Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
(Redox) 1. Synthesis 2. Decomposition 3. Single Replacement 4. Double Replacement * Combustion assigned to atoms or ions as a way to keep track of electron transfers The oxidation number for each atom in free elements is always zero (ex: H2, Na, S8) The oxidation numbers of ions are the same as the charge on the ion. In MgCl2 the Mg+2 ion has an oxidation number of +2. Each of the two Cl- ions has an oxidation number of –1. Group 1 metals always have a +1 oxidation number. Group 2 metals always have a +2 oxidation number. Oxygen has a –2 oxidation number in practically all of its compounds. The exceptions are in peroxides, such as H2O2 and Na2O2 where the oxidation number is –1 and in compounds with fluorine (OF2) in which it is +2. Hydrogen has an oxidation number of +1 in all its compounds except metal hydrides formed with Group 1 and Group 2 metals, which it is –1. The sum of the oxidation numbers in a compound must be zero. The sum of the oxidation numbers in a polyatomic ion must be equal to the charge on the ion. In the CO32- ion, the three oxygens provide a total of –6; the carbon must then be +4. Situation 1: Find the oxidation numbers of N and O in N2O. Situation 2: Find the oxidation numbers of K, Mn and O in KMnO4. (next slide for chart) Element K Each Total Mn O4 reactions in which both oxidation and reduction occur the loss of electrons by an atom or an ion, increasing the oxidation number Ex: Na0 Na+1(Na loss 1e-) Na Na+1 + 1e Ex: O-2 O0 (O loss 2e-) the gain of electrons by an atom or an ion, reducing the oxidation number Ex: F0 F-1 (F0 gain 1e-) F0 + 1eF-1 Ex: Ca+2 Ca0 (Ca+2 gained 2e-) LEO (Lose Electrons Oxidations) oxidation # increases oxidation # decreases GER (Gain Electrons Reduction) LEO the lion says GER! ◦ Electronic equations or equations in which only the atoms/ions being oxidized or reduced are shown. Ex: 2Al0 + 3Cl20 = 2Al+3Cl3-1 Oxidation ½: Al0 → Al+3 + 3e- Reduction ½: Cl20 + 2e- → 2Cl-1 Must conserve the number of electrons! Number of electrons lost equals number of electrons gained Oxidizing Agent: causes something else to undergo oxidation by taking on its electrons. This is the substance reduced. Reducing Agent: causes something else to undergo reduction by giving its electrons. This is the substance oxidized. ions that are not changed during a redox reaction Identifying Redox Reactions Once oxidation numbers are assigned, the atom that has shown an increase can be identified as the one that has undergone oxidation. The atom that has a decrease in oxidation number can be identified as the one that has undergone reduction. *Double Replacement reactions are never redox reactions because NOTHING CHANGES! Ex: HCl + NaOH = NaCl + H2O Table J – Reactivity Series of Selected Metals and Nonmetals Lithium will react more readily with a nonmetal than will any other metal on the list. Hydrogen (not a metal!!) is shown on the metal side to illustrate what metals will react with acids and which will not. All metals above hydrogen will react spontaneously with acids, while those below hydrogen will not. if a reaction proceeds to completion without adding energy sources once it is started, then it is spontaneous You must have something to oxidize and something to reduce. Mg(s) and Al(s) can only be oxidized Mg+2 (aq) and Al+3(aq) can only be reduced *One metal must be neutral, the other an ion!!! Cu(s) + Al(s) = no rxn both cannot be oxidized 2Cr+3 + 3Mn(s) = 3Mn+2 + 2Cr(s) Cr+3 + Ni(s) = no rxn Ni is lower than Cr on Table J, must be higher. Again, there must be something to oxidize and something to reduce: F2 and Cl2 can only be reduced F- and Cl- can only be oxidized Cl2 + Br2 = no rxn – can’t both be reduced Cl2 + 2Br- = 2Cl- + Br2 I2 + Br- = no rxn I is lower on the table than Br, must be higher HCl + Au(s) = Au will not react with H HCl + Sr(s) = SrCl2 + H2 Yes because Sr is more likely to oxidize than H2 2Na(s) + ZnCl2(aq) = NaCl + Zn Yes because Na is higher on Table J so it will oxidize while Zn reduces. Voltaic Cells: ½ cells of 2 different metals connected by wires – produces electricity. redox is spontaneous. Electrons flow from the metal being oxidized (reducing agent) to the ion being reduced (oxidizing agent). Salt Bridge: These are batteries which functions to complete the circuit by allowing ions to flow from reducing side to oxidizing side. Diagram: Which reaction is occurring in the cell? Zn(s) + Cu+2 = Zn+2 + Cu(s) or Zn+2 + Cu(s) = Zn(s) + Cu+2 The anode is Zn0 and the cathode is Cu+2 lectrochemical Cells electricity is used to produce a nonspontaneous redox reaction. Electricity is required! This is not spontaneous!! Requires the use of a Battery. The anode is connected to the and the cathode is connected to the Examples: Electroplating Diagram: *The only way to get a group one metal in its elemental form is by electrolysis of their fused salts. Example: KBr = K(s) + Br2(s)