Redox
Reactions
 Have
you ever
drank from an
aluminum can?
 Ever
used a
flashlight?
 Use
your calculator
on a test recently?
 Enjoy
 Are
exercising?
you alive?
What do all of these
things have in common?
ENERGY!!!
And all forms of energy harnessing require an
understanding of Redox Reactions.
Redox Notes
Part I: Define oxidation and
reduction.
Oxidation-Reduction Reactions
 Aka
Redox Rxns
 Chemical
reactions that transfer electrons from one
chemical to another at the same time

OXIDATION – the LOSS of electrons in a chemical rxn

REDUCTION – the GAIN of electrons in a chemical rxn
Redox is a tandem process
 Oxidation
cannot happen without
reduction & vice versa
 Thus,
you will have two types of agents (chemicals)

OXIDIZING AGENT – substance that oxidizes another.
It gets reduced.

REDUCING AGENT - substance that reduces another.
It gets oxidized.
Why can oxidation not
happen without reduction?

One substance that donates the
electrons needs a place for the electrons
to travel to.
 Electrons
don’t just vanish; they attach to
another atom.
 Therefore,
one substance donates and
the other accepts the electron.
LEO the Lion
and his Little e’s
Meet Leo’s family. He loves
them very much.
LEO goes
hunting and
loses his little
e’s.
LEO goes GER
Listen to him ROAR!!
LEO’s pride
searches
everywhere…
And finally finds
them playing at
the watering
hole
LEO gains his
little ‘e’s back
again.
LEO the lion goes GER
Loss of Electrons = Oxidation
Gain of Electrons = Reduction
Example 1:
 Identify
the movement of the electron(s)
and then label the reaction as oxidation
or reduction:
Al
 Al+3 + 3 e-
Example 2:
 Identify
the movement of the electron(s)
(If +  - …gained e-. If -  +…lost e-) and
then label the reaction as oxidation or
reduction:
2Br -
+ Cl2  Br2 + 2Cl -
REDOX
Atoms of elements, ions or
compounds
outermost
gain or lose
electrons
during reactions that form a new set
of elements, ions or compounds.
 SC.912.P.8.10:
Describe oxidationreduction reactions in living
and non-living systems.
Energy
REDOX Reactions
A transfer of electrons
occurs during all
single replacement or
Combustion reactions
Sometimes in Double
Replacement and
Decomposition reactions

Review definition of oxidation
and reduction.
LEO the lion says GER
 Loss
of Electrons is Oxidation _____
Gain of Electrons is Reduction
OIL RIG
 Oxidation
Is Loss of electrons causing the
oxidation number to pump up to a higher
value.
 Reduction Is Gain of electrons causing a
reduction in the value of the oxidation
number.
Living Systems
 Photosynthesis
and cellular respiration are
biological
examples of
Redox reactions.
 Write the chemical
equation for these
reactions.
Non-Living Systems
 Fires,
rusting, and metals reacting in acid
are also examples of RedOx reactions
 Lets see why these reactions are classified
as “REDOX.”
Think-Pair-Share
 When
unbonded elements react to form
compounds, one of the elements gains
electrons (would that be the metal or the
nonmetal?) while the other loses electrons
(would that be the metal or the nonmetal?).
 Think about the formation of NaCl.
1.
2.
3.
Write your thoughts in your notebook.
Share your ideas with your partner.
Share with the class.
REDOX Reaction Example
Metals
of e-
undergo Oxidation: loss
Ex
Na  Na+ +e-
Ex
Cl2 + 2e- 2Cl-
Reduction:
The
gain of e-
Oxidizing Agent-substance
reduced: Cl2
The Reducing Agent-substance
oxidized: Na
Identify the REDOX reactions.
 1)
__Mg
+ __N2
 __Mg3N2__
 2)
__Fe
+ __O2
 3)
__Ca
+ __ O2  __CaO__
 4)
__H2 + __O2
→
__Fe2O3__ (rust)

