1. CHAPTER 20
“Oxidation-Reduction Reactions”
LEO SAYS GER

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

4. Section 20.1
The 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)
0 0 +1 −1
2 Na + Cl 2 → 2 Na Cl
Each sodium atom loses one electron:
0 +1
−
Na → Na + e
Each chlorine atom gains one electron:
0 −1
−
Cl + e → Cl

10. LEO says GER :
Lose Electrons = Oxidation
0 +1
−
Na → Na + e Sodium is oxidized
Gain Electrons = Reduction
0 −1
−
Cl + e → Cl 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 is the Mg is oxidized: loses e-, becomes a Mg2+ ion
reducing
agent
Mg(s) + S(s) → MgS(s)
S is the oxidizing agent S is reduced: gains e- = S2- ion

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:
+1 +5 −2 +1 −1 +1 −1 +1 +5 −2
Ag N O 3 (aq ) + Na Cl (aq ) → Ag Cl ( s ) + Na N O 3 (aq)
+1 −2 +1 +1 +6 −2 +1 +6 −2 +1 −2
2 Na O H (aq ) + H 2 S O 4 (aq ) → + Na 2 S O 4 (aq ) + H 2 O(l )

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

18. Section 20.2
Oxidation 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 atom’s 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
1) The oxidation number of any
uncombined element is zero.
2) The oxidation number of a
monatomic ion equals its charge.
0 0 +1 −1
2 Na + Cl 2 → 2 Na Cl

21. Rules for Assigning Oxidation Numbers
3) The oxidation number of oxygen in
compounds is -2, except in
peroxides, such as H2O2 where it is -1.
4) The oxidation number of hydrogen in
compounds is +1, except in metal
hydrides, like NaH, where it is -1.
+1 −2
H2O

22. Rules for Assigning Oxidation Numbers
5) The sum of the oxidation numbers of the
atoms in the compound must equal 0.
+1 −2 +2 −2 +1
H2O Ca (O H ) 2
2(+1) + (-2) = 0 (+2) + 2(-2) + 2(+1) = 0
H O Ca O H

23. Rules for Assigning Oxidation Numbers
6) The sum of the oxidation numbers in
the formula of a polyatomic ion is equal
to its ionic charge.
? −2 ? −2
− 2−
N O3 S O4
X + 3(-2) = -1 X + 4(-2) = -2
N O S O
thus X = +5 thus X = +6

24. Reducing Agents and Oxidizing Agents
• Conceptual Problem 20.2, page 641
• An increase in oxidation number = oxidation
• A decrease in oxidation number = reduction
0 +1
−
Na → Na + e
Sodium is oxidized – it is the reducing agent
0 −1
−
Cl + e → Cl
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

26. Section 20.3
Balancing Redox Equations
OBJECTIVES
Describe how oxidation
numbers are used to identify
redox reactions.

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

28. Section 20.3
Balancing 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:
1) oxidation-reduction, in which
electrons are transferred:
• Single-replacement, combination,
decomposition, and combustion
2) 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? YES!
•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

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:
1) Use oxidation-number changes
2) 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. Choosing a Balancing Method
1) 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.
2) The half-reaction method works
best for reactions taking place in
acidic or alkaline solution.

35. 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

36. Electrochemical (Voltaic) Cells
• An apparatus that allows a redox
reaction to occur by transferring
electrons through an external
connector
• Redox reactions that occur
spontaneously may be employed to
provide a source of electrical
energy

37. • When the two half cells of a
redox reaction are connected by
an external conductor and a salt
bridge that allows the migration
of ions, a flow of electrons
(electric current) is produced
• In a voltaic cell, a chemical
reaction is used to produce a
spontaneous electric current by
converting chemical energy to
electrical energy

38. Basic Concepts of
Electrochemical Cells

39. Cathode
Cathode - The electrode where
reduction occurs
• In an electrochemical (voltaic)
cell, the cathode is the
POSITIVE electrode

40. Anode
Anode – the electrode where
oxidation occurs
• In an electrochemical (voltaic)
cell, the anode is the NEGATIVE
electrode
• The anode is the more active
metal (according to table J)

41. Electrolytic Cells
• Sometimes, in combining half reaction, the
potential (Eo) for the overall reaction is
negative. In this case, the reaction will not
take place spontaneously.
• Redox reactions that do not occur
spontaneously can be forced to take place
by supplying energy with an externally
applied electric current
• The use of an electric current to bring
about a chemical reaction is called
electrolysis
• In an electrolytic cell, an electric current
is used to produce a chemical reaction

42. Basic Concepts of
Electrolytic Cells

43. Electrodes
Cathode – negative electrode (reduction
takes place here)
• In electrolytic cells POSITIVE ions
are REDUCED at the cathode

44. Electrodes
Anode – Positive anode (oxidation takes
place here)
• In electrolytic cells NEGATIVE ions
are OXIDIZED at the anode.
**NOTE: The charge of the electrode
is the opposite in electrochemical
cells**