IB topic 9 Oxidationreduction

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IB topic 9 Oxidation-reduction

• Define oxidation and reduction in terms of
  electron loss and gain.

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• 2Mg O2 ? 2MgO
• Reduction-charge goes down
• OILRIG
• Redox always occurs together
• Question 1

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9.2Redox equations

• Deduce simple oxidation and reduction
  half-equations given the species involved in a
  redox reaction.
• 2Fe 3Cl2? 2FeCl3

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Deduce the oxidation number of an element in a
compound.

• means loss, gain of e-
• Rules
• Elements Na, O2, S8 0
• Group 1 1 H1
• O -2 halogens -1
• Many exceptions
• Ox. add up to the charge on the species
• In covalent compounds more electronegative is
  ie NH3, CCl4

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• Give O numbers to each element
• H2SO4, SO32-
• NH4, Fe2O3, K2Cr2O7, CuCl2,
• Question 2
• Question 3

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State the names of compounds using oxidation
numbers.

• MnO2, FeO, CuCl, Na2O
• Manganese (IV) oxide, iron (II) oxide, Copper (I)
  chloride sodium oxide
• Cu(H2O)62 CuCl42-
• Hexaaquacopper(II) ion
• Tetrachlorocopper (II) ion

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Deduce whether an element undergoes oxidation or
reduction in reactions using oxidation numbers.

• Ca Sn2 ? Ca2 Sn
• 4NH3 5O2? 4NO 6H2O
• Disproportionation
  Cl2 H2O?HCl HClO
• Question 4

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Deduce redox equations using half-equations.

• Steps
• Assign O numbers and write half reactions
• Balance atoms other than H and O
• Balance O by adding H2O as needed
• Balance H by adding H as needed
• Balance Charges by adding e- to the side
• Equalize the e- by multiplying
• Add the half reactions together

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• Try NO3- Cu ? NO Cu2
• 5, -2, 0 on left 2,-2,2 on right
• Cu ? Cu2 ox NO3-? NO red
• 4H NO3- ? NO 2 H2O
• Cu ? Cu2 2e-
• 4H NO3- 3e-? NO 2 H2O
• 8H 2NO3- 6e- 3Cu ? 2NO 4H2O
  3Cu2 6e-

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Fe3 MnO4- ? Fe3 Mn2
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SO32- Cr2O72-? SO42- Cr3
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• Internet example
• Question 5

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Define the terms oxidizing agent and reducing
agent.

• A substance that gets reduced causes oxidation so
  it is an oxidizing agent

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• Identify the oxidizing and reducing agents in
  redox equations.
• Fe2O3 3C ? 2Fe 3CO2
• Fe oxidizing C reducing
• IO3- 5I- 6H ? 3I2 3H2O
• IO3- oxidizing I- reducing
• questions 6

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9.3Reactivity

• Deduce a reactivity series based on the chemical
  behavior of a group of oxidizing and reducing
  agents.
• More reactive metals lose their e- more readily
  becoming a strong reducing agent
• Zn CuSO4
• Stronger Mg, AL, Zn, Fe, Pb, Cu, Ag
• simulation

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• Non metals F2 strongest oxidizing agent, most
  readily becomes reduced Cl2, Br2, I2
• Question 7

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Deduce the feasibility of a redox reaction from a
given reactivity series.

• Yes or no
• ZnCl2 Ag
• 2FeCl3 3 Mg
• Cl2 2KI
• question 8

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94 Voltaic Cells (battery)

• Explain how a redox reaction is used to produce
  electricity in a voltaic cell
• Zn(s) ? Zn2(aq) 2e- red. Agent
• Other half cell
• Cu2 2e- ? Cu(s) reduced
• This combination is a voltaic cell

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State that oxidation occurs at the negative
electrode (anode) and reduction occurs at the
positive electrode (cathode)

• Which is the anode (where e- leave) /cathode?
• Draw this setup.
• Animation
• Draw Zn/Zn2 and Ag/Ag and give the potential,
  show the flow of e-
• Where is oxidation and reduction

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• Connect these half cells with a salt bridge
• this is a spontaneous reaction
• Animation
• Question 9

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95 Electrolytic cells

• Describe, using a diagram, the essential
  components of an electrolytic cell.
• Opposite of a voltaic cell
• Requires electrical energy
• ? means then in diagrams
• Animation

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State that oxidation occurs at the negative
electrode and reduction occurs at the positive
electrode

• The power source pushes e- to the electrode
• -electrode attracts ions
• -electrode is the cathode cations gain e- so are
  reduced
• Show the electrolysis of MgF2

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Describe how current is conducted in an
electrolytic cell

• Do questions 10

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Deduce the products of the electrolysis of a
molten salt

• Diagram the electrolysis of molten(melted) NaCl
• Tell where oxidation and reduction occurs
• Do question 11

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19.1 Standard electrode potentials

• Describe the standard hydrogen electrode.

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Standard H cell

• Conditions - Pt electrode
• H2 gas at 1 atm pressure
• 1 mol dm-3 H
• 298 K or 25oC
• 0.00 V
• Attach a half cell if e- flows to H2 it is
• Like Zn which is -0.76 V

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conventions

• Zn(s)/Zn2H(aq)/1/2 H2(g) (Pt)
• Oxidation on left side
• More value of electrode potentials give off e-

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Define the term standard electrodepotential (E Ö
) .

• relative electrode potential compared under
  standard conditions with the standard hydrogen
  electrode
• Look at your data booklet

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19.1.3 Calculate cell potentials usingstandard
electrode potentials.

• Try Cr2O72- and Br2
• Answer 0.26 V

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Predict whether a reaction willbe spontaneous
using standardelectrode potential values.

• Can a solution of tin II ions reduce a solution
  of iron III ions?
• Sn4 2 e- ? Sn2 Eo 1.33
• Yes 0.62V

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• Can a solution of Sn4 ions reduce Fe3 to Fe
• No what does work
• 
  
• Fe2 and 0.59 V
• Do question 12

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19.2 Electrolysis

• Predict and explain the products of electrolysis
  of aqueous solutions.
• For water need DC in a dilute solution of H2SO4
• H to H2 given off at the electrode
• OH- to O2 at the electrode
• 2H2O ? 4H O2 4 e-

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• Electrolysis of NaCl(aq)
• - electrode H2
• electrode dilute OH- to O2 conc Cl- to Cl2
• Write half reactions
  
  
• Do question 13

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Determine the relative amounts of theproducts
formed during electrolysis.

• Position in the electrochemical series
• ions lower in the series will gain e- at the
  electrode (cathode) in perference to those higher
• Hydroxide ions release e- to form oxygen and H2O
  in preference to other anions at the positive
  electrode

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• In some cases concentrations ( more concentrated
  may be discharged)
• Nature of the electrode C and Pt are inert
• List all the cations and anions
• Cations lower in the series gain e- more readily

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Describe the use of electrolysis
inelectroplating.

• CuSO4(aq) with copper electrodes
• Cu2 goes to electrode and plates Cu
• electrode Cu goes to Cu2(use impure ore)

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Electroplating

• Object to be electroplated is put at the negative
  electrode and is placed in a solution of ions of
  the metal used to plate it.

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Factors affecting relative amounts

• Charge on the ions
• Na Cu2 Al3 Al takes more energy to make
• Quantity of e- (amperage and time) charge
  current x time
• Do question 14

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• Do questions 1-14 on chapter 10 in your IB Study
  Guide and turn in