7. The sum of oxidation numbers in a neutral compound is zero and in a polyatomic ion the sum is equal to its net charge.
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1 General Information: Rules for Determining Oxidation Numbers 1. Elements in molecules consisting of just that element are assigned an oxidation number of zero Na, Fe, Cl 2, N 2, P 4, S 8 2. Monatomic ions and ionic compounds have an oxidation number equal to the individual ion charge Cl -, S -2, Al +3, MnCl 2, PbBr 2, TiCl 4 3. In compounds: Group IA metals are assigned an oxidation number of +1 Group IIA metals are assigned an oxidation number of +2 Group IIIA metals are assigned an oxidation number of In a compound, fluorine is assigned an oxidation number of In a compound, hydrogen is assigned an oxidation number of +1 Exception: Metal hydrides (Rule 3 takes priority) 6. In a compound, oxygen is assigned and oxidation number of -2 Exception: Peroxides- H 2 O 2, Na 2 O 2, BaO 2 (O = -1) Superoxides KO 2, RbO 2, CsO 2 (O = -1/2) Others OF 2 (O = +2) 7. The sum of oxidation numbers in a neutral compound is zero and in a polyatomic ion the sum is equal to its net charge. 8. Apply the oxidation number rules in the order given. Activity 1: Determine the oxidation number of the underlined element in the following compounds. Oxidation Number Oxidation Number Oxidation Number HClO NaClO H 2 SO 3 NH 4 + SO 4 HNO 3 CO 3 S 6 Ca 2+ CO MnO 4 H 3 PO 4 AlO 2 HClO P 4 PbO 2 LiH Fe(OH) 3 UF 6 K 2 Cr 2 O 7 Na 4 P 2 O 7 Ag ZnO 2 UO 2 2+ HNO 2 Fe(OH) 2 HAsO 4 Na 2 S 2 O 3 O 3 BaH 2 N 2 VO 2 + S 2 O 3 Na 2 O HBr NO 2 1
2 BeO Al 3+ Na + P 2 O 3 Cl 2 Sb 2 O 5 General Information: Balancing Redox Reactions via the Oxidation Number[Arrow] Method 1. Assign oxidation numbers for every species in the reaction. 2. Draw arrows connecting the species that have been oxidized and reduced. 3. Determine the electrons gained and lost. 4. Set the gained electrons equal to the lost electrons (you will have to multiply). 5. The multiplier becomes the new coefficient for everything associated with the arrow. See below for Acidic or Basic conditions. 6. Re-write the reaction with the new coefficients. 7. Check the mass and charge. Acidic and Basic Conditions: a. Acidic Conditions: Add H 2 O to balance oxygen atoms then add H + to balance hydrogen atoms. b. Basic Conditions: Add OH to balance charge then add H 2 O to balance H atoms (or O atoms) Activity 2: Balance the following redox reactions using the oxidation state/ arrow method. Fe 3+ + Tl + Fe 2+ + Tl 3+ Al + Cd +2 Al 3+ + Cd MnO 4 + Fe +2 Mn 2+ + Fe 3+ [Acidic solution] 2
3 General Information: Balancing Simple, Acidic & Basic Redox Reactions via the HalfReaction Method Step 1: Split the equation up into a reduction half reaction and an oxidation half reaction. Ignoring hydrogen and oxygen, be sure to match pairs with a common element, in other words, don't create elements out of thin air. Step 2: Balance all elements that are not H or O. Step 3: Balance O atoms by adding H 2 O's to one side of the equation. Step 4: Balance H atoms by adding H + ions to one side of the equation. Step 5: Balance the electric charge by adding e to the more positive side. Note that you are not trying to get both sides equal to zero, just equal to each other. At this point, the reduction reaction has e as a reactant and the oxidation reaction has e as a product. Step 6: Make sure the number of e that were added to the total reduction reaction is the same as the number of e that were added to the total oxidation reaction. If not, multiply either one or both total reactions by an integer that will make them equal. Step 7: Combine the two half reactions and cancel species that appear on both sides and reduce the coefficients to the smallest whole numbers. No e should appear in the final equation. Activity 3 Fe 3+ + Tl + Fe 2+ + Tl 3+ Al + Cd +2 Al 3+ + Cd MnO 4 + Fe +2 Mn 2+ + Fe 3+ [Acidic solution] 3
4 Activity 4: Identify the oxidizing and the reducing agents for the following reactions Reaction Oxidizing Agent Reducing Agent 1. FeS 2 + O 2 Fe(OH) 3 + SO 4 2. Al + NO 3 3. Cr 2 O 7 4. MnO 4 AlO 2 + NH 3 + Fe 2+ Cr 3+ + Fe 3+ + Fe +2 Mn 2+ + Fe Na + H 2 O Na + + H 2 6. PbO 2 + Cr 3+ Pb 2+ + CrO 4 7. Fe 3+ + Tl + Fe 2+ + Tl Al + Cd +2 Al 3+ + Cd Activity 5 Balance the following reaction via the half-reaction method and present your results to the whole group. Cr 2 O 7 + Fe 2+ Cr 3+ + Fe 3+ [acidic] Now balance the same reaction via the oxidation state/ arrow method and present your results to the whole group. Cr 2 O 7 + Fe 2+ Cr 3+ + Fe 3+ [acidic] As a group discuss the advantages and disadvantages of each balancing method as you see them. What strategy will you employ on the exam when asked to balance redox reactions? 4
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