OXIDATION - REDUCTION REACTIONS a.k.a Redox Reactions
Textbook: 4.4-4.6 & 20.1-20.2
Good website: www.wfu.edu/~ylwong/redox/
Oxidation – The __________________________ by a substance (element, compound, ion)
Reduction – The _________________________ by a substance (element, compound, ion)
O. I. L. R. I. G.
L.E.O. the lion says G.E.R
Oxidation --------->
A-2 ---------> A-1 + electron
A-1 ---------> A + electron
---------> A+1 + electron
A
+1
---------> A+2 + electron
A
<--------- Reduction
Redox reaction – a process where electrons are transferred from one substance to another
How can you tell when a redox reaction is taking place?
1) Assign __________________________________to atoms in substances.
2) ________________________________________________________________
___________________________ to determine if the atom has lost or gained electrons.
Rules For Assigning Oxidation Numbers (Use the rules in order)
1.
The oxidation number of an atom in an element is 0.
Examples: Na, H2, Br2, S8, Ne
Ox. #
2.
The oxidation number of a monatomic ion is the same as its charge.
Examples: Na+1, Ca+2, Al+3, Cl-1, O-2
Ox #
3.
The sum of the oxidation numbers of all atoms in a neutral compound is zero.
4.
The sum of the oxidation numbers of all atoms in an ion is equation to the charge on the
ion.
5.
In compounds, fluorine is always assigned an oxidation number of -1. (The most
electronegative element in a compound always has a negative oxidation number.)
6.
Hydrogen’s oxidation number will be:
+1 when bonded to a nonmetal (HCl)
-1 when bonded to a metal (NaH)
Examples: Na—H, H—Ca—H, H—Cl,
Ox #
H—S—H
7.
Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be 1 and when combined with only F, it will be +2 and in O2 it will be 0
Examples: H—O—H, Ca—O, H—O—O—H, [O—O]-2, F—O—F ,
O2
Ox #
8.
Halogens usually have an oxidation number of -1.
Examples: Na—Cl, I—Mg—I, Cl—O—Cl, H—O—Br
Ox #
**If none of the above rules help you get started…look for an atom with a known charge and use
that charge as its oxidation number.
Example: Cd-S
Ox #
Use algebra to determine oxidation numbers of “difficult” atoms:
Example:
H2SO4
H is +1 x 2 = +2
O is -2 x 4 = -8
Example:
ClO4-1
O is -2 x 4 = -8
Example:
NH4+1
H is +1 x 4 = +4
Example:
FeSO4
O is -2 x 4 = -8
Example:
C3H8
H is +1 x 8 = +8
Oxidation Numbers do NOT have to be integers!!
Once oxidation numbers have been assigned, compare them before and after the reaction.
4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s)
Fe is ____________________________________
O is ______________________________________
Notice that a total of 12 electrons were lost and 12 electrons were gained
2 Fe2O3 (s)
+
3 C (s)

4 Fe (s)
+
3 CO2 (g)
Fe is _____________________________________
C is ______________________________________
O ______________________________________
12 electrons lost and 12 electrons gained
As seen in the above examples, oxidation and reduction ______________ occur together
Reducing agent
Oxidizing agent
Examples: Assign oxidation numbers, indicate what is oxidized and reduced, indicate what is
the oxidizing agent and reducing agent
1)
Ca (s)
2)
2 Fe+2 (aq)
In general,
+
2 H+1
+

Cl2 (aq)
Ca+2 (aq)

