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IB topic 9 Oxidation-reduction

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1 IB topic 9 Oxidation-reduction
Define oxidation and reduction in terms of electron loss and gain.

2 Taking notes Do not copy, write in your own words or draw diagrams or pictures 10 % better performance if you copy 35% better performance if you write in your own words

3 2Mg + O2  2MgO Reduction-charge goes down OILRIG Redox always occurs together Question 1

4 A:  

5 B

6 C

7 D

8 9.2Redox equations Deduce simple oxidation and reduction half-equations given the species involved in a redox reaction. 2Fe + 3Cl2 2FeCl3

9 Deduce the oxidation number of an element in a compound.
+ means loss, – gain of e- Rules Elements Na, O2, S8 = 0 Group 1 = +1 H=+1 O = -2 halides -1 Many exceptions Ox. # add up to the charge on the species In covalent compounds more electronegative is – ie NH3, CCl4

10 Give Ox. numbers to each element
H2SO4, SO32- NH4+, Fe2O3, K2Cr2O7, CuCl2, Question 2 Question 3

11 State the names of compounds using oxidation numbers.
MnO2, FeO, CuCl, Na2O Manganese (IV) oxide, iron (II) oxide, Copper (I) chloride, sodium oxide [Cu(H2O)6]2+ [CuCl4]2- Hexaaquacopper(II) ion Tetrachlorocopper (II) ion

12 Deduce whether an element undergoes oxidation or reduction in reactions using oxidation numbers.
Ca + Sn2+  Ca2+ + Sn 4NH3 + 5O2 4NO + 6H2O Disproportionation Cl2 + H2OHCl + HClO Question 4

13 Deduce redox equations using half-equations.
Steps Assign O numbers and write half reactions Balance atoms other than H and O Balance O by adding H2O as needed Balance H by adding H+ as needed Balance Charges by adding e- to the + side Equalize the e- by multiplying Add the half reactions together

14 Try NO3- + Cu  NO + Cu2+ +5, -2, 0 on left +2,-2,+2 on right Cu  Cu2+ ox NO3- NO red 4H+ + NO3-  NO H2O Cu  Cu2+ + 2e- 4H+ + NO3- + 3e- NO H2O 8H+ + 2NO3- + 6e- + 3Cu  NO + 4H2O + 3Cu2+ + 6e-

15 Fe+2 + MnO4-  Fe+3 + Mn+2

16 SO32- + Cr2O72- SO42- + Cr3+

17 Internet example Question 5

18 Define the terms oxidizing agent and reducing agent.
A substance that gets reduced causes oxidation so it is an oxidizing agent

19 Identify the oxidizing and reducing agents in redox equations.
Fe2O3 + 3C  2Fe + 3CO2 Fe oxidizing C reducing IO3- + 5I- + 6H+  3I2 +3H2O IO3- oxidizing I- reducing question 6

20 9.3Reactivity Deduce a reactivity series based on the chemical behavior of a group of oxidizing and reducing agents. More reactive metals lose their e- more readily becoming a strong reducing agent Zn + CuSO4 Stronger Mg, AL, Zn, Fe, Pb, Cu, Ag simulation

21 Non metals F2 strongest oxidizing agent, most readily becomes reduced Cl2, Br2, I2
Question 7

22 Deduce the feasibility of a redox reaction from a given reactivity series.
Yes or no ZnCl2 + Ag 2FeCl3 + 3 Mg Cl2 + 2KI question 8

23 9:4 Voltaic Cells (battery)
Explain how a redox reaction is used to produce electricity in a voltaic cell Zn(s) → Zn2+(aq) + 2e- red. Agent Other half cell Cu2+ + 2e- → Cu(s) reduced This combination is a voltaic cell

24 State that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode) Which is the anode (where e- leave) /cathode? Draw this setup. Animation Draw Zn/Zn2+ and Ag/Ag+ and give the potential, show the flow of e- Where is oxidation and reduction

25 Connect these half cells with a salt bridge
this is a spontaneous reaction Animation Question 9

26 9:5 Electrolytic cells Describe, using a diagram, the essential components of an electrolytic cell. Opposite of a voltaic cell Requires electrical energy │‌ means + then – in diagrams Animation

27 The power source pushes e- to the – electrode
-electrode attracts + ions -electrode is the cathode cations gain e- so are reduced Show the electrolysis of MgF2

28 Describe how current is conducted in an electrolytic cell
Do questions 10

29 Deduce the products of the electrolysis of a molten salt
Diagram the electrolysis of molten(melted) NaCl Tell where oxidation and reduction occurs Do question 11

30 19.1 Standard electrode potentials
Describe the standard hydrogen electrode.

31 Standard H cell Conditions - Pt electrode H2 gas at 1 atm pressure
1 mol dm-3 H+ 298 K or 25oC 0.00 V Attach a half cell if e- flows to H2 it is – Like Zn which is V

32 conventions Zn(s)/Zn+2││H+(aq)/1/2 H2(g) (Pt) Oxidation on left side
More – value of electrode potentials give off e-

33 Define the term standard electrode potential (E Ö ) .
relative electrode potential compared under standard conditions with the standard hydrogen electrode Look at your data booklet

34 19.1.3 Calculate cell potentials using standard electrode potentials.
Try Cr2O72- and Br2 Answer 0.26 V

35 Predict whether a reaction will be spontaneous using standard electrode potential values.
Can a solution of tin II ions reduce a solution of iron III ions? Yes 0.91V

36 Can a solution of Sn4+ ions reduce Fe3+ to Fe
Sn e- → Sn2+ Eo = +1.33 No what does work Do question 12

37 19.2 Electrolysis Predict and explain the products of electrolysis of aqueous solutions. For water need DC in a dilute solution of H2SO4 H+ to H2 given off at the – electrode OH- to O2 at the + electrode 2H2O  4H+ + O2 + 4 e-

38 Electrolysis of NaCl(aq)
- electrode H2 + electrode dilute OH- to O2 conc Cl- to Cl2 Write half reactions Do question 13

39 Determine the relative amounts of the products formed during electrolysis.
Position in the electrochemical series + ions lower in the series will gain e- at the – electrode (cathode) in preference to those higher Hydroxide ions release e- to form oxygen and H2O in preference to other anions at the positive electrode

40 In some cases concentrations ( more concentrated may be discharged)
Nature of the electrode C and Pt are inert List all the cations and anions Cations lower in the series gain e- more readily

41 Describe the use of electrolysis in electroplating.
CuSO4(aq) with copper electrodes Cu2+ goes to – electrode and plates Cu + electrode Cu goes to Cu2+(use impure ore)

42 Electroplating Object to be electroplated is put at the negative electrode and is placed in a solution of ions of the metal used to plate it.

43 Factors affecting relative amounts
Charge on the ions Na+ Cu2+ Al3+ Al takes more energy to make Quantity of e- (amperage and time) charge = current x time Do question 14

44 Do questions 1-14 on chapter 10 in your IB Study Guide and turn in

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