrode easily.Gold and platinum are called noble metals because they are very resistant to losing their e- by corrosion.Other metals lose electrons easily but are protected from extensive corrosion by the oxide coating formed on their surface.• Aluminum oxidizes quickly in air to form a coating of very tightly packed aluminum oxide particles.
  • 25. Resistance to CorrosionIron also forms a coating when it corrodesBut the coating of iron oxide that forms is not tightly packed.Water and air can penetrate the coating and attack the iron metal below it.Corrosion continues until the iron object becomes only a pile of rust.
  • 26. Controlling CorrosionTo prevent corrosion, the metal surface can be coated with oil, paint, plastic or another metal.Coatings exclude air and water from the surface, preventing corrosion.If coating is scratched or worn away, however, the exposed metal will begin to corrode.
  • 27. Controlling Corrosion Another method of corrosion controlOne metal is “sacrificed” or allowed to corrode, in order to save a second metal.To protect an iron object, a piece of magnesium (or other active metal) may be placed in electrical contact with the iron.When oxygen and water attack the iron object, the iron atoms lose electrons as the iron being to be oxidized.
  • 28. Controlling CorrosionBecause magnesium is a better reducing agent than iron and is more easily oxidizedthe magnesium immediately transfers electrons to the iron, preventing their oxidation to iron ions.Magnesium is “sacrificed” by oxidation and protects the iron in the process.
  • 29. Controlling CorrosionSacrificial zinc and magnesium blocks are sometimes attached to piers and ship hulls to prevent corrosion damage in areas submerged in water.Underground pipelines and storage tanks may be connected to magnesium block for protectionIt is easier and cheaper to replace a block of magnesium or zinc than to replace a bride or a pipeline.
  • 30. Question Can you identify the common chemical characteristic of all metal corrosion?The transfer of electrons from metals to oxidizing agents.
  • 31. QuestionsDefine oxidation and reduction in terms of the gain or loss of oxygen. Oxidation is the gain of oxygen Reduction is the loss of oxygenDefine oxidation and reduction in terms of the gain or loss of electrons. LEO the lion goes GER Loss of Electrons is Oxidation Gain of Electrons is Reduction
  • 32. QuestionsWhat happens to the atoms in an iron nail that corrodes? Iron atoms are oxidized when iron corrodesHow do you identify the oxidizing agent and the reducing agent in a redox reaction? The species reduced is the oxidizing agent. The species oxidized is the reducing agent.
  • 33. QuestionsUse electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) 2Na(s) + Br2(l) 2NaBr(s) Na oxidized, Br2 reduced H2(g) + Cl2(g) 2HCl(g) H2 oxidized, Cl2 reduced
  • 34. QuestionsUse electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) 2Li(s) + F2(g) 2LiF(s) Li oxidized, F2 reduced S(s) + Cl2(g) SCl2(g) S oxidized, Cl2 reduced
  • 35. QuestionsUse electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) N2(g) + 2O2(g) 2NO2(s) N2 oxidized, O2 reduced Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s) Mg oxidized, Cu reduced
  • 36. QuestionsIdentify the reducing agent and the oxidizing agent for each reaction. 2Na(s) + Br2(l) 2NaBr(s) Na reducing agent, Br2 oxidizing agent H2(g) + Cl2(g) 2HCl(g) H2 reducing agent, Cl2 oxidizing agent
  • 37. QuestionsIdentify the reducing agent and the oxidizing agent for each reaction. 2Li(s) + F2(g) 2LiF(s) Li reducing agent, F2 oxidizing agent S(s) + Cl2(g) SCl2(g) S reducing agent, Cl2 oxidizing agent
  • 38. QuestionsIdentify the reducing agent and the oxidizing agent for each reaction. N2(g) + 2O2(g) 2NO2(s) N2 reducing agent, O2 oxidizing agent Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s) Mg reducing agent, Cu oxidizing agent
  • 39. End of Section 20.1
  • 40. Oxidation NumbersOxidation number is a + or – number assigned to an atom to indicate its degree of oxidation or reduction.General Rule A bonded atom’s oxidation # is the charge that it would have if the e- in the bond were assigned to the atom of the more electronegative element.
  • 41. Rules for Assigning Oxidation Numbers1. The oxidation number of a monatomic ion is = in magnitude and sign to its ionic charge. Bromide (Br1-) is -1 Iron III (Fe3+) is +32. The oxidation number of hydrogen in a compound is +1, except in metal hydrides, such as NaH, where it is -13. The oxidation number of oxygen in a compound is -2 except in peroxides, such as H2O2, where it is -1 and in compounds with the more electronegative fluorine, where it is positive.
