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Redox Reactions: a rxn that involves an exchange of electrons between two species
Oxidation: increase in oxidation number, loss of electrons, becomes more +
Reduction: decrease in oxidation number, gain of electrons, becomes more LEO goes GER or OIL RIG
NO NET CHANGE IN # OF ELECTRONS
Oxidizing Agent: ion or molecule that accepts e-, brings about oxidation, itself is reduced
the reactant containing the element that is reduced
Reducing Agent: species that donates e-, brings about reduction, itself is oxidized
the reactant containing the element that is oxidized
A. Oxidation Number
RULES FOR ASSIGNING OXIDATION NUMBERS
1. The oxidation number of any free element is 0.
2. The oxidation number of a monatomic ionic is equal to the charge on the ion.
3. The oxidation number of each hydrogen atom in most compounds is 1 +,
exception in hydrides.
4. The oxidation number of each oxygen atom in most compounds is 2-, exception
is peroxides.
5. The sum of the oxidation numbers of all the atoms in a particle must equal the
apparent charge of that particle.
6. In compounds, the elements of Group IA and Group IIA and aluminum have
positive oxidation numbers numerically equal to their group number in the
periodic table.
1. Pseudocharge assigned according to arbitrary rules.
2. Oxidation number of element in elementary substance (e. g.; F 2, O2 ) is 0.
3. Oxidation number of element in monatomic ion is the charge of that ion:
a. Al is 3+ in Al2O3
b.O is 2- in Al2O3
4. Oxidation number of Group I elements in their compounds is 1+
a. Group II is 2+
b. Group VII is 1c. O is 2-, exception in peroxides
d. H is 1 +, exception in hydrides
5. Sum of oxidation numbers of all the atoms in a particle must equal the apparent
charge of that particle.
a. oxidation number of sulfur in Na2SO4
6+
b. oxidation number of Mn in MnO47+
Steps to Balancing Redox Reactions using the Ion-Electron Method
ION-ELECTRON METHOD
for
Balancing Redox Equations
1) Write unbalanced ionic equations for the two half-reactions.
2) Balance atoms other than H and O.
3) Balance O with H2O.
4) Balance H with H+.
5) Balance charge with appropriate number of electrons.
6) If in acidic solution, then skip to step 10.
7) If the reaction is occurring in basic solution, the hydrogen ions (H +) must be
neutralized by adding equal numbers of OH- ions as there is H+ ions to BOTH
sides of the equations.
8) Combine H+ ions and OH- ions to from water.
9) Subtract water where possible.
10)Rewrite the balanced half reactions.
11)Multiply balanced half-reactions by appropriate coefficients so that the numbers
of electrons are equal.
12)Add the half-reactions together.
13)Cancel species that appear on both sides to get the balanced Net Ionic Equation.
14)If necessary, add spectator ions to get the balanced molecular equation.
Check the Final Balance (atoms and charges)!