Copyright© by Houghton Mifflin Company. All rights reserved. Oxidation-Reduction Reactions (Redox)

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  • Copyright by Houghton Mifflin Company. All rights reserved.Oxidation-Reduction Reactions(Redox)

    Copyright by Houghton Mifflin Company. All rights reserved.

  • What is the difference between acid/base reactions and redox reactions?Acid/base reactions proton transfer (p+)

    Redox reactions electron transfer (e-)

  • Flow of electronsElectrons respond to differences in potential by moving from the region of high potential to the region of low potential.

    -+ High EpLow Epe-

  • Flow of electronsCl low electronegativityhigh electronegativitye-LiLithium loses the e- tug-of-war with chloride.+-

  • TerminologyCations: positively charged ions generally metalsNH4+ is the exceptionAnions: negatively charged ions non-metalscomplex ions

  • Oxidation: When a substances loses e-

    Reduction: When a substance gains e-

  • Electron Transfer and TerminologyLose electrons: OxidationGain electrons: Reduction.

  • oxidizedreduced

  • Half-reactionsCa(s) Ca2+(aq) + 2e- oxidation half reaction

    2H+(aq) + 2e- H2(g) reduction half reaction

  • Half-reactions add together Ca(s) Ca2+(aq) + 2e- 2H+(aq) + 2e- H2(g)

    Ca(s) + 2H+ + 2e- Ca2+ + 2e- + H2(g)Ca(s) + 2H+(aq) Ca2+(aq) + H2(g)

    +

  • Ca(s) + 2H+(aq) Ca2+(aq) + H2(g)

    Ca(s) has lost two e- to 2 H+(aq) to become Ca2+(aq). Ca(s) has been oxidized to Ca2+(aq)At the same time 2 electrons are gained by 2 H+(aq) to form H2(g) . We say H+(aq) is reduced to H2(g) .

  • Half-reactions add together Cu(s) Cu2+(aq) + 2e- Ag+(aq) + e- Ag(s)

    Cu(s) + 2Ag+(aq) + 2e- Cu2+(aq) + 2e- + 2Ag(s)Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

    + ( ) x 2

  • Iron comes from iron ore which is taken out of the ground by mining.

  • The pure iron is obtained by heating the ore at very high temperatures in a furnace with limestone to remove impurities.

  • This heap of iron ore pellets will be used in steel production.

  • 1. Hot blast from Cowper stoves 2. Melting zone 3. Reduction zone of ferrous oxide 4. Reduction zone of ferric oxide 5. Pre-heating zone 6. Feed of ore, limestone and coke 7. Exhaust gases 8. Column of ore, coke and limestone 9. Removal of slag 10. Tapping of molten pig iron 11. Collection of waste gases

  • Why is gaining electrons called reduction?Reduction originally meant the loss of oxygen from a compound.2 Fe2O3(s) + C(s) 4 Fe(s) + 3 CO2(g)Iron ore is reduced to metallic iron. The size of the pile gets smaller, hence the word reduction.

  • Why is losing electrons called oxidation? Oxidation originally meant the combination of an element with oxygen. 4 Fe(s) + 3 O2(g) 2 Fe2O3(g) C(s) + O2(g) CO2(g)

  • It Takes Two: Oxidation-Reduction In all reduction-oxidation (redox) reactions, one species is reduced at the same time as another is oxidized.

  • Oxidizing Agent: the species which causes oxidation is called the oxidizing agent. substances that gains electrons the oxidizing agent is always reduced

  • Reducing Agent: the species which causes reduction is called the reducing agent. the reducing agent is always oxidized. substances that give up electrons

  • Cu(s) + 2 Ag+(aq) Cu2+(aq) + Ag(s)oxidated reduced R.A. O.A.

  • A redox reaction is a chemical reaction in which electrons are transferred.Number of electrons lost by one species equals number of electrons gained by the other species.Reduction is a process in which e- are gained.Oxidation is a process in which e- are lostA reducing agent donates e- and is oxidized.A oxidizing agent gains e- and is reduced.WS 15-1

  • Electric potential (V), EoWork that must be done to move an electric charge between specified points. Electric potential differences are measured in volts . Standard conditions:

    At 25oC with all ions at 1 mol/L concentrations and all gases at 1.00 atm pressure

  • Standard Reduction PotentialsWe cannot measure the potential of an individual half-cell! We assign a particular cell as being our reference cell and then assign values to other electrodes on that basis. ( H2 half cell )

  • [H+] = 1.00 mol/L H2 (g) e-Pt gauzeThe Standard Hydrogen electrodeEo (H+(aq)/H2(g)) half-cell = 0.000 Vp{H2(g)} = 1.00 atm

  • Electric potential (V), EoIf the net potential is a positive number then the reaction is spontaneous. Products are favoured.

