IB Topic 9: Oxidation and Reduction 9'1 Introduction to oxidation and reduction

1
IB Topic 9 Oxidation and Reduction9.1
Introduction to oxidation and reduction

• 9.1.1 Define oxidation and reduction in terms of
  electron loss and gain.
• 9.1.2 Deduce the oxidation number of an element
  in a compound.
• 9.1.3 State the names of compounds using
  oxidation numbers.
• 9.1.4 Deduce whether an element undergoes
  oxidation or reduction in reactions using
  oxidation numbers.

1
2
9.1.1 Define oxidation and reduction in terms of
electron loss and gain.

• Many important chemical reactions involve a
  transfer of electrons
• Mg(s) 2H(aq)? Mg2(aq) H2(g)
• In this reaction, the Mg will _________________
  electrons to become Mg 2
• In this reaction, the 2H will _________________
  electrons to become H2

2
3
9.1.1 Define oxidation and reduction in terms of
electron loss and gain.

• Oxidation a loss of electrons
• Reduction a gain of electrons
• LEO goes GER
• Mg(s) 2H(aq)? Mg2(aq) H2(g)
• Which substance is being oxidized?
• _________
• Which substance is being reduced?
• _________

3
4
You cant have one without the other!

• Reduction (gaining electrons) cant happen
  without an oxidation to provide the electrons.
• You cant have 2 oxidations or 2 reductions in
  the same equation. Reduction has to occur at the
  cost of oxidation

LEO the lion says GER!
ose
lectrons
xidation
ain
lectrons
eduction
GER!
5
9.1.2 Deduce the oxidation number of an element
in a compound.

• All elements in a compound (even covalent
  compounds) can be assigned oxidation numbers. It
  is a convenient form of bookkeeping and will
  assist in balancing complex equations.
• Rules for assigning oxidation numbers (o.n)
• For any atom in its elemental form, the o.n. is
  0.
• Mg, S, H2, Cl2, P4 all are 0
• For any monatomic ion, the o.n. equals the charge
  on the ion.
• Mg2 is 2 Cl- is -1
• The o.n. of oxygen is usually -2 (except
  peroxides, -1)
• The o.n. of hydrogen is 1 when bonded to a
  nonmetal and -1 when bonded to a metal.
• The o.n of F- is -1. The o.n. of the other
  halogens is -1 except when combined with oxygen.
• The sum of the o.n. of all atoms in a neutral
  compound is zero.
• The sum of the o.n. in a polyatomic ion equals
  the charge on the ion.

5
6
9.1.2 Deduce the oxidation number of an element
in a compound.

• Determine the oxidation number (state) of each
  element in the following compounds.
• H2S H is 1 (rule 4) S is -2 (rule 2)
• S8 S is 0 (rule 1)
• SCl2 Cl is -1 (rule 5) S is 2 (rule 6)
• Na2SO3 Na is 1 (rule 2) O is -2 (rule 3)
• S is 4 (rule 6)
• SO42- O is -2 (rule 3) S is 6 (rule 7)
• H2O2 H is 1 (rule 4) O is -1 (rule 3-peroxide)

6
7
9.1.2 Deduce the oxidation number of an element
in a compound.

• Determine the oxidation number (state) of each
  element in the following compounds.
• P2O5 P is _____ O is _____
• NaH Na is _____ H is _____
• Cr2O7-2 Cr is_____ O is _____
• SnBr4 Sn is_____ Br is_____
• HClO4 H is _____ Cl is _____ O is _____
• NO2- N is _____ O is _____
• N2 N is _____
• Ca(NO3)2 Ca is _____ N is _____ O is _____
• BaO2 Ba is _____ O is _____

7
8
9.1.3 State the names of compounds using
oxidation numbers.

• Many metals can have more than one oxidation
  number.
• Exceptions are alkali metals (all 1), alkaline
  earth metals (all 2), zinc (2), aluminum (3)
  silver (1).
• These are identified using Roman numerals to
  denote the charge.
• Iron(II) is Fe2 (ferrous) Iron(III) is Fe3
  (ferric)
• Copper(I) is Cu1 Copper(II) is Cu2
• Name the following compounds using oxidation
  numbers.
• SnCl4 _________________________________
• Cr(NO3)3 _________________________________
• KOH _________________________________
• PbSO4 _________________________________
• CuBr _________________________________

8
9
9.1.3 State the names of compounds using
oxidation numbers.

