(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: H 2, Na, S 8 )  The oxidation numbers of ions are the same as the charge on the ion. In MgCl 2 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 H 2 O 2 and Na 2 O 2 where the oxidation number is –1 and in compounds with fluorine (OF 2 ) 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 CO 3 2- 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 N 2 O.

 Situation 2: Find the oxidation numbers of K, Mn and O in KMnO 4. (next slide for chart)

ElementKMnO4O4 Each Total

 reactions in which both oxidation and reduction occur

 the loss of electrons by an atom or an ion, increasing the oxidation number  Ex: Na 0 Na +1 (Na loss 1e - )  Na Na e -  Ex: O -2 O 0 (O loss 2e - )

 the gain of electrons by an atom or an ion, reducing the oxidation number  Ex: F 0 F -1 (F 0 gain 1e - )  F 0 + 1e - F -1  Ex: Ca +2 Ca 0 (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: 2Al 0 + 3Cl 2 0 = 2Al +3 Cl 3 -1

 Oxidation ½: Al 0 → Al e -  Reduction ½: Cl e - → 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 + H 2 O

 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 Mn(s) = 3Mn Cr(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:  F 2 and Cl 2 can only be reduced  F - and Cl - can only be oxidized

 Cl 2 + Br 2 = no rxn – can’t both be reduced  Cl 2 + 2Br - = 2Cl - + Br 2  I 2 + 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) = SrCl 2 + H 2  Yes because Sr is more likely to oxidize than H 2   2Na(s) + ZnCl 2 (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.  These are batteries which redox is spontaneous.  Electrons flow from the metal being oxidized (reducing agent) to the ion being reduced (oxidizing agent).  Salt Bridge: 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 Zn 0 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) + Br 2 (s)