Redox worksheet.docx
Oxidation-Reduction Summary and Study Assignment
The atoms in elements, ions or molecules involved in electrochemical reactions are characterized by the number of
protons the atom has compared to the number of electrons it is assigned; we refer to the difference between the number of
protons and the number of assigned electrons as the oxidation state of the atom. For monatomic ions, the oxidation state of
the ion is simply the ion’s charge. In the case of covalent bonding (electrons are shared) between atoms of differing
electronegativity, the bonding electrons are assigned to the atom with the largest electronegativity when determining the
oxidation state of the bonded atoms. For polyatomic ions the sum of the oxidation states must equal the total charge of the
ion. For molecules, the sum of the oxidation states must equal zero. In electrochemical reactions, also called oxidationreduction or redox reactions, atoms undergo changes in their oxidation state. In other words, redox reactions involve the
transfer of VALENCE electrons from one reactant to another.
Definitions
1.
2.
3.
4.
5.
Oxidation: The loss or apparent loss of electrons; an increase in oxidation number for the atom.
Reduction: The gain or apparent gain of electrons; a decrease in oxidation number for the atom.
Oxidation number (O.N.): A “bookkeeping” number system to determine which atom is oxidized or
reduced. Equals the difference between the number of protons and the number of assigned electrons an atom
has; sign can be – or +.
Oxidizing Agent: Oxidizes another substance. The oxidizing agent is reduced.
Reducing Agent: Reduces another substance. The reducing agent is oxidized.
Oxidation (O.N.) number rules
1.
2.
3.
4.
5.
6.
7.
For an atom in its elemental form (K, S8, N2, Au, etc.) O.N. = 0
For a monatomic ion O.N. = the charge on the ion. (Ti3+ O.N. = +3)
Fluorine is always –1 in compounds.
Cl, Br, and I are –1 as well except when combined with fluorine or oxygen.
Hydrogen is +1, except for the metal hydrides, MHx, where H is –1 (for example calcium hydride, CaH2).
Oxygen is –2 in most compounds. An exception is for peroxides, O22–, where O is –1 each.
The sum of oxidation numbers must equal zero for a neutral compound or equal the ionic charge for a
polyatomic ion.
Examples
FeCl2: Ionic compound. The O.N. of Cl is –1. The O.N. of Fe is +2
C2H6: Covalent compound. The O.N. of H is +1. The O.N. of each C is –3
HSO4–: Polyatomic ion. The O.N. of H is +1. The O.N. of each O is –2. The O.N. of S is +6.
Change in oxidation numbers and redox reactions
In a oxidation-reduction reaction one element is oxidized and one element is reduced. By assigning oxidation numbers
to the elements one can then see which element is oxidized (has an increase in oxidation number) and which element
is reduced (has a decrease in oxidation number).
The oxidized element has a loss or an apparent loss of electrons. The change in oxidation number indicates the number
of electrons lost by the element.
The reduced element has a gain or an apparent gain of electrons. The change in oxidation number indicates the number
of electrons gained by the element.
Examples
1. N2 + 2O2 –> 2NO2
Assign oxidation numbers to each atom:
N2: each N has an O.N. of 0.
O2: each O has an O.N. of 0.
Foothill College Chemistry 1C-Larson/Daley
NO2: O has an O.N. of -2.; N has an O.N. of +4.
1
Last Modified 8/9/2017
Redox Worksheet
Oxidation: N2 is oxidized.
N has an increase in O.N. from 0 to +4. N is oxidized by losing 4 e–.
(N2 is the reducing agent)
Reduction: O2 is reduced.
O has a decrease in O.N. from 0 to –2. O is reduced by gaining 2 e–.
(O2 is the oxidizing agent)
2. ClO3–(aq) + 6H+(aq) + 6Br–(aq) –> Cl–(aq) + 3H2O(l) + 3Br2(l)
Assign oxidation numbers to each atom:
ClO3–: each O has a O.N. of -2.; Cl has a O.N. of +5. Cl–: Cl has a O.N. of -1.