  • 42. Rules for Assigning Oxidation Numbers4. The oxidation number of an atom in uncombined (elemental) form is 0. Potassium metal (K) is 0 Nitrogen Gas (N2) is 05. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 06. For polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.
  • 43. Some ThoughtDetermining oxidation numbers of elements in compounds is a way for chemists to keep track of electron transfer during redox reactions What are other examples where items are numbered to keep track of movement? The numbers on a sports player’s jerseyThe area codes assigned to telephone numbers in different regions
  • 44. Binary Ionic CompoundsIn binary ionic compounds, such as NaCl and CaCl2, the oxidation numbers of the atoms equal their ionic charges. Na1+ + Cl-1 NaCl oxidation # +1 -1 neutral Ca2+ + Cl-1 CaCl2 oxidation # +2 -1 neutral Note the sign I put before the oxidation number
  • 45. Molecular CompoundsNo ionic charges are associated with atoms of molecular compounds.However, oxygen is reduced in the formation of water for example.In water the two shared e- in the H – O bond are shifted toward the O and away from the H.Imagine the e- contributed by the two H atoms are completely transferred to the O.
  • 46. Molecular CompoundsThe charge that would result from the transfer are the oxidation numbers of the bonded elements.The oxidation number of O is -2 and the oxidation number of each hydrogen is +1Oxidation numbers are often written above the chemical symbols in a formula. +1 -2 H2O
  • 47. Multiple Oxidation NumbersMany elements can have several different oxidation numbers. +1 +6 -2 K2CrO4 – Potassium Chromate +1 +12 -2 K2CrO7 – Potassium Dichromate
  • 48. Sample ProblemsWhat is the oxidation number of each kind of atom in the following ions and compounds? +4 -2 SO2 +4 -2 CO32- +1 +6 -2 Na2SO4 -3 +1 -2 (NH4)2S
  • 49. Sample ProblemsDetermine the oxidation number of each element in the following. +3 -2 S2O3 +1 -1 Na2O2 +5 -2 P2O5 +5 -2 NO3-
  • 50. Sample ProblemsDetermine the oxidation number of chlorine in each of the following substances. +1 +5 -2 KClO3 0 Cl2 +2 +7 -2 Ca(ClO4)2 +1 -2 Cl2O
  • 51. Sample ProblemsWhat are the oxidation numbers of iodine in the following? +1 +7 -2 HIO4 +1 +5 -2 HIO3 +1 +1 -2 HIO 0 I2 +1 -1 HI
  • 52. Oxidation Number ChangesWhen copper wire is placed in a solution of silver nitrate the following reaction occurs +1 +5 -3 0 +2 +5 -2 0 2AgNO3(aq) + Cu(s) Cu(NO3)2(aq) + 2Ag(s) Ag is reduced from +1 to 0 Cu is oxidized from 0 to +2
  • 53. Oxidation Number ChangesAn increase in the oxidation number of an atom or ion indicates oxidation.A decrease in the oxidation number of an atom or ion indicates reduction. +2 +6 -2 0 0 +2 +6 -2 2CuSO4(aq) + Fe2(s) 2Cu(s) + 2FeSO4 Cu is reduced from +2 to 0 Fe is oxidized from 0 to +2
  • 54. Sample ProblemIdentify which atoms are oxidized and which are reduced in the following reaction. Also identify the oxidizing agent and the reducing agent. +1 -1 0 +1 -1 0 2HBr(aq) + Cl2(g) 2HCl(aq) + Br2(l)Cl is reduced from 0 to -1, so Cl2 is the oxidizing agentBr is oxidized from -1 to 0, so Br1- is the reducing agent
  • 55. Sample ProblemIdentify which atoms are oxidized and which are reduced in each reaction 0 0 +1 -2 O2(g) + 2H2(g) 2H2O(l) O2 is reduced from 0 to -2 H2 is oxidized from 0 to +1
  • 56. Sample ProblemIdentify which atoms are oxidized and which are reduced in each reaction +1 +5 -2 +1 +3 -2 0 2KNO3(s) 2KNO2(s) + O2(g) N is reduced from +5 to +3 O is oxidized from -2 to 0
  • 57. Sample ProblemIdentify the oxidizing agent and the reducing agent in each equation. 0 0 +1 -2 O2(g) + 2H2(g) 2H2O(l) O2 is reduced, thus O2 is the oxidizing agent H2 is oxidized, thus H2 is the reducing agent
  • 58. Sample ProblemIdentify the oxidizing agent and the reducing agent in each equation. +1 +5 -2 +1 +3 -2 0 2KNO3(s) 2KNO2(s) + O2(g) N is reduced, thus N is the oxidizing agent O is oxidized, thus O is the reducing agent
  • 59. QuestionsWhat is the general rule for assigning oxidation numbers?The oxidation number is the charge a bonded atom would have if the electrons in the bond were assigned to the more electronegative elementHow is a change I oxidation number related to the process of oxidation and reduction?An increase in oxidation number indicates oxidation; a decrease in oxidation number indicates reduction.
