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Chapter 20 Oxidation-Reduction Reactions

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Title: Chapter 20 Oxidation-Reduction Reactions


1
CHAPTER 20
Oxidation-Reduction Reactions
LEO SAYS GER
Pre-AP Chemistry Charles Page High School Stephen
L. Cotton
2
Section 20.1The Meaning of Oxidation and
Reduction (called redox)
  • OBJECTIVES
  • Define oxidation and reduction in terms of the
    loss or gain of oxygen, and the loss or gain of
    electrons.

3
Section 20.1The Meaning of Oxidation and
Reduction (Redox)
  • OBJECTIVES
  • State the characteristics of a redox reaction
    and identify the oxidizing agent and reducing
    agent.

4
Section 20.1The Meaning of Oxidation and
Reduction (Redox)
  • OBJECTIVES
  • Describe what happens to iron when it corrodes.

5
Oxidation and Reduction (Redox)
  • Early chemists saw oxidation reactions only as
    the combination of a material with oxygen to
    produce an oxide.
  • For example, when methane burns in air, it
    oxidizes and forms oxides of carbon and hydrogen,
    as shown in Fig. 20.1, p. 631

6
Oxidation and Reduction (Redox)
  • But, not all oxidation processes that use oxygen
    involve burning
  • Elemental iron slowly oxidizes to compounds such
    as iron (III) oxide, commonly called rust
  • Bleaching stains in fabrics
  • Hydrogen peroxide also releases oxygen when it
    decomposes

7
Oxidation and Reduction (Redox)
  • A process called reduction is the opposite of
    oxidation, and originally meant the loss of
    oxygen from a compound
  • Oxidation and reduction always occur
    simultaneously
  • The substance gaining oxygen (or losing
    electrons) is oxidized, while the substance
    losing oxygen (or gaining electrons) is reduced.

8
Oxidation and Reduction (Redox)
  • Today, many of these reactions may not even
    involve oxygen
  • Redox currently says that electrons are
    transferred between reactants
  • Mg S ? Mg2 S2-

(MgS)
  • The magnesium atom (which has zero charge)
    changes to a magnesium ion by losing 2 electrons,
    and is oxidized to Mg2
  • The sulfur atom (which has no charge) is changed
    to a sulfide ion by gaining 2 electrons, and is
    reduced to S2-

9
Oxidation and Reduction (Redox)
Each sodium atom loses one electron
Each chlorine atom gains one electron
10
LEO says GER
Lose Electrons Oxidation
Sodium is oxidized
Gain Electrons Reduction
Chlorine is reduced
11
LEO says GER
- Losing electrons is oxidation, and the
substance that loses the electrons is called the
reducing agent. - Gaining electrons is
reduction, and the substance that gains the
electrons is called the oxidizing agent.
Mg(s) S(s) ? MgS(s)
Mg is oxidized loses e-, becomes a Mg2 ion
Mg is the reducing agent
S is reduced gains e- S2- ion
S is the oxidizing agent
12
Oxidation and Reduction (Redox)
  • Conceptual Problem 20.1, page 634
  • It is easy to see the loss and gain of electrons
    in ionic compounds, but what about covalent
    compounds?
  • In water, we learned that oxygen is highly
    electronegative, so
  • the oxygen gains electrons (is reduced and is
    the oxidizing agent), and the hydrogen loses
    electrons (is oxidized and is the reducing agent)

13
Not All Reactions are Redox Reactions
- Reactions in which there has been no change in
oxidation number are NOT redox reactions.
Examples
14
Corrosion
  • Damage done to metal is costly to prevent and
    repair
  • Iron, a common construction metal often used in
    forming steel alloys, corrodes by being oxidized
    to ions of iron by oxygen.
  • This corrosion is even faster in the presence of
    salts and acids, because these materials make
    electrically conductive solutions that make
    electron transfer easy

15
Corrosion
  • Luckily, not all metals corrode easily
  • Gold and platinum are called noble metals because
    they are resistant to losing their electrons by
    corrosion
  • Other metals may lose their electrons easily, but
    are protected from corrosion by the oxide coating
    on their surface, such as aluminum Figure 20.7,
    page 636
  • Iron has an oxide coating, but it is not tightly
    packed, so water and air can penetrate it easily

16
Corrosion
  • Serious problems can result if bridges, storage
    tanks, or hulls of ships corrode
  • Can be prevented by a coating of oil, paint,
    plastic, or another metal
  • If this surface is scratched or worn away, the
    protection is lost
  • Other methods of prevention involve the
    sacrifice of one metal to save the second
  • Magnesium, chromium, or even zinc (called
    galvanized) coatings can be applied

17
Section 20.2Oxidation Numbers
  • OBJECTIVES
  • Determine the oxidation number of an atom of any
    element in a pure substance.

18
Section 20.2Oxidation Numbers
  • OBJECTIVES
  • Define oxidation and reduction in terms of a
    change in oxidation number, and identify atoms
    being oxidized or reduced in redox reactions.