H+: H has a O.N. of +1. H2O: O has a O.N. of -2.; each H has a O.N. of +1.
Br–: Br has a O.N. of -1. Br2: each Br has a O.N. of 0.
Oxidation: Br– is oxidized.
Br has an increase in O.N. from -1 to 0. Br is oxidized by losing 1 e–.
(Br– is the reducing agent)
Reduction: ClO3– is reduced.
Cl has a decrease in O.N. from +5 to –1. Cl is reduced by gaining 6 e–.
(ClO3– is the oxidizing agent)
Balancing Oxidation-Reduction Reactions
Steps for balancing using the half-reaction method
1. Rewrite the equation in NET-IONIC form.
2. Assign oxidation numbers to each element.
3. Identify the oxidized and reduced species from the changes in oxidation number between reactants and products.
4. Divide the reaction into oxidation and reduction half-reactions. Include only the reactants (and their products) that have a
change oxidation number.
5. Balance each half-reaction in acid medium by:
a) balancing elements other than O and H by inspection
b) balancing O atoms by adding H2O to the side with fewer O atoms
c) balancing H atoms by adding H+ ions to the side with fewer H atoms
–
d) balancing charge by adding e to the side with the most positive ionic charge.
6. Find the common multiple of electrons between half-reactions by multiplying each half-reaction by an integer to make
e– lost equal to e– gained.
7.
8.
9.
Add the half-reactions together and cancel identical substances (including phases!) appearing on each side.
If the reaction occurs in basic medium, add OH– ions to both sides and convert each H+ to H2O.
If possible, add the spectator ions back after balancing the overall equation.
Foothill College Chemistry 1C-Larson/Daley
2
Last Modified 8/9/2017
Redox Worksheet
Example of the half-reaction method
Step 1:
KMnO4(aq) + KI(aq) ® MnO2(s) + KIO3(aq)
(NOT BALANCED)
–
–
–
MnO4 (aq) + I (aq) ® MnO2(s) + IO3 (aq)
(Skeleton net ionic equation, not balanced,
spectator ions omitted!)
Step 2 and 3:
Mn changes from a +7 oxidation state to a +4 oxidation state. Mn is reduced.
I– changes from a –1 oxidation state to a +5 oxidation state. I– is oxidized.
Step 4:
Reduction half-reaction
MnO4– ® MnO2
Oxidation half-reaction
I– ® IO3–
Step 5:
For the reduction half-reaction the steps are:
a) Atoms other than O and H MnO4– ® MnO2
b) O atom with H2O
MnO4– ® MnO2 + 2H2O
c) H atom with H+
4H+ + MnO4– ® MnO2 + 2H2O
d) Charge with e–
3e– + 4H+ + MnO4– ® MnO2 + 2H2O
Repeating for the oxidation half-reaction gives: I– + 3H2O ® IO3–+ 6H+ + 6e–
Step 6.: Multiply the reduction half-reaction by 2x:
2 [3e– + 4H+ + MnO4– ® MnO2 + 2H2O]
=
6e– + 8H+ + 2MnO4– ® 2MnO2 + 4H2O
Step 7:
+
=
6e– + 8H+ + 2MnO4–
® 2MnO2 + 4H2O
–
I + 3H2O
® IO3–+ 6H+ + 6e–
6e– + 8H+ + 2MnO4– + I– + 3H2O ® 2MnO2 + 4H2O + IO3– + 6H+ + 6e–
2H+(aq) + 2MnO4–(aq) + I–(aq) ® 2MnO2(s) + H2O(l) + IO3–(aq)
(acidic solution)
Step 8: ONLY If in basic medium add OH–(aq) to both sides to neutralize any H+
2OH– + 2H+ + 2MnO4– + I– ® 2MnO2 + H2O + IO3– + 2OH–
2H2O + 2MnO4– + I– ® 2MnO2 + H2O + IO3– + 2OH–
H2O(l) + 2MnO4–(aq) + I–(aq) ® 2MnO2(s) + IO3–(aq) + 2OH–(aq)
(basic solution)
Step 9: Add back in spectator ions. Let’s do the basic media solution.
