1. Oxidation Numbers and Redox Reactions Section 7.2

2. Oxidation Numbers Oxidation number: the number of electrons that must be added to or removed from an atom in a combined state to convert the atom into the elemental form AKA oxidation state Do not have an exact physical meaning

4. Usage Useful for naming compounds, writing formulas, and balancing chemical equations There are specific rules for assigning oxidation numbers to a species Shared electrons are assumed to belong to the more electronegative atom in each bond

5. Rules 1.Atoms in a pure element have an oxidation number of 0 2.F has an oxidation number of -1 3.O has an oxidation number of -2, except in peroxides (-1) or bonded with F (+2) 4.H is +1 when in a compound where the other element is more electronegative than it is

6. Rules Continued 5.H is -1 when it is bonded with metals 6.The more-electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion, the other gets the positive charge it would get as a cation

7. Rules Continued 7.The algebraic sum of the oxidation numbers of all atoms in a neutral compound is 0 8.The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion 9.A monoatomic ion has an oxidation number equal to its charge

8. Practice Assign oxidation numbers to each atom: UF 6 We know F has an oxidation number of - 1, so the oxidation number for F 6 is -6 Since the compound is neutral, the U must be +6 See the board for how to write this

9. Practice Assign oxidation numbers to each atom: H 2 SO 4 H is bonded with more electronegative atoms, so it is +1 for each (total +2) O is -2 for each for a total of -8 S must be +6 to balance the neutral formula

10. Practice Assign oxidation numbers to each atom: ClO 3 - O is -2 for a total of -6 The charge on the ion is -1, so Cl must be +5 Practice problems on page 234 in the book at school.

11. Redox Reaction Reduction and oxidation occurs This is a reaction in which electrons are transferred from one atom to another Oxidation: loss of electrons Also the increase in oxidation number Reduction: gain of electrons Also the decrease in oxidation number LEO goes GER or OILRIG

12. Example of Redox Reaction Zn (s) + I 2(aq) → Zn 2+ (aq) + 2 I - (aq) Half reactions are used to show redox Zn (s) → Zn 2+ (aq) + 2e- (oxidized) I 2 (aq) + 2e- → 2 I - (aq) (reduced)

13. Oxidation Number in Redox Reactions 2KBr (aq) + Cl 2(aq)  2KCl (aq) + Br 2(aq) Assign oxidation numbers to all elements in the balanced equation K in KBr = +1, K in KCl = +1 (no change) Br in KBr = -1, Br in Br 2 = 0 (increase) Cl in Cl 2 = 0, Cl in KCl = -1 (decrease) Bromine is oxidized, chlorine is reduced

14. Reducing and Oxidizing Agents The species that is reduced is called the oxidizing agent The species that is oxidized is called the reducing agent. 2Zn + O 2  2ZnO Zn in 2Zn = 0, Zn in ZnO = +2, Zn is oxidized and is the reducing agent

15. Continued O in O 2 = 0, O in ZnO = -2, O is reduced and is the oxidizing agent Practice: identified which species is reduced, oxidized, the oxidizing agent, and the reducing agent N 2(g) + 3H 2(g)  2NH 3(g) N is reduced, oxidizing agent, H is oxidized, reducing agent

16. Disproportionation Disproportionation: a process by which a substance acts as both the reducing agent and the oxidizing agent 2 Cu + (aq) → Cu 2+ (aq) + Cu (s) +1 +2 0