Oxidation/Reduction

1
Oxidation/Reduction

• Chapter 20

2
Two Types of Chemical Rxns

• Exchange of Ions no change in charge/oxidation
  numbers
• Acid/Base Rxns
• NaOH HCl

3
Two Types of Chemical Rxns

• Precipitation Rxns
• Pb(NO3)2(aq) KI(aq)
• Dissolving Rxns
• CaCl2(s) ?

4
Two Types of Chemical Rxns

• Exchange of Electrons changes in oxidation
  numbers/charges
• Fe(s) CuSO4(aq) ? FeSO4(aq) Cu(s)
• Remove spectator ions
• Fe Cu2 ? Fe2 Cu
• Protons
• Electrons

5
Review of Oxidation Numbers

• Oxidation numbers the charge on an ion or an
  assigned charge on an atom.
• Al Cl2 P4
• Mg2 Cl-

6
Review of Oxidation Numbers

• Calculate the oxidation numbers for
• HClO Cr3
• S8 Fe2(SO4)3
• Mn2O3 SO32-
• KMnO4 NO3-
• HSO4-

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Oxidation

• Classical Definition addition of oxygen
• Modern Definition an increase in oxidation
  number
• Fe O2 ? Fe2O3
• CO O2 ? CO2
• CH3CH2OH ? CH3CHO ? CH3COOH

8
Oxidation

• Fe O2 ? (limited oxygen)
• Fe O2 ? (excess oxygen)

9
Oxidation

• C O2 ? (limited oxygen)
• C O2 ? (excess oxygen)

10
Reduction

• Classical addition of hydrogen
• Modern decrease (reduction) in oxidation number
• N2 3H2 ? 2NH3 (Haber process)
• R-CC-R H2 ?
• H H
• (unsaturated fat) (saturated fat)

11
Oxidizing/Reducing Agents

• Oxidation and Reduction always occur together.
• Oxidizing Agents
• Get reduced
• Gain electrons
• Reducing Agents
• Get oxidized
• Lose electrons

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Oxidizing/Reducing Agents

• 0 2
• Cu O2 ? CuO
• 0 -2

Got oxidized, reducing agent
Got reduced, oxidizing agent
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• Identify the Oxidizing/Reducing Agents in the
  following (Calculate the ox. numbers also).
• Cu S8 ? Cu2S
• H2 O2 ? H2O
• Cu AgNO3 ? Cu(NO3)2 Ag
• H2O Al MnO4- ? Al(OH)4- MnO2

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• Identify the Oxidizing/Reducing Agents in the
  following (Calculate the ox. numbers also).
• Al O2 ? Al2O3
• Li N2 ? Li3N
• Cu(NO3)2 Fe ? Fe(NO3)3 Cu
• KMnO4 FeSO4? Fe2(SO4)3 Mn K

16
Balancing Redox Reactions

• Half-Reaction Method
• Break eqn into oxidation half and reduction half
• Easy Examples
• Al Fe2 ? Fe Al3
• Cu Zn2 ? Cu2 Zn
• Mg Na ? Mg2 Na

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Balancing Redox Reactions

• Whats really happening
• Cu2 Zn ? Cu Zn2

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Balancing Redox Reactions

• Steps for more complicated examples
• Balance all atoms except H and O
• Balance charge with electrons
• Balance O with water
• Balance H with H
• 5. (Add OH- to make water in basic solutions)

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Balancing Redox Reactions

• Example 1
• MnO4- C2O42- ? Mn2 CO2
• 1. Separate into half reactions
• MnO4- ? Mn2
• C2O42- ? CO2

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Balancing Redox Reactions

• MnO4- ? Mn2
• 7
  2
• 5e- MnO4- ? Mn2
• 5e- MnO4- ? Mn2 4H2O
• 8H 5e- MnO4- ? Mn2 4H2O

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Balancing Redox Reactions

• C2O42- ? CO2
• C2O42- ? 2CO2
• 3
  4 (1 e- per carbon)
• C2O42- ? 2CO2 2e-

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Balancing Redox Reactions

• C2O42- ? 2CO2 2e- (X
  5)
• 8H 5e- MnO4- ? Mn2 4H2O (X 2)
• 5C2O42- ? 10CO2 10e-
• 16H 10e- 2MnO4- ? 2Mn2 8H2O
• 16H5C2O42-2MnO4- ? 2Mn2 10CO2 8H2O