• Binary covalent compounds can also be named using
  Roman numerals but most use prefixes.
• CO Carbon (II) oxide or carbon monoxide
• CO2 Carbon (IV) oxide or carbon dioxide
• PCl3 Phosphorus (III) chloride or phosphorus
  trichloride
• Name the following compounds using oxidation
  numbers. What are their common names?
• SO3 _________________________________________
• PCl5 _________________________________________
• N2O _________________________________________

9
10
9.1.4 Deduce whether an element undergoes
oxidation or reduction in reactions using
oxidation numbers.

• Deduce what is being oxidized and what is being
  reduced in
• Mg(s) 2HCl(aq) ? MgCl2(aq) H2(g)
• Assign oxidation numbers to both reactants and
  products
• Mg 2HCl ? MgCl2 H2
• Mg0 H1 Cl-1 ? Mg2 Cl-1 H0
• Which species increased in o.n.? That is the one
  being oxidized (losing electrons)
• Mg is going from 0 to 2 so is oxidized
• Which species decreased in o.n.? That is the one
  being reduced (gaining electrons)
• H is going from 1 to 0 so is reduced

10
11
9.1.4 Deduce whether an element undergoes
oxidation or reduction in reactions using
oxidation numbers.

• Deduce what is being oxidized and what is being
  reduced in
• Cu(s) 2AgNO3(aq) ? Cu(NO3)2(aq) 2Ag(s)
• Assign oxidation numbers to both reactants and
  products
• Cu 2AgNO3 ? Cu(NO3)2 2Ag
• Cu0 Ag1 N5 O-2 ? Cu2 N5 O2 Ag0
• Which species increased in o.n.? That is the one
  being oxidized (losing electrons)
• Cu is going from 0 to 2 so is oxidized
• Which species decreased in o.n.? That is the one
  being reduced (gaining electrons)
• Ag is going from 1 to 0 so is reduced

11
12
9.1.4 Deduce whether an element undergoes
oxidation or reduction in reactions using
oxidation numbers.

• Deduce what is being oxidized and what is being
  reduced in
• I2O5(s) 5CO(g) ? I2(s) 5CO2(g)
• Assign oxidation numbers to both reactants and
  products
• I2O5(s) 5CO(g) ? I2(s) 5CO2(g)
• I__ O__ C__ O__? I__ C__ O__
• Which species increased in o.n.? That is the one
  being oxidized (losing electrons)
• Which species decreased in o.n.? That is the one
  being reduced (gaining electrons)

12
13
9.1.4 Deduce whether an element undergoes
oxidation or reduction in reactions using
oxidation numbers.

• Deduce what is being oxidized and what is being
  reduced in
• 2Hg2(aq) N2H4(aq) ? 2Hg(l) ? N2(g)
  4H(aq)
• Deduce what is being oxidized and what is being
  reduced in
• Cl2(g) H2O(l) ? HCl(aq) HClO(aq)

13
14
IB Topic 9 Oxidation and Reduction9.2 Redox
equations

• 9.2.1 Deduce simple oxidation and reduction
  half-equations given the species involved in a
  redox reaction.
• 9.2.2 Deduce redox equations using
  half-reactions.
• 9.2.3 Define the terms oxidizing agent and
  reducing agent.
• 9.2.4 Identify the oxidizing and reducing agents
  in redox equations.

14
15
9.2.1 Deduce simple oxidation and reduction
half-equations given the species involved in a
redox reaction.

• Consider the reaction when copper metal is placed
  in a solution of silver ions.
• Half-reactions
• Copper metal loses 2 electrons Cu ? Cu2
  2e-
• Silver ion gains 1 electron Ag e- ? Ag

15
16
9.2.1 Deduce simple oxidation and reduction
half-equations given the species involved in a
redox reaction.

• The of e- ____ has to equal the of e-
  _________
• The silver half-reaction has to be multiplied by
  2
• Cu ? Cu2 2e-
• 2Ag 2e- ? 2Ag
• Overall

16
17
9.2.1 Deduce simple oxidation and reduction
half-equations given the species involved in a
redox reaction.

• Consider the reaction when iodide ions are added
  to chlorine water.
• Which is more reactive, chlorine or
  iodine?__________
• Since they are nonmetals will the more reactive
  species gain or lose electrons? _________

17
18
9.2.1 Deduce simple oxidation and reduction
half-equations given the species involved in a
redox reaction.

