1. Oxidation & Reduction IB Chemistry

2. Oxidation Numbers Rules for Assigning Oxidation Numbers 1Oxidation numbers always refer to single ions 2The oxidation number of an atom is always 0 3The sum of the oxidation numbers in a neutral compound is 0 4The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion 5The oxidation number of Hydrogen is usually +1 (-1 when bonded to a metal) 6The oxidation number of Oxygen is always -2 (except for H 2 O 2 ) 7The oxidation numbers of alkali metals are +1 8The oxidation numbers of alkaline earth metals are +2

3. Examples NO 2 N 2 O 5 HClO 3 Ca(NO 3 ) 2 Fe(OH) 3 CO 3 -2 SO 4 -2 O= -2 Ø = N +(-2x2)  N= +4 O= -2 Ø = 2N +(-2x5)  N= +5 O= -2, H= +1 Ø = Cl +(-2x3) +1  Cl = +5 O= -2, Ca= +2 Ø = 2N+(-2x6)+2  N=+5 O= -2, H= +1 Ø = Fe+(-2x3)+(1x3)  Fe=+3 O= -2 -2 = C+(-2x3)  C =+4 O= -2 -2 = S+(-2x4)  S =+6

4. Determining What’s Happened… Careful examination of the oxidation numbers of atoms in an equation allows us to determine what is oxidized and what is reduced in an oxidation-reduction reaction

5. Example

6. Exercise For each of the following reactions find the element oxidized and the element reduced Cl 2 + KBr  KCl + Br 2 Cu + HNO 3  Cu(NO 3 ) 2 + NO 2 + H 2 O HNO 3 + I 2  HIO 3 + NO 2

7. Exercise For each of the following reactions find the element oxidized and the element reduced Cl 2 + KBr  KCl + Br 2 0 +1-1 +1-1 0 Br increases from –1 to 0 -- oxidized Cl decreases from 0 to –1 -- Reduced K remains unchanged at +1

8. Exercise For each of the following reactions find the element oxidized and the element reduced Cu + HNO 3  Cu(NO 3 ) 2 + NO 2 + H 2 O 0 +1+5-2 +2 +5 -2 +4 –2 +1-2 Cu increases from 0 to +2. It is oxidized Only part of the N in nitric acid changes from +5 to +4. It is reduced The nitrogen that ends up in copper nitrate remains unchanged

9. Exercise For each of the following reactions find the element oxidized and the element reduced HNO 3 + I 2  HIO 3 + NO 2 1 +5 -2 0 +1+5-2 +4-2 N is reduced from +5 to +4. It is reduced I is increased from 0 to +5 It is oxidized The hydrogen and oxygen remain unchanged.

10. Oxidation-Reduction Reactions All oxidation reduction reactions have one element oxidized and one element reduced Occasionally the same element may undergo both oxidation and reduction. This is known as an auto-oxidation reduction

11. Agents The oxidizing agent takes the electron(s) and is itself reduced The reducing agent loses the electron(s) and is itself oxidized.

12. Monday 4/15/13 Objective: SWBAT review redox concepts HW: no hw Warm up: Determine oxidation #’s, which is oxidized/reduced, as well as agents

13. Half Reactions Show what is happening to the oxidized species or the reduced species (tells ½ the story) +2 -20 0 Zinc is oxidized & is the reducing agent Copper is reduced & is the oxidizing agent Sulfate acts as a spectator ion – doesn’t do anything

14. Movement of e- Write a half reaction for the reduced species Write a half reaction for the oxidized species Or…

15. Ionic Equations Add the half reactions together The e- must cancel out

16. What if e-’s don’t cancel? Half equation 2 needs to be multiplied by 2 to achieve equal amounts of e- on both sides Now they can be added together:

17. Activity Series – reducing agents Activity series allow you to predict whether a redox reaction will happen or not. – Elements higher on the chart will displace a metal ion of an element lower on the chart. – Mg + Zn 2+ will react to form Zn and Mg 2+ – More reactive elements are stronger reducing agents

18. Redox reactions in Acidified Solutions For each half equation: 1.If a metal is present, add coefficients to balance 2.Add water to balance oxygens 3.Add hydrogen ions to balance H 4.Add electrons to balance the charge 5.Balance half equations 6.Add half equations 7.Simplify (now an algebraic expression)

19. Wednesday 3/26/14 Objective: SWBAT determine the flow of electrons in a voltaic cell compared to an electrolytic cell HW: Finish packet Warm up: Determine the oxidation states of each element HNO 3 + I 2  HIO 3 + NO 2 Quiz Monday

20. Voltaic Cells Two half cells connected together – allows for electrons to be transferred during the redox reaction Produces energy in the form of electricity Half cells: a metal in contact with an aqueous solution of its own ions Spontaneous redox reaction taking place Zn is higher in the activity series Electrons will flow towards Cu Zn is higher in the activity series Electrons will flow towards Cu

21. Voltaic Cells Oxidation occurs at the anode Reduction occurs at the cathode Electrons flow towards the cathode Electron movement

22. Voltaic Cells: Purpose of Salt Bridge Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would not be able to continue.

23. Voltaic Cells: Ions, not electrons, move through salt bridge Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. – Cations move toward the cathode. – Anions move toward the anode.

24. How can you remember? a RED CAT and AN OX REDuction = CAThode ANode = OXidation

25. Electrolysis Used to make non-spontaneous redox reactions occur – Provide energy in the form of electricity from an external source Electricity is passed through an electrolyte – Electrical energy is converted to chemical energy Reverse of Voltaic cells Electrolytes: conduct energy in solution

26. Anodes & Cathodes Voltaic CellsElectrolytic Cells Anode Oxidation occurs here - + Cathode Reduction occurs here + - Electrolytic cells are “pumping” electrons – not spontaneous Voltaic cells are spontaneously occurring

27. Electrolysis of Molten NaCl Goal – to produce Na Also produces chlorine gas Na reduced, Cl oxidized

28. Electroplating The use of electrolysis to coat one metal with a thin layer of another metal Item to be plated placed at cathode Immersed in solution of metal ions of the plating material Metal using to plate with connected to the anode Electrons from the metal to the object to be plated When the metal loses an electron, its cation goes into solution The electron moves towards the object to be plated When the cation finds the electron, it turns back to its solid form – this time on the object to be plated