IB Chemistry
*Topics 9, 19  Oxidation & Reduction/Electrochemistry
9.1 Introduction to Oxidation and Reduction
Historical definitions:
*Oxidation  to combine with ___________________ [e.g. 4Fe(s) + 3O2(g) 
]
*Reduction the opposite of oxidation (i.e. the formation of a ______________ from its compounds / the
process of obtaining a metal from an ore) [e.g. 2Fe2O3(s) 
]
*Ore  a naturally occurring solid compound or mixture of compounds from which a _________ can be extracted
Modern Definitions: (*based on the idea of ________________ transfers)
*Oxidation  the ________of electrons (LEO =
)
*Reduction  the ________ of electrons (GER=
)
*Redox Reaction (i.e Reduction – Oxidation rxn)  a rxn involving both an _______________ and a ________________
Example: Redox Reactions (e.g. a Single Displacement Rxn)
Ex.1: Overall Equation: Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
 Total ionic equ’n:
 Net ionic equ’n:
 Oxidation half – reaction:
[Oxidation(LEO) = e- on ________ side]
 Reduction half – reaction:
[Reduction(GER) = e- on ________side]
 Sum of half –reactions =
More Important Terms:
*The oxidizing agent (OA) is the entity that causes _____________; it is the entity that removes electrons from another; it is the entity
that gets __________________. In the above example, the OA =_______.
*The reducing agent (RA) is the entity that causes _____________; it is the entity that donates electrons to another; it is the entity
that gets ________________. In the above example, the RA = _______.
Common Oxidizing Agents
MnO4- (permanganate ion)
Acidified Cr2O72- (dichromate ion)
H2O2 (hydrogen peroxide)
Common Reducing Agents
NaHSO3 (sodium hydrogen sulfite)
LiAlH4 (lithium aluminum hydride)
S2O32- (
*A half -reaction is a balanced equation that shows the number of _____________ involved in either oxidation
or reduction. *Two half – reactions are needed to represent a ____________ reaction. The two half –reaction
“add up to” the ________________ redox equation.
Practice: a) Write balanced half - reactions for the following. b) Identify the OA and RA in each.
a) Al(s) + Fe3+(aq)  Al3+(aq) + Fe(s)
b) Cd(s) + 2Ag+(aq)  Cd2+(aq) + 2Ag(s)
c) 3Zn(s) +Fe2(SO4)3(aq)  3ZnSO4(aq) +2Fe(s)
OXIDATION NUMBERS/STATES (What are Oxidation Numbers?)
 Oxidation numbers are used to keep track of the electron _________________ taking place in redox reactions.
 For reactions involving covalent reactants and products, you cannot use ionic charges to describe the transfer of electrons.
However, to assign oxidation numbers to the elements in a covalent molecule or polyatomic ion, you can pretend the bonds are ionic.
 Oxidation numbers, therefore, are hypothetical charges, assigned using a set of arbitrary rules. In a covalently bonded molecule or
polyatomic ion, the more electronegative atoms are considered to be negative and the less electronegative atoms are considered to be
positive.
 The oxidation state of an atom in an entity is defined as the apparent net electric charge that an atom would have if electron pairs in
covalent bonds belonged entirely to the more electronegative atom. [e.g Consider H2O]
*Common Oxidation Numbers (Table 1 p. 658)
(**Memorize this table.)
Atom or ion
Oxidation #
1. All atoms in elements
Examples
Na in Na(s)
Br in Br2(l)
P in P4(s)
2. a) Hydrogen in all compounds
H in HCl
H in H2S
H in CH4
b)Except hydrogen in metal
H in NaH
hydrides
H in CaH2
H AlH3
3. a) Oxygen in all compounds
O in H2O
O in Li2O
O in KNO3
O in MnO4b) Except oxygen in peroxides
O in H2O2
4. All monatomic ions
Na+ is +1
S2- is –2
Ca2+ is +2
*Remember: oxidation numbers are not units of electric charge. For this reason, the sign __________ the number.
*Determining Oxidation Numbers
Step 1: Assign _______________ oxidation numbers.
Step 2: The total of the oxidation numbers of atoms in a molecule or ion ________ the value of the net electric charge on the molecule
or ion.
a) The sum of the oxidation numbers for a compound is ________.
b) The sum of the oxidation numbers for a polyatomic ion equal the __________ on the ion.
Step 3: Any unknown oxidation number is determined algebraically from the sum of the known oxidation numbers
and the net charge on the entity.
Oxidation Numbers and Redox Reactions:
*An oxidation involves an ______________ in the oxidation number.
*A reduction involves a _____________ in the oxidation number.
*A redox reaction is one that involves _____________ in oxidation numbers. There are lots of reactions which
are NOT redox reactions (e.g. precipitations, acid/base reactions)!
Reactions that are always going to be redox reactions:
 Combustion reactions – elemental oxygen becoming a compound.
 Synthesis reactions – any element reacting with another to produce a compound.
 Ions changing charge – Fe2+ becoming Fe3+
 Ions of the oxy-acids changing the number of oxygen atoms – SO32- becoming SO42*A fast way to pick up 99% of redox reactions is to look for an element (O2, Cl2, Fe, Na etc.) as a reactant or product. This works
because the oxidation state of any free element is zero, and in a compound, it is unlikely zero.
BALANCING REDOX EQUATIONS:
METHOD #1  USING OXIDATION NUMBERS
Procedure:
Step 1: Assign oxidation numbers and identify the atoms/ions whose oxidation numbers change.
Step 2: Using the change in oxidation numbers, write the number of electrons transferred per atom.
Step 3: Using the chemical formulas, determine the number of electrons transferred per reactant.
Step 4: Calculate the simplest whole number coefficients for the reactants that will balance the total number of electrons transferred.
Balance the reactants and products.
Step 5: Balance the O atoms using H2O(l), and then balance the H atoms using H +(aq)
Examples: Balance the following using the oxidation number method.
1. H2S(g) + O2(g)  SO2(g) + H2O(g)
2. ClO3-(aq) + I2(aq)  Cl-(aq) + IO3-(aq)
3. CH3OH(aq) + MnO4-(aq)  CO32-(aq) + MnO42-(aq)
[In acidic solution.]
[In acidic solution.]
METHOD #2  USING HALF-REACTIONS
“Getting Ready” Procedure for Writing Half-Reactions:
Step 1: Write the chemical formulas for the reactants and products.
Step 2: Balance all atoms, other than O and H.
Step 3: Balance O by adding H2O(l).
Step 4: Balance H by adding H+(aq).
Step 5: Balance the charge on each side by adding e- and cancel anything that is the same on both sides.
*Practice: Balance the following half- reactions. (all in acidic solution)
a) dinitrogen oxide to nitrogen gas
b) nitrite ions to nitrate ions
c) silver oxide to silver metal
d) nitrate ions to nitrous acid
e) hydrogen gas to water
*Balancing Redox Reactions Using Half-Reactions
Procedure:
Step 1: Separate the skeleton equation into the start of two half-reaction equations.
Step 2: Balance each half-reaction equation.
Step 3: Multiply each half-reaction equation by simple whole numbers to balance the electrons lost and gained.
Step 4: Add the half-reaction equations, cancelling the electrons and anything else that is exactly the same on both sides of the
equation.
Examples: Balance the following using the half-reaction method:
1. Fe2+(aq) + CrO72-(aq)  Fe3+(aq) + Cr3+(aq)
2. MnO4-(aq) + C2O42-(aq)  CO2(g) + MnO2(s)
[In acidic solution.]
[In acidic sollution.]