• Consider the reaction when iodide ions are added
  to chlorine water.
• Half-reactions
• 2I- ? I2 2e-
• Cl2 2e- ? 2Cl-

18
19
9.2.1 Deduce simple oxidation and reduction
half-equations given the species involved in a
redox reaction.

• Since electrons lost electrons gained, the
  overall reaction is

19
20
9.2.2 Deduce redox equations using half-reactions.

• Balancing more complex red-ox reactions using
  half-reactions
• Write the unbalanced equation in ionic form.
• Write separate half-reactions for the oxidation
  and reduction processes, adding the appropriate
  electrons.
• Balance the half-reactions. You may add H2O and
  H to balance oxygen and hydrogen as needed.
• Multiply each half-reaction by an appropriate
  number to make the number of electrons equal in
  both.
• Add the half-reactions to show an overall
  equation.
• Add the spectator ions and balance the equation.

20
21
9.2.2 Deduce redox equations using half-reactions.

• S(s) HNO3(aq) ? SO2(g) NO(g) H2O(l)
• Write the unbalanced equation in ionic form.
• Write separate half-reactions for the oxidation
  and reduction processes, adding the appropriate
  electrons.

21
22
9.2.2 Deduce redox equations using half-reactions.

• S(s) HNO3(aq) ? SO2(g) NO(g) H2O(l)
• 3) Balance the half-reactions. You may add H2O
  and H to balance oxygen and hydrogen as needed.
• 4) Multiply each half-reaction by an appropriate
  number to make the number of electrons equal in
  both.

22
22
23
9.2.2 Deduce redox equations using half-reactions.

• S(s) HNO3(aq) ? SO2(g) NO(g) H2O(l)
• 5) Add the half-reactions to show an overall
  equation.
• 6) Add the spectator ions and balance the
  equation.

23
23
24
9.2.2 Deduce redox equations using half-reactions.

• Cr2O72-(aq) Cl-(aq) H(aq) ? Cr3(aq)
  Cl2(g) H2O(l)
• Write the unbalanced equation in ionic form.
• Write separate half-reactions for the oxidation
  and reduction processes, adding the appropriate
  electrons.
• Balance the half-reactions. You may add H2O and
  H to balance oxygen and hydrogen as needed.
• Multiply each half-reaction by an appropriate
  number to make the number of electrons equal in
  both.
• Add the half-reactions to show an overall
  equation.
• 6) Add the spectator ions and balance the
  equation.

24
24
25
9.2.2 Deduce redox equations using half-reactions.

• Cu(s) NO3-(aq) H(aq)? Cu2(aq) NO2(g)
  H2O(l)
• Write the unbalanced equation in ionic form.
• Write separate half-reactions for the oxidation
  and reduction processes, adding the appropriate
  electrons.
• Balance the half-reactions. You may add H2O and
  H to balance oxygen and hydrogen as needed.
• Multiply each half-reaction by an appropriate
  number to make the number of electrons equal in
  both.
• Add the half-reactions to show an overall
  equation.
• 6) Add the spectator ions and balance the
  equation.

25
25
26
9.2.3 Define the terms oxidizing agent and
reducing agent.

• Oxidizing Agent (oxidant)
• Substances that are able to oxidize other
  substances
• They are themselves reduced
• Substances that readily gain electrons
• Reducing Agent (reductant)
• Substances that are able to reduce other
  substances
• They are themselves oxidized
• Substances that readily lose electrons

26
27
9.2.4 Identify the oxidizing and reducing agents
in redox equations.

• Consider the following reactions
• 1) S(s) HNO3(aq) ? SO2(g) NO(g) H2O(l)
• 2) Cr2O72-(aq) Cl-(aq) H(aq) ? Cr3(aq)
  Cl2(g) H2O(l)
• 3) Cu(s) NO3-(aq) H(aq)? Cu2(aq) NO2(g)
  H2O(l)
• For each reaction, state the oxidizing agent and
  the reducing agent

27
27
28
9.2.4 Identify the oxidizing and reducing agents
in redox equations.

• Consider the following reaction
• Cd(s) NiO2(s) H2O(l) ? Cd2(aq) Ni2(aq)
  OH-(aq)
• 1. What is the oxidizing agent?
• 2. What is the reducing agent?
• 3. Balance the equation

28
28
29
IB Topic 9 Oxidation and Reduction9.3 Reactivity

• 9.3.1 Deduce a reactivity series based on the
  chemical behavior of a group of oxidizing and
  reducing agents.
• 9.3.2 Deduce the feasibility of a redox reaction
  from a given reactivity series.