  • 60. Sample ProblemIdentify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent. 0 0 +1 -2 2Na(s) + Cl2(g) 2NaCl(s) Cl2 is reduced, thus Cl2 is the oxidizing agent Na is oxidized, thus Na is the reducing agent
  • 61. Sample ProblemIdentify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent. 0 0 +1 -2 2HNO3(aq) + 6HI(aq) 2NO(g) + 3I2(s) + 4H2O(l) O2 is reduced, thus O2 is the oxidizing agent H2 is oxidized, thus H2 is the reducing agent
  • 62. End of section 20.2
  • 63. Identifying Redox Reactions In general, all chemical reaction can be assigned to one of two classes1. Redox reactions in which electrons are transferred from one reacting species to another. a. Many single-replacement reactions, combination reactions, decomposition reactions and combustion reactions are redox reactions.2. All other reactions in which no electron transfer occurs. a. Double-replacement reactions and acid-base reactions are not redox reactions
  • 64. What Kind of Reaction Is It? 2Mg(s) + O2(g) 2MgO(s)Combination reaction – two or more substances react to form a single new substance 2HgO (s) 2Hg (l) + O2 (g)Decomposition reaction – a single compound breaks down into two or more simpler products. Many are redox reactions
  • 65. What Kind of Reaction Is It?Zn(s) + Cu(NO3)2(aq) Cu(s) + Zn(NO3)2 (aq) Single Replacement Reaction – one element replaces a second element in a compound. 2C6H18(l) + 25O2 (g) 16CO2(g) + 18H2O(l) Combustion Reactions – an element or a compound reacts with oxygen often producing energy in the form of heat and light.. Many are redox reactions
  • 66. What Kind of Reaction Is It?Na2S(aq) + Cd(NO3)2(aq) CdS(s) + 2NaNO3(aq)Double Replacement Reaction – involving anexchange of positive ions between two compounds. 2C6H18(l) + 25O2 (g) 16CO2(g) + 18H2O(l)Combustion Reactions – an element or a compoundreacts with oxygen often producing energy in theform of heat and light.. Many are not redox reactions
  • 67. Identifying Redox ReactionsIf the oxidation number of an element in a reacting species change, then that element has undergone either oxidation or reduction. Many reactions in which color changes occur are redox reactions.
  • 68. Balancing Redox Equations Many redox reaction are too complex to be balanced by trial and error.Two systematic methods are available tobalance redox reactionsThe two methods are based on the fact that thetotal number of electrons gained in reductionmust equal the total number of electrons lost inoxidation.One method used oxidation number changes,and the other used half reactions.
  • 69. Using Oxidation Number Changes to Balance Redox Equations Oxidation Number MethodYou balance by comparing the increases and the decreases in oxidation numbers. +3 -2 +2 -2 0 +4 -2 Fe2O3(s) + CO(g) Fe(S) + CO2(g)Step 1 – assign oxidation numbers to all the atoms in the equation.
  • 70. Using Oxidation Number Changes to Balance Redox Equations Oxidation Number Method C oxidized (+2) +3 -2 +2 -2 0 +4 -2 Fe2O3(s) + CO(g) Fe(S) + CO2(g) Fe reduced (-3)Step 2 – Identify which atoms are oxidized and which are reduced
  • 71. Using Oxidation Number Changes to Balance Redox Equations Oxidation Number Method (+2) oxidized +3 -2 +2 -2 0 +4 -2 Fe2O3(s) + CO(g) Fe(S) + CO2(g) (-3) reducedStep 3 – Use one bracketing line to connect the atoms that undergo oxidation and another such line to connect those that undergo reduction.
  • 72. Using Oxidation Number Changes to Balance Redox EquationsIn a balanced redox equation, the total increase in oxidation number of the species oxidized must be balance by the total decrease in the oxidation number of the species reduced.
  • 73. Using Oxidation Number Changes to Balance Redox Equations Oxidation Number Method 3 x (+2) = +6 +3 -2 +2 -2 0 +4 -2 Fe2O3(s) + 3CO(g) 2Fe(S) + 3CO2(g) 2 x (-3) = -6Step 4 – Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients.