+
H2 (g)
2 Fe+3 (aq)
+
2 Cl-1 (aq)
________________________________ act as reducing agents (are oxidized)
________________________________ act as oxidizing agents (are reduced).
Periodic table – in general, metals on left of table are more active, metals become less active as
you move to the right side of the table
Predicting the Products of Redox Reactions
The simple ones you already know:
• Decomposition (except of acids)
• Composition (except of acids)
•
Combustion
•
Replacement
A Closer look at Replacement Reactions:
Replacement Reactions are redox reactions.
General pattern: A + BX  AX + B
Example:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
_________ is oxidized and the ___________ is reduced.
Example:
Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s)
The net ionic equation shows the redox chemistry well:
Fe(s) + Ni2+(aq)  Fe2+(aq) + Ni(s)
____________ is oxidized to ___________
_____________ is reduced to ___________
Always keep in mind that whenever one substance is oxidized, some other substance must be
reduced.
The Activity Series
The activity series is a list of metals in order of decreasing ease of oxidation.
The metals at the top of the activity series are called active metals and are easily oxidized – they
want to be ions.
The metals at the bottom of the activity series are called noble metals and NOT easily oxidized –
they want to be atoms.
A metal in the activity series can only be oxidized by a metal ion below it.
If we place Cu into a solution of Ag+ ions, then Cu2+ ions can be formed because Cu is above Ag
in the activity series:
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
or
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
More on this later !
Special Redox Reactions (Memorize – they don’t follow the expected rule)
1. Hydrogen reacts with a hot metallic oxide to produce the metal element and water.
Ex: H2 + MgO  Mg + H2O
2. A metal sulfide reacts with oxygen to produce the metallic oxide and sulfur dioxide.
Ex: 2MgS + 3O2  2MgO + 2SO2
3. Chlorine gas reacts with dilute sodium hydroxide to produce sodium hypochlorite, sodium chloride,
and water.
Cl2 + 2NaOH  NaClO + NaCl + H2O
4. Copper reacts with concentrated sulfuric acid to produce copper(II) sulfate, sulfur dioxide, and water.
Cu + 2H2SO4  CuSO4 + SO2 + 2H2O
5. Copper reacts with dilute nitric acid to produce copper(II) nitrate, nitrogen monoxide, and water.
Cu + HNO3  Cu(NO3)2 + NO + H2O
6. Copper reacts with concentrated nitric acid to produce copper(II) nitrate, nitrogen dioxide, and water.
Cu + HNO3  Cu(NO3)2 + NO2+ H2O
Memorize… These are reduction reactions (oxidation number decreases)
Memorize… These are oxidation reactions (oxidation number increases)
Redox HW #1
Name___________________________________________Score_______
Assign oxidation numbers to each element in each of the following:
SnCl4
CrO3
VOCl3
V2O3
HNO3
FeSO4
SO3
COCl2
CH2Cl2
KClO3
Mn2O7
OsO4
H2PtCl4
ClO3-1
SO3-2
C2O4-2
MnO4-1
S2O3-2
Fe3O4
I3-1
KMnO4
Fe2O3
NiO2
XeOF4
SF4
P4O6
(NH4)2HPO4
CO
Na2C2O4
HBr
HBrO
Br2
HBrO4
BrF3
Li3N
NO2
NO2-1
NO3-1
N2
N2O
Some of the following reactions are NOT redox reactions so leave the blanks empty for those.
Ca (s) + Sn+2 (aq)  Ca+2 (aq) + Sn (s)
Sub. Oxidized
____________
Sub. Reduced
____________
Oxid. Agent
___________
Red. Agent
__________
Oxid. Agent
___________
Red. Agent
__________
Si (s) + 2 Cl2 (g)  SiCl4 (l)
Sub. Oxidized
____________
Sub. Reduced
____________
Cl2 (g) + 2 NaBr (aq)  Br2 (aq) + 2 NaCl (aq)
Sub. Oxidized
____________
Sub. Reduced
____________
Oxid. Agent
___________
Red. Agent
__________
CH4
+

2O2
CO2
+
Sub. Oxidized
____________
Zn
+
Sub. Reduced
____________
2HCl

ZnCl2
Sub. Oxidized
____________
Cr2O72-

Sub. Oxidized
____________

2H2O2
2H2O
+

2Ag+
+
+
NH3
+
Sub. Reduced
____________
SiCl4

2H2O
Sub. Oxidized
____________
H2O
Oxid. Agent
___________
Red. Agent
__________
Oxid. Agent
___________
Red. Agent
__________
Oxid. Agent
___________
Red. Agent
__________
Oxid. Agent
___________
Red. Agent
__________
Oxid. Agent
___________
Red. Agent
__________
NH4Cl
Sub. Oxidized
____________
+
Red. Agent
__________
Cu2+
Sub. Reduced
____________

Oxid. Agent
___________
Cu
2Ag
Sub. Oxidized
____________
HCl
+
Sub. Reduced
____________

Red. Agent
__________
O2
+
Sub. Oxidized
____________
Cu
2CrO42-
Sub. Reduced
____________
CuCl2
Oxid. Agent
___________
H2
Sub. Reduced
____________
Sub. Oxidized
____________
2CuCl
+
Sub. Reduced
____________
2OH-
+
2H2O
4HCl
+
Sub. Reduced
____________
SiO2
Oxid. Agent
___________
Red. Agent
__________