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Balancing Redox Reactions (Acidic Solutions)

• Cr2O72- Cl- ? Cr3 Cl2
• 14 H Cr2O72- 6Cl- ? 2Cr3 3Cl2 7H2O
• Cu NO3- ? Cu2 NO2
• Cu 4H 2NO3- ? Cu2 2NO2 2H2O
• Mn2 BiO3- ? Bi3 MnO4-
• 14H 2Mn2 5BiO3- ? 5Bi3 2MnO4- 7H2O

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Balancing Redox Reactions (Basic Solutions)

• Add OH- AT THE VERY END ONLY!!!!!
• NO2- Al ? NH3 Al(OH)4-
• 5H2O OH- NO2- 2Al ? NH3 2Al(OH)4-
• Cr(OH)3 ClO- ? CrO42- Cl2
• 2Cr(OH)3 6ClO- ? 2CrO42- 3Cl2 2OH- 2H2O

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• Balance in Both Acidic and Basic Solutions
• F- MnO4- ? MnO2 F2
• HNO2 H2O2 ? O2 NO
• H is 1

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• F- MnO4- ? MnO2 F2
• 8H 6F- 2MnO4- ? 2MnO2 3F2
  4H2O
• 4H2O 6F- 2MnO4- ? 2MnO2 3F2
  8OH-
• HNO2 H2O2 ? O2 NO
• 2HNO2 H2O2 ? O2 2NO 2H2O

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Voltaic (Galvanic) Cells

• Voltaic(Galvanic) Cells redox reactions that
  produce a voltage
• Spontaneous reactions (DGlt0)
• Voltage of the cell (Eocell) is positive
• Batteries
• Electrolytic cells redox reactions that must
  have a current run through them.
• DGgt0 and Eocell is negative.
• Often used to plate metals

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Voltaic (Galvanic) Cells

• History
• Galvani (died 1798) uses static electricity to
  move the muscles of dead frogs
• Volta (1800) Created the first battery

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Voltaic (Galvanic) Cells

• Voltaic cell
• Anode Oxidation site
• Cathode Reduction site (RC cola)
• Salt bridge completes the circuit

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Voltaic (Galvanic) Cells

• Cell Notation
• Zn Zn2(aq) Cu2(aq) Cu
• Anode Zn ? Zn2 2e-
• Cathode Cu2 2e- ? Cu
• Cell Cu2 Zn ? Cu Zn2

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Voltaic (Galvanic) Cells
34
Hydrogen Electrode

• Standard Electrode
• Voltage(potential) 0 Volts
• 2H(aq) 2e- ? H2(g) 0 volts
• H2(g) ? 2H(aq) 2e- 0 volts
• 3. Often used in electrodes (like pH)

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Standard Reduction Potentials

• Rules
• Flipping an equation changes the sign of E
• Multiplying an equation does not change the
  magnitude of E

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Calculating Cell Potential

• A cell is composed of copper metal and Cu2(aq)
  on one side, and zinc metal and Zn2(aq) on the
  other. Calculate the cell potential.
• Zn2 2e- ? Zn -0.76 V
• Cu2 2e- ? Cu 0.34 V
• flip the zinc equation
• Zn ? Zn2 2e- 0.76 V
• Cu2 2e- ? Cu 0.34 V
• Zn Cu2 ? Zn2 Cu 1.10 V

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• What is the cell emf of a cell made using Cu and
  Cu2 in one side and Al and Al3 in the other?
  Write the complete cell reaction.
• ANS 2.00 V

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• Calculate the standard emf for the following
  reaction. Hint break into half-reactions.
• 2Al(s) 3I2(s) ? 2Al3(aq) 6I-(aq)

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• A voltaic cell is based on the following half
  reactions.
• In(aq) ? In3(aq) 2e-
• Br2(l) 2e- ? 2Br-(aq) 1.06 V
• If the overall cell voltage is 1.46 V, what is
  the reduction potential for In3?

