8.1B: Oxidation States

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Rules for assigning oxidation numbers: 1. The oxidation number of a free element = 0. 2. The oxidation number of a monatomic ion = charge on the ion. 3. The oxidation number of hydrogen = + 1 and rarely - 1. 4. The oxidation number of oxygen = - 2 and in peroxides - 1. 5. The sum of the oxidation numbers in a polyatomic ion = charge on the ion. Elements in group 1, 2, and aluminum are always as indicated on the periodic table. K2CO3 The sum of all the oxidation numbers in this formula equal 0. Multiply the subscript by the oxidation number for each element. To calculate O.N. of C K = (2) ( + 1 ) = + 2 O = (3) ( - 2 ) = - 6 therefore, C = (1) ( + 4 ) = + 4 HSO4 - To calculate O.N. of S The sum of all the oxidation numbers in this formula equal -1. Multiply the subscript by the oxidation number for each element. H = (1) ( + 1 ) = + 1 O = (4) ( - 2 ) = - 8 therefore, S = (1) ( + 6 ) = + 6 Calculate O.N. in following compounds: 1. Sb in Sb2O5 2. N in Al(NO3)3 3. P in Mg3(PO4)2 4. S in (NH4)2SO4 5. Cr in CrO4 -2 6. Hg in Hg(ClO4)2 7. B in NaBO3 8. Si in MgSiF6 9. I in IO3 - 10. N in (NH4)2S 11. Mn in MnO4 - 12. Br in BrO3 - 13. Cl in ClO - 14. Cr in Cr2O7 -2 15. Se in H2SeO3 Reducing Agents and Oxidizing Agents  Reducing agent - the reactant that gives up electrons.  The reducing agent contains the element that is oxidized (looses electrons).  If a substance gives up electrons easily, it is said to be a strong reducing agent.  Oxidizing agent - the reactant that gains electrons.  The oxidizing agent contains the element that is reduced (gains electrons).  If a substance gains electrons easily, it is said to be a strong oxidizing agent. Example: Fe2O3 (cr) + 3CO(g) 2Fe(l) + 3CO2 (g)  Notice that the oxidation number of C goes from +2 on the left to +4 on the right.  The reducing agent is CO, because it contains C, which loses e - .  Notice that the oxidation number of Fe goes from +3 on the left to 0 on the right.  The oxidizing agent is Fe2O3, because it contains the Fe, which gains e - . Practice Problems: In any Redox equation, at least one particle will gain electrons and at least one particle will lose electrons. This is indicated by a change in the particle's oxidation number from one side of the equation to the other. For each reaction below, draw arrows and show electron numbers as in the example here. The top arrow indicates the element that gains electrons, reduction, and the bottom arrow indicates the element that looses electrons, oxidation. An arrow shows what one atom of each of these elements gaines or looses. 1. Mg + O2 MgO 2. Cl2 + I - Cl - + I2 3. MnO4 - + C2O4 -2 Mn+2 + CO2 4. Cr + NO2 - CrO2 - + N2O2 -2 5. BrO3 - + MnO2 Br - + MnO4 - 6. Fe+2 + MnO4 - Mn+2 + Fe+3 7. Cr + Sn+4 Cr+3 + Sn+2 8. NO3 - + S NO2 + H2SO4 9. IO4 - + II2 10. NO2 + ClO - NO3 - + Cl - Balancing Redox Equations by the Half-reaction Method 1. Decide what is reduced (oxidizing agent) and what is oxidized (reducing agent).  Do this by drawing arrows as in the practice problems. 2. Write the reduction half-reaction.  The top arrow in step #1 indicates the reduction half-reaction.  Show the electrons gained on the reactant side.  Balance with respect to atoms / ions.  To balance oxygen, add H2O to the side with the least amount of oxygen. THEN: add H + to the other side to balance hydrogen. 3. Write the oxidation half-reaction.  The bottom arrow in step #1 indicates the oxidation half-reaction.  Show the electrons lost on the product side.  Balance with respect to atoms / ions.  To balance oxygen, add H2O to the side with the least amount of oxygen. THEN: add H + to the other side to balance hydrogen. 4. The number of electrons gained must equal the number of electrons lost.  Find the least common multiple of the electrons gained and lost.  In each half-reaction, multiply the electron coefficient by a number to reach the common multiple.  Multiply all of the coefficients in the half-reaction by this same number. 5. Add the two half-reactions.  Write one equation with all the reactants from the half-reactions on the left and all the products on the right.  The order in which you write the particles in the combined equation does not matter. 6. Simplify the equation.  Cancel things that are found on both sides of the equation as you did in net ionic equations. Rewrite the final balanced equation. Check to see that electrons, elements, and total charge are balanced.  There should be no electrons in the equation at this time.  The number of each element should be the same on both sides. It doesn't matter what the charge is as long as it is the same on both sides. Practice Problems: 1. Identify the oxidizing agent and reducing agent in each equation:  H2SO4 + 8HI H2S + 4I2 + 4H2O  CaC2 + 2H2O C2H2 + Ca(OH)2  Au2S3 + 3H2 2Au + 3H2S  Zn + 2HCl H2 + ZnCl2 2. To make working with redox equations easier, we will omit all physical state symbols. However, remember that they should be there. An unbalanced redox equation looks like this: MnO4 - + H2SO3 + H + Mn+2 + HSO4 - + H2O Study how this equation is balanced using the half-reaction method. It is important that you understand what happens in each step. Be prepared to ask questions about this process in class tomorrow.