19
Assigning Oxidation Numbers
  • An oxidation number is a positive or negative
    number assigned to an atom to indicate its degree
    of oxidation or reduction.
  • Generally, a bonded atoms oxidation number is
    the charge it would have if the electrons in the
    bond were assigned to the atom of the more
    electronegative element

20
Rules for Assigning Oxidation Numbers
  • The oxidation number of any uncombined element is
    zero.
  • The oxidation number of a monatomic ion equals
    its charge.

21
Rules for Assigning Oxidation Numbers
  • The oxidation number of oxygen in compounds is
    -2, except in peroxides, such as H2O2 where it is
    -1.
  • The oxidation number of hydrogen in compounds is
    1, except in metal hydrides, like NaH, where it
    is -1.

22
Rules for Assigning Oxidation Numbers
  • The sum of the oxidation numbers of the atoms in
    the compound must equal 0.

2(1) (-2) 0 H O
(2) 2(-2) 2(1) 0 Ca O H
23
Rules for Assigning Oxidation Numbers
  • The sum of the oxidation numbers in the formula
    of a polyatomic ion is equal to its ionic charge.

X 4(-2) -2 S O
X 3(-2) -1 N O
thus X 6
thus X 5
24
Reducing Agents and Oxidizing Agents
  • Conceptual Problem 20.2, page 641
  • An increase in oxidation number oxidation
  • A decrease in oxidation number reduction

Sodium is oxidized it is the reducing agent
Chlorine is reduced it is the oxidizing agent
25
Trends in Oxidation and Reduction
  • Active metals
  • Lose electrons easily
  • Are easily oxidized
  • Are strong reducing agents
  • Active nonmetals
  • Gain electrons easily
  • Are easily reduced
  • Are strong oxidizing agents

Conceptual Problem 20.3, page 643 Technology
Society page 644
26
Section 20.3Balancing Redox Equations
  • OBJECTIVES
  • Describe how oxidation numbers are used to
    identify redox reactions.

27
Section 20.3Balancing Redox Equations
  • OBJECTIVES
  • Balance a redox equation using the
    oxidation-number-change method.

28
Section 20.3Balancing Redox Equations
  • OBJECTIVES
  • Balance a redox equation by breaking the
    equation into oxidation and reduction
    half-reactions, and then using the half-reaction
    method.

29
Identifying Redox Equations
  • In general, all chemical reactions can be
    assigned to one of two classes
  • oxidation-reduction, in which electrons are
    transferred
  • Single-replacement, combination, decomposition,
    and combustion
  • this second class has no electron transfer, and
    includes all others
  • Double-replacement and acid-base reactions

30
Identifying Redox Equations
  • In an electrical storm, nitrogen and oxygen
    react to form nitrogen monoxide
  • N2(g) O2(g) ? 2NO(g)
  • Is this a redox reaction?
  • If the oxidation number of an element in a
    reacting species changes, then that element has
    undergone either oxidation or reduction
    therefore, the reaction as a whole must be a
    redox.
  • Conceptual Problem 20.4, page 647

YES!
31
Balancing Redox Equations
  • It is essential to write a correctly balanced
    equation that represents what happens in a
    chemical reaction
  • Fortunately, two systematic methods are
    available, and are based on the fact that the
    total electrons gained in reduction equals the
    total lost in oxidation. The two methods
  • Use oxidation-number changes
  • Use half-reactions

32
Using Oxidation-Number Changes
  • Sort of like chemical bookkeeping, you compare
    the increases and decreases in oxidation numbers.
  • start with the skeleton equation
  • Step 1 assign oxidation numbers to all atoms
    write above their symbols
  • Step 2 identify which are oxidized/reduced
  • Step 3 use bracket lines to connect them
  • Step 4 use coefficients to equalize
  • Step 5 make sure they are balanced for both
    atoms and charge Problem 20.5, 649

33
Using half-reactions
  • A half-reaction is an equation showing just the
    oxidation or just the reduction that takes place
  • they are then balanced separately, and finally
    combined
  • Step 1 write unbalanced equation in ionic form
  • Step 2 write separate half-reaction equations
    for oxidation and reduction
  • Step 3 balance the atoms in the half-reactions

(More steps on the next screen.)
34
Using half-reactions
  • continued
  • Step 4 add enough electrons to one side of each
    half-reaction to balance the charges
  • Step 5 multiply each half-reaction by a number
    to make the electrons equal in both
  • Step 6 add the balanced half-reactions to show
    an overall equation
  • Step 7 add the spectator ions and balance the
    equation
  • Rules shown on page 651 bottom
  • Conceptual Problem 20.6, page 652

35
Choosing a Balancing Method
  • The oxidation number change method works well if
    the oxidized and reduced species appear only once
    on each side of the equation, and there are no
    acids or bases.
  • The half-reaction method works best for
    reactions taking place in acidic or alkaline
    solution.

36
End of Chapter 20
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