41

• Calculate the standard emf for the following
  reaction.
• Cr2O72- 14H 6I- ? 2Cr3 3I2 7H2O

42

• Two half reactions in a voltaic cell are
• Zn2(aq) 2e- ? Zn(s)
• Li(aq) e- ? Li(s)
• Calculate the cell emf.
• Which is the anode? Which is the cathode?
• Which electrode is consumed?
• Which electrode is positive?
• Sketch the cell, indicating electron flow.

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• Given the following half-reactions
• Pb2 2e-? Pb
• Ni2 2e- ? Ni
• Calculate the cell potential (Eo).
• Label the cathode and anode.
• Identify the oxidizing and reducing agents.
• Which electrode is consumed?
• Which electrode is plated?
• Sketch the cell, indicating the direction of
  electron flow.

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Strengths of Oxidizing and Reducing Agents

• larger the reduction potential, stronger the
  oxidizing agent
• Wants to be reduced, can oxidize something else.
• lower the reduction potential, stronger the
  reducing agent
• Would rather be oxidized

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• F2 2e- ? 2F- 2.87 V
• Stronger Cl2 2e- ? 2Cl- 1.36 V
• Oxidizing .
• Agents .
• .
• .
• .
• Li e- ? Li -3.05V

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Example 1

• Which of the following is the strongest oxidzing
  agent? Which is the strongest reducing agent?
• NO3- Cr2O72- Ag

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• Which of the following is the strongest reducing
  agent? Which is the strongest oxidizing agent?
• I2(s) Fe(s) Mn(s)

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• Can copper metal (Cu(s)) act as an oxidizing
  agent?

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Spontaneity
Voltaic Cells Electrolytic Cells
Positive emf Spontaneous Can produce electric current Batteries Negative emf Not spontaneous Must pump electricity in Electrolysis
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Example 1

• Are the following cells spontaneous as written?
• Cu 2H ? Cu2 H2
• Cl2 2I- ? 2Cl- I2
• I2 5Cu2 6H2O ? 2IO3- 5Cu 12H
• Hg2 2I- ? Hg I2

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EMF and DGo

• DG -nFE
• n number of electrons transferred
• E Cell emf
• F 96,500 J/V-mol (Faradays Constant)
• Positive Voltage gives a negative DG (spont)

52

• Calculate the cell potential and free energy
  change for the following reaction
• 4Ag O2 4H ? 4Ag 2H2O
• ANS 0.43 V, -170 kJ/mol

53

• Calculate DG and the EMF for the following
  reaction. Also, calculate the K.
• 3Ni2 2Cr(OH)3 10OH- ? 3Ni 2CrO42- 8H2O
• ANS 87 kJ/mol, -0.15 V, 6 X 10-16

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EMF and K

• DGo -RTlnK (DGo -nFEo)
• -nFEo -RTlnK
• lnK nFEo (assume 298 K)
• RT
• log K nEo
• 0.0592

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Example 1

• Calculate DG, cell voltage and the equilibrium
  constant for the following cell
• O2 4H 4Fe2 ? 4Fe3 2H2O
• ANS -177 kJ/mol, 0.459 V, 1 X 1031

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Example 2

• If the equilibrium constant for a particular
  reaction is 1.2 X 10-10, calculate the cell
  potential. Assume n 2.

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Concentration Cells Nernst Equation

• DG DGo RT lnQ
• -nFE -nFEo RT lnQ
• E Eo - RT lnQ (assume 298 K)
• nF
• E Eo - 0.0592 log Q
• n
• Can adjust the voltage of any cell by changing
  concentrations

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Using the Nernst Eqn

• Suppose in the following cell, the concentration
  of Cu2 is 5.0 M and the concentration of Zn2 is
  0.050 M. Calculate the cell voltage.
• Zn(s) Cu2(aq)?Zn2(aq) Cu(s) 1.10 V

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• E Eo - 0.0592 V log Q
• n
• E 1.10V - 0.0592 V log CuZn2
• 2 Cu2Zn
• E 1.10V - 0.0592 V log Zn2
• 2 Cu2
• E 1.10V - 0.0592 V log 0.050
• 2 5.0
• E 1.16 V

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Example 1

• Calculate the emf at 298 K generated by the
  following cell (Eo 0.79 V) where Cr2O72- 2.0
  M, H 1.0 M, I-1.0 M and Cr3 1.0 X
  10-5M.
• Cr2O72- 14H 6I- ? 2Cr3 3I2 7H2O
• ANS 0.89 V

