Transcript Chapter 20
Chapter 20 Oxidation Numbers 1. 2. 3. 4. 5. 6. 7. Elements always have oxidation number = 0. Column I alkalai metals in compounds always have oxidation number = +1. Column II alkaline earth metals in compounds always have oxidation number = +2. Aluminum and gallium oxidation numbers are +3, zinc and cadmium are +2, and silver is +1. Hydrogen normally has oxidation number = +1 in compounds except when combined with Colunm I or Column II elements; then rules 2 and 3 apply. Oxygen normally has oxidation number = -2 except in H2O2 (rule 4 has priority), or Column I and Column II oxides, where rules 2 and 3 apply. To calculate other oxidation numbers pretend oxidation numbers are per atom charges and make all “charges” on all atoms total up to overall charge on ion or molecule containing atoms. Calculating Oxidation Numbers Each oxide ion has a charge of -2 7 oxide ions have a subtotal charge of -2 x 7 = -14 Since the formula has to be uncharged the 2 manganese ions have to have a +14 subtotal The +14 subtotal divided evenly over 2 manganese ions gives each manganese +14 / 2 = +7 This compound is manganese(VII) oxide Work oxidation numbers of Cr and S in Cr2(SO4)3 (Hint: treat SO42- as a single particle) Oxidation and Reduction 1. In simple chemical reduction-oxidation (redox) reactions one reactant substance contains an atom whose oxidation number increases when product is created. This substance becomes oxidized in the reaction. This substance is called a reducing agent or reductant because it causes another substance to become reduced. 2. The other reactant substance contains an atom whose oxidation number decreases when product is created. This substance becomes reduced in the reaction. This substance is called an oxidizing agent or oxidant because it causes another substance to become oxidized. Oxidation -Reduction Reactions I Play Video On YouTube Oxidation - Reduction Reactions II Play Video On YouTube Common Oxidation States Play Video On YouTube Balancing Redox Reactions An alternative to textbook method called the REDOHX method is outlined below. 1. 2. 3. 4. 5. 6. 7. Balance Reducing and oxidizing atoms first. Update total “charges” on reducing and oxidizing atoms by multiplying oxidation numbers by all appropriate subscripts and coefficients. Balance Electrons by adjusting coefficients in front of oxidizing and reducing agents. Balance atoms which Don’t fit other categories. Balance Oxygens by adding H2O. Balance H by adding H+ Xtra work only when balancing in base solution. Add OH to both sides to destroy H+ and then eliminate redundant H2O. 2+ Cu + 0 Fe Redox Reaction Play Video In YouTube CuO + C Redox Reaction Play Video On YouTube Mercury (II) Oxide Decomposition Play Video On YouTube SnCl2 + Zn0 Redox Reaction Play Video On YouTube Standard Cell Potentials A standard reduction potential (previous slide) is a stoichiometry-independent potential energy difference measured in voltage units associated with a reduction reaction done under standard conditions. An oxidation reaction is simply the reverse of a reduction reaction and therefore has the same voltage with the sign reversed as the voltage of the (opposite) reduction reaction found in a reduction potential table. The total cell potential for a redox reaction done under standard conditions is sum of the voltages (energies) for the oxidation part plus the reduction part. Remember to change the sign of voltage for oxidation reaction before adding voltage for reduction reaction. Relationship Between E° and G° To convert an an energy in voltage units into a Gibbs free energy three things need to be done: 1. 2. 3. Voltages are stoichiometry-independent (G° is not). Need to balance redox reaction to figure out how many electrons involved in redox reaction (“e”) and multiply by this number to fix this. Sign convention of voltages are opposite that of G°. Need to change sign of energy. Voltages based on coulomb quantities and G° based on mole quantities. Need to multiply by Faraday’s constant (F = 96,500 C/mol) to fix this. In Summary: G° = -eFE° Nonstandard E Values Nernst Equation: Analogous to nonstandard ∆G equation ∆G = ∆Go + RTlnQ E = Eo - (RT/eF)lnQ • Notice effect of opposite sign convention on direction of deviation from standard value • Notice RT (kJ/mol) becomes RT/eF (J/coul) • R = 0.008314 kJ/mol-K (∆G) vs. R = 8.314 J/mol-K (E) Nonstandard E Values Batteries in real world seldom have standard conc’s of all redox reactants and products. To figure out if voltage higher or lower than calculated for standard conditions use Le Châtelier’s principle to decide if reaction more or less spontaneous than standard and interpret this in terms of more or less positive E value. Problem 20.58: Al(s) + 3 Ag+(aq) Al3+(aq) + 3 Ag(s) (a) Dilute anode cell (b) Increase Al(s) (c) Increase AgNO3 vol. (d) Add HCl to AgNO3 cell Faraday’s Law To convert between number of moles of electrolysis product (n) and amps of current (I), or t (time in seconds), or coulombs of charge (C = It), or number of electrons involved in redox reaction (“e”) Faraday’s law is used: n = It/(eF) Problem 20.86: (a) What mass of Mg formed by passsing 5.25 A thru molten MgCl2 for 2.50 days? (b) How many minutes to make 10.00 g of Mg from molten MgCl2 using I = 3.50 A?