  • 74. Using Oxidation Number Changes to Balance Redox Equations Oxidation Number Method Fe2O3(s) + 3CO(g) 2Fe(S) + 3CO2(g)Step 5 – Make sure that the equation is balanced for both atoms and charge.
  • 75. Sample Problem Oxidation Number Method S oxidized (+4) x 3= +12+1 +6 -2 +1 -2 0 +1 -2 +1 +3 -2 +4 -22K2Cr2O7(aq) + H2O(l) + 3S(s) KOH(aq) + 2Cr2O3(s) + 3SO2 (g) Cr reduced (-3) x 4 = -12 4 Cr atoms must be reduced for each 3 S atom that are oxidized
  • 76. Sample Problem Oxidation Number Method S oxidized (+4) x 3= +12 +1 +6 -2 +1 -2 0 +1 -2 +1 +3 -2 +4 -2 2K2Cr2O7) + 2H2O + 3S 4KOH + 2Cr2O3 + 3SO2 Cr reduced (-3) x 4 = -12Check the equation and balance by inspection4 in front of KOH balances potassium2 in front of H2O balances hydrogen and oxygen.
  • 77. Sample ProblemsBalance each redox equation using the oxidationnumber change method KClO3(s) KCl (s) + O2 (g) 2KClO3(s) 2KCl (s) + 3O2 (g) HNO2 (aq) + HI (aq) NO (g) + I2 (s) + H2O (l)2HNO2 (aq) + 2HI (aq) 2NO (g) + I2 (s) + 2H2O (l)
  • 78. Sample ProblemsBalance each redox equation using the oxidationnumber change method Bi2S3 + HNO3 Bi(NO3)3 + NO + S + H2O Bi2S3 + 8 HNO3 2Bi(NO3)3 + 2NO + 3S + 4H2O SbCl5 + KI SbCl3 + KCl + I2 SbCl5 + 2KI SbCl3 + 2KCl + I2
  • 79. Using Half Reactions to Balance Redox Equations Half Reaction MethodHalf reaction is an equation showing just the oxidation or just the reduction that takes placeYou write and balance the oxidation and reduction half reaction separately before combining them into a balanced redox equationThen you balance the electrons gained in the reduction with the electrons lost in the oxidation.
  • 80. Using Half Reactions to Balance Redox Equations Half Reaction MethodS + HNO3 SO2 + NO + H2O (unbalanced) S + H+ + NO3- SO2 + NO + H2OIn this case only HNO3 is ionized. The products are covalent compoundsStep 1 – write the unbalanced equation in ionic form
  • 81. Using Half Reactions to Balance Redox Equations Half Reaction MethodOxidation Half Reaction 0 +2 S SO2Reduction Half Reaction +5 +2 NO3- NOStep 2 – Write separate half reactions for the oxidation and reduction processes.
  • 82. Using Half Reactions to Balance Redox EquationsHalf Reaction Method 0 +2 S SO2Sulfur is already balanced, but oxygen is not.The reaction takes place in acid solution, so H2O and H+ are present and can be used to balance oxygen and hydrogen as needed. 2H2O + S SO2 + 4H+Step 3 – Balance the atoms in the half reactions.
  • 83. Using Half Reactions to Balance Redox EquationsHalf Reaction Method +5 +2 NO3- NONitrogen is already balanced, but oxygen is not. 4H+ + NO3- NO + 2H2O
  • 84. Using Half Reactions to Balance Redox EquationsHalf Reaction Method 0 +4 2H2O + S SO2 + 4H+ + 4e -S is oxidized going from 0 to +4, a loss of 4 e- +5 +2 3e- + 4H+ + NO3- NO + 2H2ON is reduced going from +5 to +2, a gain of 3 e-Step 4 – Add enough electrons to one side of each half reaction to balance the charges.
  • 85. Using Half Reactions to Balance Redox EquationsHalf Reaction Method 0 +4 (x 3) 6H2O + 3S 3SO2 + 12H+ + 12e -S is oxidized going from 0 to +4, a loss of 4 e- +5 +2 (x 4) 12e- + 16H+ + 4NO3- 12 4NO + 8H2ON is reduced going from +5 to +2, a gain of 3 e-Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.