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Example 2

• Calculate the emf at 298 K generated by the
  following cell (Eo 2.20 V) where Al3 0.004
  M and I- 0.010 M.
• 2Al(s) 3I2(s) ? 2Al3(aq) 6I-(aq)
• ANS 2.36 V

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Example 3

• If the voltage of a Zn-H cell is 0.45 V at 298 K
  when Zn21.0 M and PH21.0 atm, what is the
  concentration of H? Note that atm can be used
  just like molarity.
• Zn(s) 2H(aq) ? Zn2(aq) H2(g)
• ANS 5.2 X 10-6M

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Example 4

• What pH is required if we want a voltage of 0.542
  V and Zn20.10 M and PH21.0 atm?
• Zn(s) 2H(aq) ? Zn2(aq) H2(g)
• ANS 5.84 X 10-5M, pH 4.22

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Batteries

• Lead Acid Battery
• 12 Volt DC
• Discharges when starting the car, recharges as
  you drive (generator). Running reaction
  backward.

PbO2(s) Pb(s) 2HSO4-(aq) 2H(aq) ?2PbSO4(s)
2H2O(l)
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Alkaline Batteries

• Basic
• Zinc can acts as the anode
• 2MnO2(s)2H2O(l)2e-?2MnO(OH)(s) 2OH-(aq)
• Zn(s) 2OH-(aq)? Zn(OH)2(s) 2e-
• Rechargable uses Ni-Cd

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Corrosion

• Iron rusts in acidic solns (not above pH9)
• Water needs to be present
• Salts accelerate the process
• O2 4H 4e- ? 2H2O
• Fe ? Fe2 2e-
• (The Fe2 eventually goes to Fe3, Fe2O3)

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Preventing Corrosion

• Paint
• Sometimes oxide layer(Al2O3)
• Galvanizing (coating Fe with Zn)
• Fe2 2e- ? Fe E -0.44 V
• Zn2 2e- ? Zn E -0.76 V
• Zinc is more easily oxidized
• (Zn ? Zn2 2e- E 0.76 V)

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• Cathodic protection (sacrificial anode)
• Magnesium used in water pipes
• Magnesium rods used in hot water heaters

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• An iron gutter is nailed using aluminum nails.
  Will the nail or the iron gutter corrode first?
• Fe2 2e- ? Fe E -0.44 V
• Al3 3e- ? Al E -1.66 V
• Al will corrode first (Al ? Al3 3e- ,E
  1.66 V)

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• Which of the following metals could provide
  cathodic protection to iron
• Al Cu Ni Zn

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Electrolysis and Electroplating

• Plating of silver on silverware

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Electrolytic cells

• Must run electricity through them
• Running a voltaic cell backwards
• Used to produce sodium metal
• Na(aq) e- ? Na (s) -2.71 V
• Cl2(g) 2e- ? 2Cl-(aq) 1.36 V

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• As a voltaic cell
• 2Na(s) ? 2Na(aq) 2e- 2.71 V
• Cl2(g) 2e- ? 2Cl-(aq) 1.36 V
• 2Na(s) Cl2 (s)?2NaCl(aq) 4.07 V
• As an electrolytic cell
• 2Na(aq) 2e- ? 2Na (s) -2.71 V
• 2Cl-(aq) ? Cl2(g) 2e- -1.36 V
• 2NaCl(aq) ? 2Na(s) Cl2 (s) -4.07 V

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Quantitative Electrolysis

• Electric current Amperes
• 1 ampere 1Coloumb I Q
• 1 second t
• 1 F 96,500 C/mol
• One mole of electrons has a charge of 96,500 C
• One electron has a charge of 1.602 X 10-19 C

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• What mass of aluminum can be produced in 1.00
  hour by a current of 10.0 A?
• Al3 3e- ? Al

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• Moles of e- (36,000C)(1 mol e-) 0.373 mol e-
• (96,500 C)
• Al3 3e- ? Al
• 0.373 mol
• Al3 3e- ? Al
• 0.373 mol 0.124 mol Al ? 3.36 g Al

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Example 2

• What mass of magnesium can be produced in 4000 s
  by a current of 60.0 A?
• Mg2 2e- ? Mg
• ANS 30.2 g Mg

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• What current is required to plate 6.10 grams of
  gold in 30.0 min?
• Au3 3e- ? Au

81

• How long would it take to plate 50.0 g of
  magnesium from magnesium chloride if the current
  is 100.0 A?