→__ H2O__
Identify the REDOX reactions.
 5)
__CH4__+__ O2  __CO2__ + __H2O__
 6)
__Cl2 + __ KI → __KCl__
 7)
__CaCO3__ + __HCl__ → __CaCl2__ + __H2CO3__
__
__
+ __I2
Identify the REDOX reactions.
 8)
 9)
__CH4__
+ __O2 → __ CO2 __+ __ H2O __
__AgNO3__ + __ Cu  __Cu(NO3)2__ + __Ag
__
__
 10) __C6H12O6__ + __O2 → __ CO2 __+ __ H2O __
__
__
END OF CHEMISTRY 1 PPT
 Single
Replacement Reactions are always
redox reactions! Combustion reactions
are always redox reactions! Any time an
oxidation number changes (which means
electrons are gained or lost) during the
reaction, a redox reaction is occurring.
Honors)
Some electron transfers are not as easy
to predict as metal with nonmetal. To
figure out what is oxidized and what is
reduced you can follow a plan.
(
You
must assign oxidation
numbers to all elements in the
reaction.
Identify which elements
oxidation number changed from
reactants to products.
Changes in oxidation number
 Oxidation
number = number of e- gained
or lost by an atom when it forms an ion
K+
 EX:
Br-
2K + Br2  2KBr
 Potassium is oxidzed from 0 to +1
 Bromine is reduced from 0 to -1
Here are the Rules
 1.
The oxidation # of an uncombined
atom = 0.
 2. The oxidation number of a monatomic
ion is equal to the charge on the ion.
 3. The oxidation number of the most
electronegative element in a molecule is
equal to the charge it would have if it
were a ion with noble gas configuration.
 4. F is always -1
 5. O is always -2 (except in peroxides
and when attached to F)
Rules cont…
 6.
H is always +1 (except when attached
to more electronegative metals, Li, Na,
Ca, and Al
 7. Group 1A, 2A, and 3A always have an
oxidation number equal to the group
number (equal to the charge it would
have if it were a ion with noble gas
configuration.)
 8. Sum of all oxidation numbers in a
neutral compound is 0
 9. If not neutral, sum of all oxidation
numbers is equal to the overall charge on
ion
Examples
 Neutral
elements that are not bonded to any
other element have oxidation number of zero.
Examples: Na(s), Cl2(g), Hg(l) all have
oxidation numbers of zero.
 Group
1 metals that are bonded to other
elements have an oxidation number of +1
(positive one).
Examples: Na in NaCl is +1, Li in LiOH is +1
 Group
2 metals that are bonded to other
elements have an oxidation number of +2
(positive two). What would Mg in MgO be?
Examples
 Group
2 metals that are bonded to other elements
have an oxidation number of +2 (positive two).
Examples: Ca in CaCl2 is +2, Ba in Ba(OH)2 is
+2, Therefore, Mg in MgO would be +2
 Oxygen
that is bonded to other elements has an
oxidation number of -2 (negative two) unless it is in
a peroxide or bonded to F.
Examples: O in CO2 is -2,
O in LiOH is -2,
O in Na3PO4 is -2
 BUT
WAIT… THERE ARE EXCEPTIONS
Examples
 Oxygen
that is bonded to other elements has an
oxidation number of -2 (negative two) unless it is
in a peroxide or bonded to F.
Examples: O in CO2 is -2
O in LiOH is -2
O in Na3PO4 is -2
What is O in O2?
What is O in H2O?
 Exception:
O in HOOH, hydrogen peroxide, is -1
What is O in O2? zero.
What is O in H2O? -2
Fluorine is always -1, Why?
 Halogens
that are bonded to other elements
have an oxidation number of -1 (negative
one) unless they are bonded to a more
electronegative element such as a halogen
closer to the top of the periodic table.
WAIT! What happens to oxygen when it is
bonded to fluorine?
 That’s right, oxygen must be positive in this
super rare case!
Examples
 Halogens
that are bonded to other elements have
an oxidation number of -1 (negative one) unless
they are bonded to a more electronegative
element such as a halogen closer to the top of the
periodic table.
Examples: Cl in NaCl is -1
Cl in PCl5 is -1
Cl in CaCl2 is -1
 Exception:
Cl in ClF5 is +5,
Cl in ClBr6 is -6 while Br in ClBr6 is +1
OK – Let’s face it, the larger halogens are only easily
predictable when bonded with a metal, otherwise,
much thinking is required!