    If the net potential is a negative number then the reaction is non-spontaneous. Reactants are favoured.

    Half cell potentials are not doubled or tripled as per balancing. We are only comparing potentials.

  • Only one of these two reactions is possible. Which one?Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)Cu2+(aq) + 2 Ag(s) Cu(s) + 2 Ag+(aq)

    Use data table values, electrical potential, on page 7 of your data books. (2009)

  • Compare the two half reactions that make up the reaction.Cu2+(aq) + 2Ag(s) Cu(s) + 2Ag+(aqCu2+(aq) + 2e- Cu (s) Eo = +0.34 2Ag(s) 2Ag+(aq) + 2e- Eo = -0.80

    Cu2+(aq) + 2Ag(s) Cu(s) + 2Ag+(aq) Eo = -0.46Negative potential, non-spontaneous

    +

  • Compare the two half reactions that make up the reaction.Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s) Cu(s) Cu2+(aq) + 2e- Eo = -0.342Ag+(aq) + 2e- 2Ag(s) Eo = +0.80

    Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)Eo = +0.46Positive potential, spontaneous

  • ProblemWrite the oxidation/reduction half reactions and the net ionic equation when zinc is placed in Ni(NO3)2 solution. Identify the O.A. and R.A. and state if the reaction is spontaneous or non-spontaneous.

  • ProblemNi(NO3)2 Ni2+(aq) + 2NO3- (aq)

    Zn(s) + Ni2+(aq) ?

    Oxidation: Zn(s) Zn2+(aq) + 2e-+0.76Reduction: Ni2+(aq) + 2e- Ni(s) - 0.26Spectator ionAdd half reactions

  • Problem Zn(s) + Ni2+(aq) Zn2+(aq) + Ni(s) +0.50 R.A.O.A.Positive potential, spontaneous

  • SOASRAhigh attraction for electronslow attraction for electronsdecreasing strengthdecreasing strength

  • Spontaneous shortcutLocate the O.A. on the left and the R.A. on the right of the table. If the O.A. is higher up on the table than the R.A. then the reaction is spontaneous.

  • ProblemExplain what happens when nickel is placed in a zinc nitrate solution.

    Ni(s) + Zn2+(aq) R.A.O.A.

  • Zn +2 (aq) is the strongest Oxidizing agent and therefore is Reduced Zn+2 (aq) + 2e Zn (s)Ni(s) is the strongest Reducing agent andTherefore is Oxidized Ni(s) Ni +2 (aq) + 2e (must reverse)

  • NET REDOX REACTIONADD REACTIONS/REMOVE ELECTRONSZn2+ (aq) + 2e Zn(s) - 0.76 VNi (s) Ni 2+ (aq) + 2e + 0.26 V (r)

    Zn2+ (aq) Ni(s) Zn(s) + Ni2+ (aq) - 0.50VNON SPONTANEOUS REACTION

  • On the table Ni(s) Zn2+(aq) R.A. is above the O.A.Non-spontaneousWS 15-23

  • DISPROPORTIONATIONredox reactions in which the oxidizing agent and the reducing agent are the same species2 H2O2 (l) ------> 2 H2O (l) + 2 O2 (g) O ( -1) O ( -2) O ( 0)Oxygen -------- > reduced / OAOxygen ---------------------> oxidized / RA

  • Predicting redox reactionsList all species present.Choose the strongest oxidizing and reducing agent. Watch for acids ( H +) Also water H2OWrite the reduction half reaction, as written in the data book. Write the oxidation half reaction, reverse the equation in the data book.Balance number of electrons.Add the two half reactions together to form the net ionic equation.Predict if reaction is spontaneous or not.

