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Lesson 1
OXIDATION REDUCTION REACTIONS

Earlier in the year we learned that if a copper
strip was placed in silver nitrate, the silver
would be displaced by the copper and solid
silver crystals would form on the copper. The
reaction is classified as a single
displacement reaction.
OXIDATION REDUCTION REACTIONS

Cu(s) + 2 Ag NO3(aq)  2 Ag(s) + Cu(NO3)2(aq)

If we look at the transfer of electrons in this
reaction we will see that Cu looses e- and Ag
gains eCu(s) + 2 Ag+ (aq)  2 Ag(s) + Cu2+ (aq)
Cu becomes Cu2+, a loss of 2 e Each Ag+ becomes Ag, a gain of 2 e

OXIDATION REDUCTION REACTIONS
Oxidation - A process in which chemical entities
lose e Cu was oxidized to Cu2+
OXIDATION REDUCTION REACTIONS
Reduction – A process in which chemical entities
gain e Each Ag+ was reduced to Ag
OXIDATION REDUCTION REACTIONS

reaction in which one reactant is oxidized and
the other is reduced is called a redox reaction.
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Therefore, in any redox reaction, the number of
e- lost must equal the number of e- gained. It
must balance.
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TIP:

To remember which is oxidation and which is
reaction you can use of these memory aids.
 LEO says GER
 Lose Electrons Oxidation
 Gain Electrons Reduction

OR
OIL RIG
 Oxidation Is Loss – Reduction Is Gain
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OXIDIZING AND REDUCING AGENTS

A substance that removes electrons from
another substance is known as a oxidizer or a
Oxidizing agent


A substance that gives electrons to another
substance is known as an reducer or a
reducing agent.
STEPS TO IDENTIFYING OXIDATION AND
REDUCTION
Step 1) Identify repeating entities
 Write the total ionic equation
 Eliminate ions common on both sides
Step 2) Label charges
 Uncombined have a charge of 0
Step 3) Identify loss and gain of electrons
EXAMPLE
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Step 1)
Zn + Cu2+(aq) + SO42-(aq)  Zn2+(aq) + SO42-(aq) + Cu(s)
Step 2)
 Zn0(s) + Cu2+(aq)  Zn2+(aq) + Cu0(s)
Step 3)
 Zn becomes Zn2+, loss of 2 e- (oxidation)
 Cu2+ becomes Cu, gain of 2e- (reduction)
QUESTIONS
Practice page 376
 Questions page 377 # 2-5, 7

REDOX REACTIONS OF NON-METALS
(379-383)
REDOX REACTIONS OF NON-METALS
Not all redox reactions involve metals
 2H2(g) + O2(g)  2H2O(g)
This reaction is similar to the reaction of sodium
with oxygen
 2Na(s) + Cl2(g)  2NaCl(s)
REDOX REACTIONS OF NON-METALS
The main difference is that there are NO
ionic charges.
 Earlier in the year we learned that polar
covalent bonds have partial charges δand δ+.
 The more electronegative an element is
the more it will pull the electrons of the
other element towards it. This apparent
charge is called an oxidation number,
which is the charge it would have if it
gained or lost electrons.

OXIDATION NUMBERS

The sum of all of the oxidation numbers in a
molecule must be zero because the molecule
does not have an overall charge.

If it was a polyatomic ion it would have a charge
but the oxidation numbers can still be
calculated.
OXIDATION NUMBERS
2H2(g) + O2(g)  2H2O(g)
 2Na(s) + Cl2(g)  2NaCl(s)


In the reaction seen above, O is more
electronegative than H. Since oxygen gained 2
e- to from O2- in the ionic compound. We will
assume it has a oxidation # of -2.
OXIDATION NUMBERS
To find the oxidation #of the other
element, solve for the unknown. H2O
 2( H) + (O) = 0
 2(H) + (-2) = 0
 2H = 2
H=1


Therefore, H has an oxidation number of
+1.
OXIDATION NUMBERS
Note:
 The oxidation number of an atom can change if
the atom is involved in a redox reaction.