H2O(l) + 2KMnO4(aq) + KI(aq) ® 2MnO2(s) + KIO3(aq) + 2KOH(aq)
Foothill College Chemistry 1C-Larson/Daley
3
Last Modified 8/9/2017
Redox Worksheet
Lab Section: MW or TTH
Name:
Oxidation-Reduction Exercise
Assigning Oxidation Numbers
Give the oxidation number of nitrogen in each of the following:
NH2OH
N2
NH4+
N2H4
HNO3
Give the oxidation number of manganese in each of the following:
MnO4–
Mn2O3
Mn
MnSO4
Types of Reactions
Combination, decomposition, single displacement and combustion reactions are all examples of oxidation-reduction
reactions. Note that there are reactions that do not fall into these categories that are redox reactions.
Precipitation reactions that are categorized as double displacement (Also called exchange reactions in some textbooks.)
are NOT oxidation-reduction reactions. All acid-base reactions, whether they are Arrhenius, Bronsted-Lowry or
Lewis, are NOT oxidation-reduction reactions.
Directions: For each of the following reactions classify the reaction as: (a) redox, (b) precipitation, or (c) acid-base. For
the redox reactions, identify the oxidizing and reducing agent.
Reaction
Reaction Type
1.
2Al(s) + 3Cl2(g) ® 2AlCl3(s)
2.
Ag+(aq) + Cl–(aq) ® AgCl(s)
3.
Mg(s) + 2HCl(aq) ® MgCl2(aq) + H2(g)
5.
2H+(aq) + CO32– (aq) ® CO2(g) + H2O(l)
6.
HCl(aq) + NH3(aq) ® NH4Cl(aq)
7.
2CH3CH3 (g) + 7O2(g) ® 4CO2(g) + 6H2O(g)
8.
AgCl(s) + 2NH3(aq) ® Ag(NH3)2+(aq) + Cl–(aq)
Foothill College Chemistry 1C-Larson/Daley
4
Last Modified 8/9/2017
Redox Worksheet
Lab Section: MW or TTH
Name:
Balancing Oxidation-Reduction Reactions
For each of the following reactions balance the reaction using the half-reaction method.
1.
Blood alcohol levels can be determined by a redox titration with a standard potassium dichromate solution under acidic
conditions. The unbalanced equation for the reaction is:
C2H5OH(aq) + K2Cr2O7 + HCl(aq) ® CO2 (g) + CrCl3 (aq) +KCl(aq)
Write the balanced net-ionic equation for this reaction. Add back spectator ions once the equation is balanced.
2.
One of the reactions we will use in our qualitative analysis laboratory project is the oxidation of solid chromium (III)
hydroxide by hydrogen peroxide in a strongly alkaline aqueous solution. The products of the reaction are aqueous
chromate ion and water. Write the balanced net-ionic equation for this reaction.
Foothill College Chemistry 1C-Larson/Daley
5
Last Modified 8/9/2017
Redox Worksheet
Lab Section: MW or TTH
3.
Name:
In chemistry 1A at Foothill students determine the mass percent oxalate in a “green crystal” that they first synthesize.
The oxalate analysis involves dissolving a sample of the green crystal in water, adding sulfuric acid and then titrating the
solution with a standard potassium permanganate solution. The unbalanced skeleton net-ionic equation for the reaction
that takes place is:
MnO4–(aq) + C2O42–(aq) ® Mn2+(aq) + CO2(g)
Write the balanced net-ionic equation for this reaction.
4.
One of the reactions we will use in our qualitative analysis laboratory project is the disproportionation of thiosulfate ion
to sulfide and sulfate in a weakly acidic solution. A disproportionation reaction is a special type of redox reaction where
a single reactant (thiosulfate in this case) is both oxidized and reduced. The reaction is used as a source of sulfide ions to
separate insoluble sulfides from soluble sulfides. Write the balanced net-ionic equation for the disproportionation of
thiosulfate ion.
Foothill College Chemistry 1C-Larson/Daley
6
Last Modified 8/9/2017