  • ProblemsA mixture of bromine gas and chlorine gas is added to a solution of copper (II) sulphate and a copper strip.Br2(g)Cl2(g)H20(l)Cu2+(aq)Cu(s)SO42-(aq)RAOACl2(g) + 2e- 2 Cl-(aq) Cu(s) Cu2+(aq) + 2e- Cl2(g) + Cu(s) 2 Cl-(aq) + Cu2+(aq)

  • ProblemsLead is placed in a zinc nitrate solution.NO3-(aq)H20(l)Zn2+(aq)Pb(s)RAOANon-spontaneous OA is below RA

  • ProblemsA few drops of Hg(l) are dropped into a solution which is 1.0 M in both sulphuric acid and potasium permanganate.MnO4-(aq)SO42-(aq)H20(l)K+(aq)Hg(l)H+(aq)RAOA

  • ProblemsA few drops of Hg(l) are droped into a solution which is 1.0 M in both sulphuric acid and potasium permanganate. MnO4-(aq) + 8 H+(aq) + 5e- Mn2+(aq) + 4 H2O(l) Hg(l) Hg2+(aq) + 2e- 2MnO4-(aq) + 16H+(aq) + 5Hg(l) 2Mn2+(aq) + 8H2O(l) + 5Hg2+(aq)( ) x2 ( ) x5

  • General RulesMetal (+) ions are oxidizing agents.Nonmetal (-) ions are reducing agents.Metal elements are reducing agents.Nonmetal elements are oxidizing agents.

  • Building a redox table (method one)One can use experimental evidence to determine the relative strengths of oxidizing and reducing agents.The greater the number of spontaneous reactions, the stronger the oxidizing agent.This means we can rank oxidizing agents according to the number of spontaneous reactions.By convention the strongest oxidizing agent is at the top left in a redox table and the strongest reducing agent is at the bottom right of the table.

    SOASRAReduction Table

  • Problem: Make a redox table Cu(s) Mg(s) Ag(s) Zn(s)

    Cu2+(aq)________________Mg2+(aq)________________Ag+(aq)________________Zn2+(aq)________________Virtual Lab

  • REDOX TABLEMg(s) is most reactive - 3 timesZn(s) is second - 2 timesCu(s) is third - 1 timesAg(s) is most unreactive 0 timesTHEREFORE Mg oxidizes easiest and is the strongest reducing agentMg is placed in the lower right hand side then Zn(s) Cu(s) and Ag(s) is last.Now you fill in the reduction reactions

  • Activity Series - Redox TableAg+(aq) + 1e- Ag(s) Cu2+(aq) + 2e- Cu(s)Zn2+(aq) + 2e- Zn(s)Mg2+(aq) + 2e- Mg(s)Strongest reducing agentsWeakerStrongest oxidizing agents WeakerMg is the strongest reducing agent as it oxidizes the most and is on the lower right side of the table.

  • Redox Table Building (method two)

    The spontaneity rule is used to order the oxidizing agents to produce a redox table.Consider the following redox equations which represent spontaneous reactions from an experiment. From this evidence construct a redox table.

  • Redox Table Building3 Equations givenall are spontaneous reactions Co(s) + Pd2+(aq) Co2+(aq) + Pd(s)Pd(s) + Pt2+(aq) Pd2+(aq) + Pt(s)Mg(s) + Co2+(aq) Mg2+(aq) + Co(s)

    Work with one equation at a time.

  • Redox Table Building Co(s) + Pd2+(aq) Co2+(aq) + Pd(s)Pd2+(aq) + 2 e- Pd(s) reduced/stays as isCo2+(aq) + 2 e- Co(s) oxidized/reverseOA is above RAReverse the oxidation reaction tocompare Pd(s) with Co(s)

    RAOA

  • NOTE POSITION OF THE REDUCING AGENTSPt(s) is above Pd(s) in position

    Pt(s)

    Pd(s) stronger reducing agent

  • Redox Table Building Pd(s) + Pt2+(aq) Pd2+(aq) + Pt(s)Pt2+(aq) + 2 e- Pt(s) Pd2+(aq) + 2 e- Pd(s)OA is above RAspontaneous reaction

    OARA

  • NOTE POSITION OF THE REDUCING AGENTSPt(s) is above Pd(s) in position

    Pt (s) Pd (s) Stronger reducing agent

  • Redox Table Building Mg(s) + Co2+(aq) Mg2+(aq) + Co(s)Co2+(aq) + 2 e- Co(s) ( stays as is)Mg2+(aq) + 2 e- Mg(s) (reverse)OA is above RAspontaneous reactionOARA

  • MAKING THE TABLEUSING THE 4 REDUCING REAGENTS Mg(s) Pt(s) Co(s) Pd(s) Place into the correct order using the previous information you collecteNote the position with respect to each other.Make the complete reactions MAKE THE REDOX TABLE

  • Redox Table Building Pt2+(aq) + 2 e- Pt(s) Pd2+(aq) + 2 e- Pd(s) Co2+(aq) + 2 e- Co(s) Mg2+(aq) + 2 e- Mg(s)

  • Oxidation StatesSome reactions are not adequately explained with redox theories.Chemists have developed a method of electron bookkeeping to describe the redox of molecules and complex ions.

