Oxidation and reduction reactions, also called redox reactions, involve the transfer of electrons between reactants. In redox reactions, one reactant gets oxidized by losing electrons while another reactant gets reduced by gaining electrons. The reactant that loses electrons is the oxidizing agent, while the reactant that gains electrons is the reducing agent. Redox reactions can be identified by changes in oxidation numbers, which are assigned to track electron gains and losses. Balancing redox equations requires splitting the overall reaction into separate half-reactions for oxidation and reduction, then combining the half-reactions by adding electrons to balance charges.

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

2. Section 20.1
The Meaning of Oxidation and Reduction (called “redox”)
• OBJECTIVES
o Define oxidation and reduction in terms of the loss or gain of oxygen, and
the loss or gain of electrons.
o State the characteristics of a redox reaction and identify the oxidizing
agent and reducing agent.
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
CH4 + 2O2  CO2 + 2H2O
• 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
• A process called _______________________________ 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
_______________________, while the substance losing oxygen (or gaining
electrons) is __________________________.
• Today, many of these reactions may not even involve oxygen
• Redox currently says that electrons are __________________________________
Mg + S→ Mg2+ + S2- (MgS)
• The magnesium atom (which has zero charge) changes to a magnesium ion by
losing 2 electrons, and is __________________________ to Mg2+
• The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2
electrons, and is _________________________to S2-
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3. Oxidation and Reduction (Redox)
Each sodium atom loses one electron:
Each chlorine atom gains one electron:
LEO says GER :
___ose ___lectrons = ___xidation
0 +1
Na → Na + e − Sodium is ___________________________
___ain ___lectrons = ___eduction
0 −1
Cl + e − → Cl Chlorine is __________________________
• Losing electrons is oxidation, and the substance that loses the electrons is called the
________________________________________.
• Gaining electrons is reduction, and the substance that gains the electrons is called the
________________________________________.
Example Problem:
What is oxidized and what is reduced in this single replacement reaction? What is the
oxidizing agent? What is the reducing agent?
2AgNO3 + Cu  Cu(NO3)2 + Ag
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4. Oxidation and Reduction (Redox)
• 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 _______________________ and is the
oxidizing agent), and the hydrogen loses electrons (is
___________________________ and is the reducing agent)
Not All Reactions are Redox Reactions
• Reactions in which there has been _________________________________________
__________________________________________________ are NOT redox reactions.
Examples:
+1 +5 −2 +1 −1 +1 −1 +1 +5 −2
Ag N O3 (aq ) + Na Cl (aq ) → Ag Cl ( s ) + Na N O3 (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 )
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5. Section 20.2
Oxidation Numbers
OBJECTIVES
• Determine the oxidation number of an atom of any element in a pure
substance.
• Define oxidation and reduction in terms of a change in oxidation number, and
identify atoms being oxidized or reduced in redox reactions.
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
Rules for Assigning Oxidation Numbers
1) The oxidation number of any uncombined element is ________________
2) The oxidation number of a monatomic ion _________________________________.
2Na + Cl2  2NaCl
3) The oxidation number of oxygen in compounds is ________, except in peroxides, such
as H2O2 where it is -1.
4) The oxidation number of hydrogen in compounds is ________, except in metal
hydrides, like NaH, where it is -1.
H2O
5) The sum of the oxidation numbers of the atoms in the compound _________________
H20 Ca(OH)2
6) The sum of the oxidation numbers in the formula of a polyatomic ion is ___________
_______________________________________________.
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6. NO3 SO4
Example Problem: What is the oxidizing number of each kind of atom in the following
compounds?
a. SO2 b. CO32- c. K2SO4
Reducing Agents and Oxidizing Agents
• An increase in oxidation number = ____________________________________
• A decrease in oxidation number = _____________________________________
Na  Na + e-
Sodium is ____________________________ - it is the ___________________________
Cl + e-  Cl
Chlorine is ____________________________- its is the __________________________
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
Example Problem:
Use the changes in oxidation number to identify which atoms are oxidized and which are
reduced in each reaction.
a. Cl2 + 2HBr  2HCl + Br2
b. C + O2  CO2
Name ___________________________________ Date __________________________
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7. 10.1 and 10.2 Section Review
1. State the characteristics of a redox reaction, and explain how to identify the oxidizing
agent and the reducing agent?
2. Determine which reactant is oxidized and which reactant is reduced.
a. H2 + Cl2  2HCl
b. S + Cl  SCl2
c. N2 + 2O2  2NO2
d. 2Li + F2  2LiF
e. H2 + S  H2S
3. In the equations above a-e, determine which reactant is the reducing agent and which is
the oxidizing agent.
Section 20.3
Balancing Redox Equations
OBJECTIVES
• Describe how oxidation numbers are used to identify redox reactions.
• Balance a redox equation using the oxidation-number-change method.
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8. • Balance a redox equation by breaking the equation into oxidation and
reduction half-reactions, and then using the half-reaction method.
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
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 _______________________.
Example Problem
Use the change in oxidation number to identify whether each reaction is a redox reaction
or a reaction of some other type. If a reaction is a redox reaction, identify the element
reduced, the element oxidized, the reducing agent, and the oxidizing agent.
a. N2O4  2NO2
b. Cl2 +NaBr  2NaCl + Br2
c. PbCl2 + K2SO4  2KCl + PbSO4
d. 2NaOH + H2SO4  Na2SO4 + 2H2O
Balancing Redox Equations
• It is essential to write a correctly balanced equation that represents what happens
in a chemical reaction
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9. • Fortunately, two systematic methods are available, and are based on the
fact that the total electrons gained in reduction _________________ the
total lost in oxidation. The two methods:
1. Use oxidation-number changes
2. Use half-reactions (you only need to know this method)
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
• 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 665 – 666
Example problem:
Use the half reaction method to balance the equation for the following redox reaction.
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