1. OXIDATION-REDUCTION REACTIONS Settle in, this is going to take a while…

2. What is redox?  Reaction where there is a transfer of electrons between reactants  Oxidation involves the loss of electrons (OIL)  Oxidation number/state of the element increases  Oxidized element is the reducing agent  Reduction involves the gain of electrons (RIG)  Oxidation number/state of the element decreases  Reduced element is the oxidizing agent

3. Example Complete Reaction: Mg + Zn(NO 3 ) 2  Mg(NO 3 ) 2 + Zn Net-ionic Reaction: Mg + Zn 2+  Mg 2+ + Zn The magnesium metal was oxidized by the zinc and the zinc was reduced by the magnesium.

4. Do what?!?!  The oxidation state of the magnesium changed from 0 to +2  Oxidation state increased = oxidation  Because magnesium gave its electrons away, it is the reducing agent  The oxidation state of zinc changed from +2 to 0  Oxidation state decreased = reduction  Because zinc took the electrons, it is the oxidizing agent

5. How do you know oxidation states?  The oxidation number for any pure element is zero.  Group 1 metals form +1 ions, group 2 metals form +2 ions, group 13 metals form +3 ions.  Transition metals can be all kinds of oxidation numbers (ranging from +1 to +7)  Transition metal oxidation states can be determined based on the nonmetal(s) it’s bonded to…

6. Nonmetal oxidation states  Fluoride is ALWAYS -1, the other halides are usually -1.  Oxide is usually -2, except when it’s in the peroxide ion (-1) or bonded to fluorine (+2)  Hydrogen is +1, unless it is the hydride ion (-1)

7. Putting it all together  The total charge on a compound is zero, so all oxidation numbers must cancel out.  The total charge of elements in a polyatomic ion must add to the charge on the ion

8. Practice What is the oxidation number of each element in the following compounds? 1. Zn(NO 3 ) 2 2. H 2 SO 4 3. KMnO 4 4. N 2 O 4 5. PCl 3

9. What’s the point?  When an element gains electrons, another element must accept those electrons (Newton’s 3 rd law).  If you separate the reaction into half-reactions, you can exploit this electron transfer to generate electricity.  The study of this is electrochemistry, but more on that later…

10. Half-Reactions?  You can separate a redox reaction into the reduction reaction and the oxidation reaction.  First you have to identify which element is oxidized and which is reduced.  So let’s practice identification first:

11. Practice Determine the oxidation states of all elements in the following reactions and then identify which element is oxidized and which is reduced. N 2 + 3H 2  2NH 3 2MnO 2 + Zn + 2H 2 O  2MnO(OH) + Zn(OH) 2 AgNO 3 + Cu  Cu(NO 3 ) 2 + Ag

12. N 2 + 3H 2  2NH 3

13. 2MnO 2 + Zn + 2H 2 O  2MnO(OH) + Zn(OH) 2

14. AgNO 3 + Cu  Cu(NO 3 ) 2 + Ag

15. Separating reactions  Once the oxidized and reduced elements have been identified, separate the reactions.  Use net ionic reactions instead of complete reactions 2AgNO 3 + Cu  Cu(NO 3 ) 2 + 2Ag 2Ag +1 + Cu  Cu 2+ + 2Ag

16.  The silver is reduced, so that is the reduction reaction: 2Ag +1  2Ag  The masses are balanced, but the charges are not, so add the electrons being transferred: 2Ag +1 + 2e -  2Ag  Notice that the reduction half reaction has electrons as reactants

17. 2Ag +1 + Cu  Cu 2+ + 2Ag  The copper is oxidized, so that is the oxidation reaction: Cu  Cu 2+  The masses are balanced, but the charges are not, so add the electrons being transferred: Cu  Cu 2+ + 2e -  Notice that the oxidation half reaction has electrons as products

18. 2Ag +1 + Cu  Cu 2+ + 2Ag  When you put the 2 half-reactions together they equal the complete reaction: 2Ag +1 + 2e -  2Ag Cu  Cu 2+ + 2e -  the electrons cancel out to make the net ionic reaction

19. More about balancing  Sometimes you need more than electrons to balance a half reaction.  Oxygens and hydrogens can be balanced by the addition of H +, OH -, and H 2 O depending on if the solution is acidic or basic.

20. Balancing Acidic Redox  Balance all elements that are not oxygen or hydrogen  Balance O by adding H 2 O where needed  Balance H by adding H + where needed  Balance charges by adding e - to the more positive side

21. Balancing Basic Redox  Follow the steps for acidic balancing first, then…  Count the H + used to balance the reaction and add the same number of OH - to both sides.  Combine OH - and H + to make water.  Combine/cancel water molecules as needed.

22. A LOT of practice Balance the following equation in both an acidic and a basic solution: HOCl + NO(g)  Cl 2 (g) + NO 3 -

23. More practice Balance this reaction in an alkaline environment: MnO 2 (s) + Zn(s) + H 2 O  2MnO(OH)(s) + Zn(OH) 2 (s)

24. Homework Problems Pick a few homework problems to work through together