Compounds are Neutral







A compound has an overall charge of zero, which means
all the negative charges have to equal the positive
charges.
Examples: When calculating the oxidation number of N in
NO2 , use the rules above to help you. You see that
oxygen normally has an oxidation number of -2 and there
are two oxygen atoms. 2(-2) = -4. The total number of
negative charges is 4 negatives. The only other atom that
is present is nitrogen. That means the nitrogen is
responsible for all for the positive charge.
X + -4 = 0. X = +4. Therefore, the oxidation number on N in
NO2 is +4.
The oxidation number of C in CO (carbon monoxide) is +2.
The oxidation number of C in CO2 (carbon dioxide) is +4.
The oxidation number of P in PCl3 (phosphorous trichloride)
is +3.
The oxidation number of P in P2O5 (diphosphorous
pentoxide) is +5.
Mn in MnO2 is +4.
Polyatomic Ions
 An
ion has an overall charge equal to the
charge of the ion. That means the positive
charges will NOT equal the negative
charges, but instead, when you add all
the charges together the sum will be
equal to the charge of the ion.
 Example:
The Mn in
permanganate ion, MnO4-, is =7
(Here is how: X + 4(-2)= -1)
X + -8
X=7
= -1 add 8 to both sides.
Hydrogen bonded to a metal
is assigned -1, and hydrogen
bonded to nonmetal is +1.
H in HCl is +1
H in BH3 is -1
DO NOW!
Assign oxidation
numbers to each element in the
element, compound or ion.
a) HCl
b) KNO3
OHd) Mg3N2
e) I2
c)
DO NOW!
Assign oxidation
numbers to each element
ClO3g) Al(NO3)3
h) S8
i) H2O2
j) PbO2
f)
More Practice!
Assign oxidation
numbers to each element.
NaHSO4
l) SO32m)O2
n) KMnO4
o) LiH
k)
More Practice!
Assign oxidation numbers to
each element in the
compounds listed below.
Fe2O3
q) SO3
r) NH4+
s) H2SO4
t) Na
p)
Assess yourself!
Assign oxidation numbers to
each element in the
compounds listed below.
Have you MASTERED THIS SKILL?
Try it in the context of a
chemical equation!!
Assign oxidation numbers to
each element
 1)
__Mg
+ __N2
 __Mg3N2__
 2)
__Fe
+ __O2
 3)
__Ca
+ __ O2  __CaO__
 4)
__H2 + __O2
→
__Fe2O3__
→__ H2O__
Assign oxidation numbers to
each element
 5)
__CH4__+__ O2  __CO2__ + __H2O__
 6)
__Cl2 + __ KI → __KCl__
 7)
__CaCO3__ + __HCl__ → __CaCl2__ + __H2CO3__
__
__
+ __I2
Assign oxidation numbers to
each element
 8)
__CH4__
+ __O2 → __ CO2 __+ __ H2O __
 9)
__AgNO3__ + __ Cu  __Cu(NO3)2__ +
__
__
__Ag
Balance using Half Reactions
 Fe
+ CuSO4  Cu + Fe2(SO4)3
 -omit spectator ions-ions that don’t change their
O#
Example: SO4 stays at -2
 Fe + Cu+2 Cu + Fe +3
 Split into half reactions and balance electron
transfer
 Fe  3e-1 + Fe +3
 Cu+2 +2e -1  Cu
Example: 6 electrons is the least common multiple.
Balance using
Half Reactions
Fe + CuSO4  Cu + Fe2(SO4)3
Split into half reactions and balance electron
transfer
 2(Fe  3e-1 + Fe +3)
 3(Cu+2 +2e -1  Cu)
The iron lost a total of six electrons as it was
oxidized
 2Fe  6e-1 + 2Fe +3
 3Cu+2 +6e -1  3Cu
The balanced reaction is

2Fe
+ 3Cu+2  2 Fe +3 + 3Cu
Balance each reaction
1) __Mg
+ __N2
 __Mg3N2__
2) __Fe
+ __O2
3) __Ca
+ __ O2  __CaO__
4) __H2 + __O2
→
__Fe2O3__
→__ H2O__
Balance each reaction.
5) __CH4__+__ O2  __CO2__ + __H2O__
6) __Cl2 + __ KI → __KCl__
+ __I2
7) __CaCO3__ + __HCl__ → __CaCl2__ + __H2CO3__
__
__
Balance each reaction
8) __CH4__
+ __O2 → __ CO2 __+ __ H2O __
9) __AgNO3__ + __ Cu  __Cu(NO3)2__ +
__
__
__Ag
Self Assessment.
What happens during
oxidation and reduction?
What types of reactions
are also redox reactions?
(Honors Extension) How
can you balance redox
equations?