Chapter 18—Oxidation-Reduction Reactions：18章氧化还原反应.doc

Chapter 18Oxidation-Reduction ReactionsPriceI. Section 18.1Oxidation Reduction ReactionsA. A very common type of reaction where electrons are transferred from one atom (or group of atoms) to another.B. When a metal and a nonmetal react the metal gives some electrons to the nonmetal and they form an ionic compoundC. The metal is “oxidized”D. The nonmetal is “reduced”E. Oxidation=a loss of electronsF. Reduction=a gain of electronsII. Section 18.2Oxidation StatesA. Electron Transfer is more difficult to see when the redox reaction involves 2 or more non metalsB. To make this easier we give “imaginary” charges to the atoms in covalent compoundsC. Oxygen is a part of all redox reactions involving nonmetalsD. Oxygen = -2 (the peroxide exception!)E. Hydrogen = +1F. Oxidation states depend on the electronegativity of the atoms involved1. The more electronegative element always controls the electrons2. The most electronegative elements are FONClIII. Section 18.3Redox Reactions Between NonmetalsA. Lets examine the combustion of methane (CH4)CH4 + 2O2 CO2 + 2H2OCH4 + 2O2 CO2 + 2H2OB. Pure elements have an oxidation number of 0 (zero)C. We know oxygen is always a -2 and hydrogen is always a +1D. We use these facts to deduce the oxidation numbers of the rest of the elements in the reactionCH4 + 2O2 CO2 + 2H2O C=-4 H=+1 oxygen=0 C=+4 oxygen=-2 H=+1 O=-2E. Notice that the oxidation number of the carbon increases from -4 to +4carbon appears to have lost 8 electrons. The methane has been oxidizedF. Notice that the oxidation number of the oxygen has decreased from 0 to -2. The oxygen has been reducedG. Oxidation=an increase in oxidation number H. Reduction=a reduction in oxidation numberRules for Assigning Oxidation Numbers1. The oxidation number for all uncombined elements is zero.2. The oxidation number of a single ion is the same as its charge3. Hydrogen has an oxidation number of +14. Oxygen has an oxidation number of -2 (In peroxides oxygen is a -1)5. In binary compounds without hydrogen or oxygen the atom with the greater ectronegativity is assigned the negative oxidation number based on its charge6. The overall charge of a compound MUST BE ZERO7. For polyatomic ions, the sum of the oxidation numbers must equal the overall chargeI. “Oxidizing Agents” and “Reducing Agents”2Na + Cl2 2NaCl1. Its easy to see that the sodium is oxidized and the chlorine is reduced in this simple metal/nonmetal redox reaction2. Sodium loses electrons, chlorine gains electrons3. We say that sodium is the “reducing agent” and that chlorine is the “oxidizing agent”IV. Section 18.4-Balancing Redox ReactionsA. Until now we have been able to balance redox reactions just like we would balance any other type of reactionB. Unfortunately, some redox reactions (those involving ions that take place in aqueous solutions) are so complex that we separate the reaction into two half-reactionsC. Half-reactions are equations that have electrons as reactants or products.EMAMPLE: Ce4+ + Sn2+ Ce3+ + Sn4+This reaction can be separated into a half-reaction involving the substance being reduced e- + Ce4+ Ce3+and a half-reaction involving the substance being oxidized Sn2+ Sn4+ + 2e-The key to balancing these half-reactions is that the number of electrons lost must equal the number of electrons gainedIt is easy to see that we need to double everything in the first half-reaction 2e- + 2Ce4+ 2Ce3+Now add the second half-reaction Sn2+ Sn4+ + 2e-2e- + 2Ce4+ + Sn2+ 2Ce3+ + Sn4+ + 2e-Finally, cancel out the 2 electrons on each side2e- + 2Ce4+ + Sn2+ 2Ce3+ + Sn4+ + 2e-For the final, balanced reaction2Ce4+ + Sn2+ 2Ce3+ + Sn4+ D. It turns out that most redox reactions happen in solutions that are acidic. This makes things a bit more complicatedRules for Balancing Reactions by the Half-Reaction Method1. Identify and write the equations for the oxidation and reduction half-reactions2. For each half-reactiona. balance all of the elements except hydrogen and oxygenb. Balance oxygen using H2Oc. Balance Hydrogen using H+ ionsd. Balance the charge using electrons 3. If necessary, multiply one or both half-reactions to equalize the number of electrons transferred4. Add the half-reactions and cancel out anything that appears on both sides5. Check to be sure the elements and charges balanceV. Section 18.5Electrochemistry: An IntroductionA. Definition: Electrochemistry is the study of the interchange of chemical and electrical energy.B. Electrochemistry involves 2 types of processes1. The production of an electrical current from an oxidation-reduction reaction2. The use of an electrical current to produce a chemical changeC. Question: How do we harness the electron transfer of redox reactions in order to do work for us?D. Answer: We have to separate the oxidizing agent (the electron acceptor) from the reducing agent (the electron donator) and force the electrons to travel through a wire in order to complete the transfer.E. Consider the following typical electrochemical cellHowever, when we construct the above apparatus no flow of electrons occurs. Why? The problem is that if the electrons flowed from the right to the left the left compartment would build up a negative charge and the right would build up a net positive chargethis is no good! We must connect the two solutions so that ions can be transferred in order to cancel out the charge. To do this we need a salt bridge or a semi-permeable membrane that will allow ions to move but will not let the solutions to mix extensively. This type of a battery is called a galvanic cell.F. Definitions:1. Electrode: Another name for the terminals of a battery (one negative, one positive)2. anode: The electrode where oxidation occurs3. cathode: The electrode where reduction occurs4. electrolysis: using electrical energy to produce a chemical changeVI. Section 18.6BatteriesA. 2 important types of batteries have been around for a long time:1. Lead storage battery (like the type in your car)2.Dry cell batteries (like Energizers and Duracells)Both of these types of reactions are powered by complex redox reactions.VII. Section 18.7 CorrosionA. Corrosion (what we call rust) is simply a redox reaction where the metal is slowly reacti