Redox Reactions
Chapter 18
+ O2
Oxidation-Reduction (Redox) Reactions
“redox” reactions: rxns in which electrons are transferred from one species to another
oxidation & reduction always occur simultaneously
we use OXIDATION NUMBERS to keep track of electron transfers
Rules for Assigning Oxidation Numbers:
1) the ox. state of any free (uncombined) element is zero. Ex: Na, S, O2, H2, Cl2, O3
Rules for Assigning Oxidation Numbers:
2) The ox. state of an element in a simple ion is the charge of the ion.
Mg2+ oxidation of Mg is +2
Rules for Assigning Oxidation Numbers:
3) the ox. # for hydrogen is +1
(unless combined with a metal, then it has an ox. # of –1)
Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)
Rules for Assigning Oxidation Numbers:
4) the ox. # of fluorine is always –1.
Rules for Assigning Oxidation Numbers:
5) the ox. # of oxygen is usually –2.
Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide:
H2O2
Rules for Assigning Oxidation Numbers:
6) in any neutral compound, the sum of the oxidation #’s = zero.
Rules for Assigning Oxidation Numbers:
7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.
Rules for Assigning Oxidation Numbers:
**use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:
Examples: Assign oxidation #’s to each element:
a) NaNO3
Examples: Assign oxidation #’s to each element:
b) SO32-
Examples: Assign oxidation #’s to each element:
c) HCO3-
Examples: Assign oxidation #’s to each element:
d) H3PO4
Examples: Assign oxidation #’s to each element:
e) Cr2O72-
Examples: Assign oxidation #’s to each element:
f) K2Sn(OH)6
Definitions
Oxidation: the process of losing electrons (ox # increases)
Reduction: the process of gaining electrons (ox # decreases)
Oxidizing agents: species that cause oxidation (they are reduced)
Reducing agents: species that cause reduction (they are oxidized)
To help you remember…
OIL RIG Oxidation Is Loss Reduction Is Gain
Are all rxns REDOX rxns?
a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.
Examples
MgCO3 MgO + CO2
Examples
Zn + CuSO4 ZnSO4 + Cu
Examples
NaCl + AgNO3 AgCl + NaNO3
Examples
CO2 + H2O C6H12O6 + O2
Balancing Redox Equations
Balancing Redox Equations
In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #)
There are 2 methods for balancing redox equations:1. Change in Oxidation-Number Method2. The Half-Reaction Method
1. Change in Oxidation-Number Method:
based on equal total increases and decreases in oxidation #’s
Steps:
1. Write equation and assign oxidation #’s.
2. Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.
3. Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket.
4. Choose coefficients that make the total increase in ox. # = the total decrease in ox. #.
5. Balance the remaining elements by inspection.
Example
S + HNO3 SO2 + NO + H2O
If needed, reactions that take place in acidic or basic solutions can be balanced as follows:
Acidic: Basic:
add H2O to the side needing oxygen
balance as if in acidic sol’n
then add H+ to balance the hydrogen
add enough OH- to both sides to cancel out each H+ (making H2O) & then cancel out water as appropriate
Example: Balance the following equation, assuming it takes place in acidic solution.
ClO4- + I- Cl- + I2
Example: Balance the following equation, assuming it takes place in basic solution.
ClO4- + I- Cl- + I2
2. The Half-Reaction Method:
separate and balance the oxidation and reduction half-reactions.
Steps:1. Write equation and assign oxidation #’s.2. Determine which element is oxidized and which is reduced,
and determine the change in oxidation # for each.3. Construct unbalanced oxidation and reduction half reactions.4. Balance the elements and the charges (by adding electrons
as reactants or products) in each half-reaction.5. Balance the electron transfer by multiplying the balanced half-
reaction by appropriate integers.6. Add the resulting half-reaction and eliminate any common
terms to obtain the balanced equation.
Example: Balance the following using the half-reaction method:
HNO3 + H2S NO + S +H2O
If needed, reactions that take place in acidic solutions can be balanced as follows:
Acidic: 1. Write separate eq’ns for the oxidation & reduction half-rxns2. For each half-rxn:
a) Balance all the elements except H and Ob) Balance O using H2O
c) Balance H using H+
d) Balance charge using elections3. If necessary, multiply one or both balanced half-rxns by an integer
to equalize the number of electrons transferred in the two half-rxns.4. Add the half-reactions and cancel the identical species (those
appearing in reactants and products)5. Check that the elements and charges are balanced
If needed, reactions that take place in basic solutions can be balanced as follows:
Basic:
1. Balance as if in acidic sol’n (follow ALL steps for acidic redox)
2. Add enough OH- to both sides to cancel out each H+ (making H2O) & then cancel out water as appropriate
3. Check that the elements and charges are balanced
HOMEWORK:
Balance the following using the
half-rxn method…
In acidic sol’n:
a) Cu + NO3- Cu2+ + NO
b) Cr2O72- + Cl- Cr3+ + Cl2
c) Pb + PbO2 + H2SO4 PbSO4
In basic sol’n:
a) Al + MnO4- MnO2 + Al(OH)4
-
b) Cl2 Cl- + OCl-
c) NO2- + Al NH3 + AlO2
-