  • Oxidation StatesOxidation state:apparent net charge that an atom would have if electron pairs belonged entirely to the more electronegative atomOxidation number:a positive or negative number assigned to a combined atom according to a set of arbitrary numbers.

  • Assigning Oxidation Numbers1) Oxidation numbers for all uncombined elements (elemental/standard) = 0K(s) = 0 N2(g) = 0S8(s) = 02) Oxidation number for all simple ions is equal to the charge of the ion.Br1-(aq) = -1 Fe3+(aq) = +33) Oxidation for oxygen in a compound = -2 (except for peroxides = -1)H2O(l)H2O2(l)

    -2-1

  • Assigning Oxidation Numbers4) Hydrogen in compounds = +1 H2O(l) (except hydrides = -1) NaH(s)

    5) Sum of oxidation numbers in a compound is = 0H2O(l) (2 x +1) + (1 x -2) = 06) Sum of oxidation numbers in a complex ion = charge of ion.NH4+(aq) (4 x +1) + (1 x -3) = +1

  • ExampleWhat is the oxidation number for carbon in CO32-(aq) ?CO# + 3 OO# = -2 ? + 3 (-2) = -2 ? + -6 = -2 ? = +4

  • ExampleWhat is the oxidation number for carbon in C6H12O6 ?6 CO# + 12 HO# + 6 OO# = 0 6 (?) + 12 (+1) + 6 (-2) = 0 6 (?) + 12 + -12 = 0 ? = 0

  • If you have 2 unknowns?First ionize the substance in water. Then work out the two resulting ions separately.Example: CuSO4(aq)CuSO4(s) Cu2+(aq) + SO42-(aq)Simple ion+2Solve as a complex ion4(-2) + 1(x) = -2 x = +6+6 -2

  • Assign oxidation numbers to chlorine in each of the following chemicals. HCl(aq)Cl2(g)NaClO (s)Cl-(aq)HClO3(aq)ClO3(aq)

    -10+1-1+5+6

  • Who cares about oxidation numbers?Determining oxidation numbers allows us to predict electron transfer.If there is an increase in oxidation number then oxidation occurs.If there is a decrease in oxidation number then reduction occurs.

  • ProblemDetermine the oxidation numbers for all atoms and ions in the following redox equation and indicate which substance is undergoing oxidation and reduction.

  • ProblemCH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)+1-40-2+4-2+1C is oxidizedO is reduced

  • Identifying Redox Reactions.Which of these are Redox reactions?1) N2O4(g) 2NO2(g)

    2) Cl2(g) + 2NaBr(aq) 2 NaCl(aq) + Br2(l)

    3) PbCl2(aq)+ K2SO4(aq) 2KCl(aq) + PbSO4(aq)

    4) 2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l)

    5) 2K(s) + 2H2O(l) 2KOH(aq) + H2(g)+4 -2 +4 -20 +1 -1 +1 -1 0NOYES+2 -1 +1 +6 -2 +1 -1 +2 +6 -2NO+1 -2 +1 +1 +6 -2 +1 +6 -2 +1 -2NO0 +1 -2 +1 -2+1 0YES

  • Ion electron methodUnder Acidic conditions1. Identify oxidized and reduced species Write the half reaction for each.2. Balance the half rxn separately except H & Os.Balance: Oxygen by adding H2OBalance: Hydrogen by adding H+Balance: Charge by adding e -3. Multiply each half reaction by a coefficient. Must have the same # of e- in both half-rxn.4. Add the half-rxn together, the e - should cancel.

  • Balancing Half ReactionsMnO4 Mn2+ MnO4 Mn2+ + 4 H2O 8 H+ + MnO4 Mn2+ + 4 H2O 5 e + 8 H+ + MnO4 Mn2+ + 4 H2O Note All elements balance Note The charge is balanced / LHS 2 + = R...

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