1. Introduction to Redox Mrs. Kay Chemistry 12 Chapter 18 Pages:713-729

2. Redox Reactions combustion of gas in a car rusting of metals bleaching hair reactions in batteries Cut apples turning brown What do you think Redox stands for?

3. Redox = oxidation and reduction reactions Oxidation: loss of electrons Reduction: gain of electrons Hint: Leo the lion says Ger

4. redox reactions are a family of reactions that are concerned with the transfer of electrons between species Like acid-base reactions, redox reactions are a matched set -- you don't have an oxidation reaction without a reduction reaction happening at the same time

5. Single Displacement redox reaction Cu(s) + 2AgNO 3 (aq) ---> Cu(NO 3 ) 2 (aq) + 2Ag(s) The silver nitrate solution is transparent, when the copper wire is placed in it, the reaction begins slowly. The wire is coated with silver, while the copper is broken down into ions

6. Ex: Mg(s) + O 2 (g)  MgO(s) Mg loses 2 electrons to become Mg +2 O 2 gains 2 electrons to become O -2 The total reaction = redox reaction The half reactions = oxidation and reduction reactions

7. half reactions for the overall redox reaction Mg  Mg +2 + 2e- (Magnesium loses electrons, so it is oxidized) O 2 + 4e-  2O 2- (Oxygen gains electrons, so it is reduced)

8. We’re Not Finished yet! When combining half reactions, you must make sure that the electrons gained = the electrons lost. This ensures balancing of the redox reaction. (so we multiplied the first half reaction by 2, so that the 4e- balanced out) 2Mg  2Mg +2 + 4e- O 2 + 4e-  2O 2- Total redox reaction is: 2Mg + O 2  2MgO

9. Let’s see that again… The unbalanced reaction is as follows: Look at each half reaction separately: aluminum metal being oxidized to form an aluminum ion with a +3 charge and oxygen being reduced to form two (2) oxygen ions, each with a charge of -2.

10. If we combine those two (2) half-reactions, we must make the number of electrons equal on both sides. The number 12 is a common

11. Taking care of the number of atoms, you should end up with:

12. Practice Writing the half reactions and balance the following equations. 1.Fe + Br 2  FeBr 3 2.Ni + HgCl 2  Hg + NiCl 2 3.Sn +4 + Cu  Sn +2 + Cu +2 4.CO+ I 2 O 5  CO 2 + I 2 (this one’s tricky!)

13. Vocabulary: Examine : 2Mg + O 2  2MgO Mg oxidation # increased from 0 to +2, therefore it has oxidized. O 2 oxidation # decreased from 0 to –2, therefore it has reduced.

14. Oxygen has been reduced because of magnesium, so magnesium is the reducing agent (the atom who gave its electrons away)

15. Magnesium has been oxidized because of oxygen, so the oxygen is the oxidizing agent (the atom who accepted electrons) In redox: the reducing agent is oxidized and the oxidizing agent is reduced.

16. Assigning oxidation numbers Pg 721-726 Actual or hypothetical charges assigned using a set of rules. –Used to describe redox reactions, identify redox reactions and to identify oxidizing agents and reducing agents.

17. Oxidation Numbers from Lewis structures Using electronegativities In molecules, the more electronegative atom will have the electrons found closer to its nucleus. Molecules don’t have CHARGES, but you can assign Ox # based on which atom “owns” the electrons more So the Oxygen = -2 and the Hydrogen each = +1

18. Oxidation Numbers Molecules with the same atoms will have no difference in electronegativity, so they share electrons evenly. Ex: Cl 2 occurs where Cl-Cl, then pair of electrons are shared equally, so there is NO OWNING of electrons. Cl will have an oxidation number of zero.

19. Oxidation numbers for ionic compounds The oxidation number for an ionic compound is equal to its charge Example –MgO is made of Mg 2+ and O 2- –So, Magnesium has an ox# of +2 and oxygen has an ox# of -2.

20. Oxidation Number Rules: Pg 724 RulesExample 1. A pure element has an ox # of zero. Ne, Na, Cl 2 = 0 2. A monatomic ion ox # equals its charge Ca 2+ = +2 3. Hydrogen with a metal = -1, Hydrogen with a non-metal =+1 Ex: H 2 S : H=+1 NaH: H=-1

21. Oxidation Number rules RulesExamples 4. Oxygen is usually -2, unless its in a peroxide (H 2 O 2 or OF 2 ) Li 2 O : O = -2 H 2 O 2 : O= -1 OF 2 : O=+2 5. In molecules, the more electronegative atom usually has the same charge that it would have had as an ion PCl 3 : Cl=-1, so P=+3 CS 2 : S=-2, so C=+4

22. Oxidation number rules RulesExamples 6. The sum of all ox# in a compound or molecule equal zero CF 4 : C=+4, because 4 F=-1 each 7. The sum of ox# in a polyatomic ion equal the charge of the ion NO 3 -

23. Exercise: For the following reactions 1.draw a diagram showing the loss and gain of electrons 2.identify the substance oxidized, the substance reduced, the oxidizing agent and the reducing agent. 3.Write the oxidation and reduction half reactions

24. Practice together: Example: 1,2.2Mg + O 2  2MgO 3. 2Mg  2Mg +2 + 4e- O 2 + 4e-  2O 2- 00+2-2 Gain of e- Substance reduced Oxidizing agent Loss of e- Substance oxidized Reducing agent

25. Work on the following in class, finish for homework: 1.Ni + HgCl 2  Hg + NiCl 2 2.2Na + 2H 2 O  2NaOH + H 2 3.Cl 2 +2NaBr  Br2 + 2NaCl 4.4NH 3 + 7O 2  4NO 2 + 6H 2 O

26. Practice: Page 715 # 1-4 Page 716 # 6 & 7 Page 726 # 9-12 http://staff.prairiesouth.ca/~chemistry/che m30/6_redox/redox1_2.htmhttp://staff.prairiesouth.ca/~chemistry/che m30/6_redox/redox1_2.htm