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• Given the following
• Ag(aq) e- ? Ag(s) 0.799V
• Fe3(aq) e- ? Fe2(aq) 0.771 V
• Write the reaction that occurs.
• Calculate the standard cell potential.
• Calculate DGrxn for the reaction from the cell
  potential.
• Calculate K for the reaction.
• Predict the sign of DSrxn.
• Sketch the cell, labeling anode, cathode, and the
  direction of electron flow.

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• Do SO3 and SO32- have the same molecular shape?
  How about SO2?

84

• 16. a) Not redox
• b) I oxidized (-1 to 5) , Cl reduced (1 to
  -1)
• c) S oxidized (4 to 6), N reduced (5 to 2)
• d) Br oxidized (-1 to 0), S reduced (6 to 4)
• 20 a. Mo3 3e- ? Mo
• b. H2O H2SO3 ? SO42- 2e- 4H
• c. 4H 3e- NO3- ? NO 2H2O
• d. 4H 4e- O2 ? 2H2O
• e. 4OH- Mn2 ? MnO2 2e- 2H2O
• f. 5OH- Cr(OH)3 ? CrO42- 3e- 4H2O
• g. 2H2O 4e- O2 ? 4OH-

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• 22. a. 3NO2- Cr2O72- 8H ? 3NO3- 2Cr3
  4H2O
• 2HNO3 2S H2O ? 2H2SO3 N2O
• 2Cr2O72- 3CH3OH 16H ? 4Cr3 3HCO2H 11H2O
• 2MnO4- 10Cl- 16H ? 2Mn2 5Cl2 8H2O
• NO2- 2Al 2H2O ? NH4 2AlO2-
• H2O2 2ClO2 2OH- ? O2 2ClO2- 2H2O

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• 26. a) Al oxidzed, Ni2 reduced
• b) Al ? Al3 3e- Ni2 2e- ? Ni
• c) Al anode, Ni cathode
• d) Al negative, Ni positive
• e) Electrons flow towards the Ni electrode
• f) Cations migrate towards Ni electrode

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• a) Cd is anode, Pd is cathod
• b) Ered 0.63 V
• 36 a) 2.87 V b) 3.21 V c) -1.211 V d) 0.636V
• 38 a) 1.35 V b) 0.29 V

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• 41 a) Mg b) Ca c) H2 d) H2C2O4
• 42 a) Cl2 b) Cd2 c) BrO3- d) O3
• 44. a) Ce3 (weak reductant)
• b) Ca (strong reductant)
• c) ClO3- (strong oxidant)
• d) N2O5 ( oxidant)
• a) H2O2 strongest oxidizing agent
• b) Zn strongest reducing agent
• 50.a) 3.6 X 108 b) 1041 c) 10103
• 52. 0.292 V
• 54 a) 4 X 1015 b) 2 X1065 c) 7.3 X1049

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• 62a) 2.35 V b) 2.48 V c) 2.27 V
• 64. a) 0.771 V b) 1.266 V
• 88. a) 173 g b) 378 min
• E 1.10 V Wmax -212 kJ/mol Cu
• W -1.67 X 105 J

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• 1a) 14H Cr2O72- 3Fe ? 2Cr3 3Fe2
  7H2O
• 2Br- F2 ? 2F- Br2
• 4OH- 2Cr(OH)3 ClO3- ? 2CrO42- Cl-
  5H2O
• 2b) 0.463 V c) -89.4 kJ/mol d) 4.4 X 1015
• e) 0.442 V
• F2 is str. oxidizing agent, Li, str. reducing
  agent
• b) 78 minutes c) 1.19 g d) 0.695 g

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• In a measuring cup
• 5 mL of oil
• 5 mL of ethanol
• 5 mL of 50 NaOH solution (approximately 30
  drops).
• Place in beaker
•  Heat the mixture, stirring with popsicle stick.
• Remove from heat. After 5 minutes, add 10 mL of
  saturated salt solution.
• Collect some of the solid and test the pH of your
  soap. Compare the pH to that of commercial bar
  soap and liquid detergent solution. See if it
  lathers.