29
30
9.3.1 Deduce a reactivity series based on the
chemical behavior of a group of oxidizing and
reducing agents.
30
31
9.3.1 Deduce a reactivity series based on the
chemical behavior of a group of oxidizing and
reducing agents.

• Activity Series of Metals
• Metals at the top of the chart are easily
  oxidized.
• Metals at the top of the chart are most reactive.
• Metals above hydrogen react with acids.
• Metals will react with ions of metals below them.
• Metals at the top of the chart are good reducing
  agents.
• Ions of metals at the bottom of the chart are
  good oxidizing agents.
• Copper does react with nitric acid but the
  nitrogen in HNO3 is reduced, not the H as in
  other acid-base reactions. Balance the following
  and identify the o.a. r.a.
• Cu HNO3 ? Cu(NO3)2 NO2 H2O

31
32
9.3.1 Deduce a reactivity series based on the
chemical behavior of a group of oxidizing and
reducing agents.

• Activity Series of Halogens
• F2 2e- ? 2F- Decreasing
• Cl2 2e- ? 2Cl- Reactivity
• Br2 2e- ? 2Br-
• I2 2e- ? 2I-
• Halogens at the top of the chart are easily
  reduced.
• Halogens at the top of the chart are most
  reactive.
• Halogens will react with ions of halogens below
  them.
• Halogens at the top of the chart are good
  oxidizing agents.

32
33
9.3.2 Deduce the feasibility of a redox reaction
from a given reactivity series.

• Will an aqueous solution of iron(II) chloride
  oxidize magnesium metal? If so, write the redox
  equation.
• Because Mg is above Fe2 in the activity series,
  we predict the reaction will occur
• Mg(s) Fe2(aq) ? Mg2(aq) Fe(s)
• Will bromine water, Br2(aq) displace Cl- in a
  solution of NaCl? If so, write the redox
  reaction.
• The reaction will not occur since Cl is above Br
  in the activity series.

33
34
9.3.2 Deduce the feasibility of a redox reaction
from a given reactivity series.

• Which of the following metals will be oxidized by
  Pb(NO3)2 Zn, Cu, Fe? Write any redox reaction
  that occurs.
• Predict whether a reaction occurs when the
  following reagents are mixed Cl2(aq) and KI(aq)
  Br2(aq) and LiCl(aq). Write any redox reaction
  that occurs.
• Write balanced chemical equations for the
  following reactions. If no reaction occurs,
  simply write NR
• Zinc metal is added to a solution of silver
  nitrate
• Iron metal is added to a solution of aluminum
  sulfate
• Hydrochloric acid is added to cobalt metal
• Hydrogen gas is bubbled through a solution of
  iron(II) chloride
• Fluorine gas is bubbled through a solution of
  sodium iodide
• Balance the following redox reaction. Identify
  the oxid red agents
• Zn H NO3- ? Zn2 N2O H2O

34
34
35
IB Topic 9 Oxidation and Reduction9.4 Voltaic
cells

• 9.4.1 Explain how a redox reaction is used to
  produce electricity in a voltaic cell.
• 9.4.2 State that oxidation occurs at the negative
  electrode (anode) and reduction occurs at the
  positive electrode (cathode).

35
36
9.4.1 Explain how a redox reaction is used to
produce electricity in a voltaic cell.

• A strip of zinc is placed in a copper solution.
• Write the oxidation reaction
• Write the reduction reaction
• Write the overall reaction
• Describe two observable changes

36
37
9.4.1 Explain how a redox reaction is used to
produce electricity in a voltaic cell.

• This reaction can be used to perform electrical
  work by using a voltaic (galvanic) cell.
• The transfer of electrons takes place through an
  external pathway.
• Metal strips are placed in solutions of their
  ions. The metal strips are connected by a wire
  for flow of electrons.
• The solutions are connected by a salt bridge or
  separated by a porous glass barrier. This
  maintains electrical neutrality.
• Electrons flow from the anode to the cathode.

37
37
38
9.4.2 State that oxidation occurs at the negative
electrode (anode) and reduction occurs at the
positive electrode (cathode).

• At the Zn electrode (anode)
• Oxidation occurs
• Negative electrode
• Electrons are produced and flow through the
  external circuit toward the cathode
• Zn2 ions produced and migrate away from the
  electrode
• Negative ions (anions) from the salt bridge
  migrate into the solution to balance the increase
  in positive charges.