  • 86. Using Half Reactions to Balance Redox EquationsHalf Reaction Method 6H2O + 3S 3SO2 + 12H+ + 12e - 12e- + 16H+ + 4NO3- 4NO + 8H2O6H2O + 3S + 16H+ + 4NO3- + 12e - 3SO2 + 12H+ + 4NO + 8H2O + 12e – 3S + 4H+ + 4NO3- 3SO2 + 4NO + 2H2OStep 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation
  • 87. Using Half Reactions to Balance Redox EquationsHalf Reaction Method 3S + 4HNO3- 3SO2 + 4NO + 2H2OSpectator ion – are present during a reaction, but do not participate in or change during a reaction.Because none of the ions in the reactants appear in the products, there are no spectator ions in this particular example.Step 7 – Add the spectator ions and balance the equation.
  • 88. Sample Problem Half Reaction Method KMnO4 + HCl MnCl2 + Cl2 + H2O + KClK+ + MnO4- + H+ + Cl- Mn2+ + 2Cl- + Cl2 + H2O + K+ + Cl-Step 1 – write the unbalanced equation in ionic form
  • 89. Sample Problem Half Reaction MethodReduction Half Reaction +7 +2 MnO4- Mn2+Oxidation Half Reaction -1 0 2Cl- Cl2Step 2 – Write separate half reactions for the oxidation and reduction processes.
  • 90. Sample ProblemReduction Half Reaction +7 +2 8H+ + MnO4- Mn2+ + 4H2OOxidation Half Reaction -1 0 2Cl- Cl2Solution is acidic, so H2O and H+ ions are used to balance equation. (If solution is basic, H2O and OH- are used)Step 3 – Balance the atoms in the half reactions.
  • 91. Sample ProblemHalf Reaction Method +7 +2 5e- + 8H+ + MnO4- Mn2+ + 4H2OMn is reduced going from +7 to +2, a gain of 5 e- -1 0 2Cl- Cl2 + 2e-Cl is oxidized going from -1 to 0, a loss of 2 e-Step 4 – Add enough electrons to one side of each half reaction to balance the charges.
  • 92. Sample ProblemHalf Reaction Method +7 +2 (x 2) 10e- + 16H+ + 2MnO4- 2Mn2+ + 8H2OMn is reduced going from +7 to +2, a gain of 5 e- -1 0 (x5) 10Cl- 5Cl2 + 10e-Cl is oxidized going from -1 to 0, a loss of 2 e-Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.
  • 93. Sample ProblemHalf Reaction Method 10e- + 16H+ + 2MnO4- 2Mn2+ + 8H2O 10Cl- 5Cl2 + 10e-10e- + 16H+ + 2MnO4- + 10Cl- 2Mn2+ + 8H2O + 5Cl2 + 10e- 16H+ + 2MnO4- + 10Cl- 2Mn2+ + 8H2O + 5Cl2Step 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation
  • 94. Sample ProblemHalf Reaction Method 16H+ + 6Cl- + 2MnO4- + 2K+ + 10Cl- 5Cl2 + 2Mn2+ + 4Cl- + 8H2O + 2K+ + 2Cl- 16H+ = 6Cl- + 10Cl- (from the HCl in original equation) 2MnO4- = 2K+ (from the KMnO4 in original equation) 2Mn2+ = 4Cl- (from the MnCl2 in original equation) 2K+ + 2Cl- (from the KCl in original equation)Step 7 – Add the spectator ions and balance the equation.
  • 95. Sample ProblemHalf Reaction Method 16H+ + 16Cl- + 2MnO4- + 2K+ 5Cl2 + 2Mn2+ + 6Cl- + 8H2O + 2K+Adding spectator and non-spectator Cl- on each side gives 16HCl + 2KMnO4 5Cl2 + 2MnCl2 + 8H2O + 2KClStep 7 – Add the spectator ions and balance the equation.
  • 96. Sample Problem 2The following takes place in basic solution. Balance the equation using the half reaction method Zn + NO3- NH3 + Zn(OH4)2- 4Zn + NO3- + 6H2O + 7OH- 4Zn(OH4)2- + NH3
  • 97. Sample Problem 3The following takes place in basic solution. Balance the equation using the half reaction method Zn + As2O3 AsH3 + Zn2+ 6Zn + As2O3 + 9H2O 6Zn2+ + 2AsH3 + 12OH-
  • 98. Choosing a Balancing MethodIn some redox reactions, the same element is both oxidized and reduced. (called a disproportional reaction) 3Br2 + 6KOH 5KBr + KBrO3 + 3H2O Br is reduced from 0 to -1 Br is oxidized from 0 to + 5Equations like above and equation for reactions that take place in acidic or alkaline solution are best balanced by the half reaction method.
  • 99. End of Chapter 20

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