RULES FOR ASSIGNING OXIDATION #’S
I) The oxidation number of a atom in an
uncombined element is always Zero.
 Example: Na, K, O2, H2, Cl2, S8, Li,
RULES FOR ASSIGNING OXIDATION #’S
II) The oxidation number of a simple ion is the
charge of the ion.
 Example: Ca2+ = + 2
 Na+ = + 1
 Cl- = - 1
RULES FOR ASSIGNING OXIDATION #’S
III) The oxidation number of Hydrogen in most
compounds is +1.
 Example: H2O, H2SO4, NH3, H = +1
 Except in metal hydrides: NaH, H = -1
RULES FOR ASSIGNING OXIDATION #’S
IV) The oxidation number of oxygen in most
compounds is -2
 Example : MgO, HNO3, O = -2
 Except in peroxides: H2O2, O = -1
RULES FOR ASSIGNING OXIDATION #’S
V) The oxidation number of group 1 elements is
+1
 Example: Na, Li, = + 1
The oxidation number of group 2 elements is
+2
 Example: Mg, Ca = + 2

RULES FOR ASSIGNING OXIDATION #’S
VI) The sum of oxidation numbers in a compound
must = 0
 Example: H2O = 2 (+1) + (-2) = 0
RULES FOR ASSIGNING OXIDATION #’S
The sum of oxidation numbers in a polyatomic
ion must equal the charge of the ion.
 Example: OH- = (-2) + (+1) = -1

EXAMPLES
H2SO4:
2(H) + (S) + 4(O) = 0
2 (+1) + (S) +4(-2) = 0
2 + (S) -8 = 0
S-6=0
S=6
EXAMPLES

NO3- :
(N) + 3(O) = -1
(N) + 3(-2) = -1
N – 6 = -1
N=5
IDENTIFYING REDOX REACTIONS USING
OXIDATION NUMBERS



Not all reactions are redox
By using oxidation numbers, you can tell if a
reaction is redox. A redox reaction will have a
change in oxidation numbers where a reaction
that is not redox will not.
Example: AgNO3 + HCl  HNO3 + AgCl
Ag+ + NO3- + H+ + Cl-  H+ + NO3- + Ag+ + Cl-
Questions
 Page 383# 3
THE ACTIVITY SERIES
METAL ACTIVITY SERIES

The Activity series is an arrangement of metals
in order of their tendency to react (become
oxidized).
The most reactive metals (most easily oxidized)
are at the top of the list.
 The least reactive are at the bottom.

METAL ACTIVITY SERIES
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This allows us to predict whether a single
displacement reaction will occur not. For it to
occur, the metal must be higher on the list than
the metal in the compound that it is trying to
displace.
HALOGEN ACTIVITY SERIES
The halogens are also arranged in order of
most reactive to least reactive.
 Unlike the metals, the most reactive are the
easiest to reduce. Fluorine is the strongest
oxidizing agent because of its very high
electronegativity.

ACTIVITY SERIES
Metals
Lithium
Potassium
Calcium
Sodium
Aluminum
Zinc
Chromium
Iron
Nickel
Tin
Lead
Hydrogen *
Copper
Mercury
Silver
Platinum
Gold
Decreasing Activity
Halogens
Fluorine
Chlorine
Bromine
Iodine
The Activity series is used to predict the
products of single displacement reactions.
 A + BC  AC + B
 In general, an element that is higher on the
activity series will displace an element that is
lower. The lower element is, thus left as a pure
metal.

EXAMPLE

Zn(s) + AgNO3(aq)  Zn(NO3)2(aq) + Ag

Is this a redox reaction?

Zn(s) + 2Ag+(aq) + 2NO3-(aq)  Zn2+ 2(NO3)-(aq) + Ag

Zn0(s) + Ag+(aq)  Zn2+ + Ag0

Zn becomes Zn2+, loss of 2 e- (oxidation)

Ag+ becomes Ag, gain of 1e- for 2atoms
(reduction)

Practice page 388

Activity Series Lab