38
38
39
9.4.2 State that oxidation occurs at the negative
electrode (anode) and reduction occurs at the
positive electrode (cathode).

• At the Cu electrode (cathode)
• Reduction occurs
• Positive electrode
• Electrons come from the anode and move into the
  electrode
• Cu2 ions migrate to the electrode and gain
  electrons producing Cu
• Positive ions (cations) from the salt bridge
  migrate into the solution to balance the decrease
  in positive charges.

39
39
40
9.4.2 State that oxidation occurs at the negative
electrode (anode) and reduction occurs at the
positive electrode (cathode).

• A voltaic cell similar to that shown in slide 38
  is constructed. One electrode compartment
  consists of a cadmium strip placed in a solution
  of Cd(NO3)2 and the other has a nickel strip
  placed in a solution of NiSO4. Cadmium is a more
  reactive metal than nickel.
• Write the half-reactions that occur in the two
  electrode compartments. Write the overall
  reaction.
• Which electrode is the anode and which is the
  cathode?
• Indicate the signs of the electrodes.
• Which way do electrons flow?
• In which directions do the cations and anions
  migrate through the solution?

40
40
41
9.4.2 State that oxidation occurs at the negative
electrode (anode) and reduction occurs at the
positive electrode (cathode).

• A voltaic cell similar to that shown in slide 38
  is constructed. One electrode compartment
  consists of a silver strip placed in a solution
  of AgNO3 and the other has a nickel strip placed
  in a solution of NiSO4. Nickel is a more reactive
  metal than silver.
• Write the half-reactions that occur in the two
  electrode compartments. Write the overall
  reaction.
• Which electrode is the anode and which is the
  cathode?
• Indicate the signs of the electrodes.
• Which way do electrons flow?
• In which directions do the cations and anions
  migrate through the solution?

41
41
42
IB Topic 9 Oxidation and Reduction9.5
Electrolytic cells

• 9.5.1 Describe, using a diagram, the essential
  components of an electrolytic cell.
• 9.5.2 State that oxidation occurs at the positive
  electrode (anode) and reduction occurs at the
  negative electrode (cathode).
• 9.5.3 Describe how current is conducted in an
  electrolytic cell.
• 9.5.4 Deduce the products of the electrolysis of
  a molten salt.

42
43
9.5.1 Describe, using a diagram, the essential
components of an electrolytic cell.

• In an electrolytic cell, electricity is supplied
  from an external source and is used to make a
  non-spontaneous reaction take place.
• The substance that conducts electricity in the
  cell is an electrolyte (substance containing
  ions).
• Electrolytes do not conduct when solid because
  ions are not free to move and they have no
  delocalized electrons.
• Electrolytes do conduct when molten or dissolved
  in water because the ions are free to move toward
  opposite charged electrodes

43
44
9.5.1 Describe, using a diagram, the essential
components of an electrolytic cell.

• The electrolyte conducts electricity by the
  movement of ions within it.
• Chemical reactions occur at each electrode so
  that the electrolyte is decomposed in the
  process.
• 2NaCl(l) ? 2Na(l) Cl2(g)

44
44
45
9.5.2 State that oxidation occurs at the positive
electrode (anode) and reduction occurs at the
negative electrode (cathode).

• Oxidation occurs at the positive electrode
  (anode) because negative ions (anions) are
  attracted to it,
• 2Cl-(l) ? Cl2(g) 2e-
• Reduction occurs at the negative electrode
  (cathode) because positive ions (cations) are
  attracted to it.
• Na e- ? Na(l)

45
45
46
9.5.3 Describe how current is conducted in an
electrolytic cell.

• The electrochemical cell is a voltaic cell
  producing electricity from a chemical reaction.
  The anode produces electrons.
• Electrons flow from the anode of the voltaic cell
  to the cathode of the electrolytic cell.
• Positive ions flow toward the cathode and gain
  electrons (become reduced).
• Negative ions flow toward the anode and lose
  electrons (become oxidized).
• Electrons flow from the anode of the electrolytic
  cell to the cathode of the voltaic cell

46
46
47
9.5.4 Deduce the products of the electrolysis of
a molten salt.

• Sketch a cell for the electrolysis of molten
  MgBr2.
• Indicate the directions in which ions and
  electrons move.
• Label the anode and cathode, indicating the
  charge and the type of reaction occurring.
• Give the electrode reactions.

47
47