Redox and Electrochemistry

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Dwayne Dorsey
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1 Redox and Electrochemistry 1

2 Electrochemistry in Action! 2

3 Rules for Assigning Oxidation Numbers The oxidation number of any uncombined element is 0. The oxidation number of a monatomic ion equals the charge on the ion. The more-electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion. The oxidation number of fluorine in a compound is always 1. 3

4 Oxygen has an oxidation number of 2 unless it is combined with F, when it is +2, or it is in a peroxide, such as H 2 O 2, when it is 1. The oxidation state of hydrogen in most of its compounds is +1 unless it is combined with a metal, in which case it is 1 In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation numbers of +1, +2 and +3 respectively. The sum of the oxidation numbers of all atoms in a neutral compound is 0. The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge of the ion. 4

5 The oxidation number of any uncombined element is 0 What is the oxidation number of H in H 2? What is the oxidation number of Ne in Ne? What is the oxidation number of O in O 2? What is the oxidation number of S in S 8? Zero 5

6 The oxidation number of a monatomic ion equals the charge on the ion What is the oxidation number of Cl in Cl-? What is the oxidation number of Na in Na+? What is the oxidation number of Al in Al +3? What is the oxidation number of Ca in Ca +2? -1, +1, +3, +2 6

7 The more-electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion. What is the oxidation number of Cl in HCl? What is the oxidation number of S in H 2 S? -1, -2 7

8 The oxidation number of fluorine in a compound is always 1. What is the oxidation number of F in CF 4? What is the oxidation number of F in FCl? What is the oxidation number of F in CH 2 F 2? -1 8

9 Oxygen has an oxidation number of 2 unless it is combined with F, when it is +2, or it is in a peroxide, such as H 2 O 2, when it is 1. What is the oxidation number of O in H 2 O? What is the oxidation number of O in H2O2? What is the oxidation number of O in CaO? What is the oxidation number of O in Al 2 O 3? What is the oxidation number of O in OF 2? -2, -1, -2, -2, +2 9

10 The oxidation state of hydrogen in most of its compounds is +1 unless it is combined with a metal, in which case it is 1 What is the oxidation number of H in H 2 O? What is the oxidation number of H in HCl? What is the oxidation number of H in H 3 P? What is the oxidation number of H in CaH 2? +1, +1, +1, -1 10

11 In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation numbers of +1, +2 and +3 respectively. What is the oxidation number of Li in Li 2 O? What is the oxidation number of Ca in CaO? What is the oxidation number of Al in Al 2 O 3? +1, +2, +3 11

12 The sum of the oxidation numbers of all atoms in a neutral compound is 0. What is the sum of the oxidation numbers in H 2 O? What is the sum of the oxidation numbers H 2 SO 4? What is the sum of the oxidation numbers in CaH 2? Zero 12

13 The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge of the ion. What is the sum of the oxidation numbers in PO 4-3? What is the sum of the oxidation numbers SO 4-2? What is the sum of the oxidation numbers in OH-? -3, -2, -1 13

14 Assign the oxidation number to all elements F 2 Cl - F Cl Na 2 O H 2 O H 2 O 2 Na O H O H O 14

15 Assign the oxidation number to all elements F 2 Cl - F 0 Cl -1 Na 2 O H 2 O H 2 O 2 Na +1 O -2 H +1 O -2 H +1 O -1 15

16 Now let s do Redox Ws #2 Oxidation Numbers 16

17 Redox Terminology OXIDATION loss of electron(s) by a species; increase in oxidation number. REDUCTION gain of electron(s); decrease in oxidation number. OXIDIZING AGENT electron acceptor; species is reduced. REDUCING AGENT electron donor; species is oxidized. 17

18 Mg + HCl MgCl 2 + H 2 Assign oxidation numbers to all elements Find the element with an oxidation number increases (this element is oxidized) Find the element with an oxidation number decreases (this element is reduced) 18

19 Mg + HCl MgCl 2 + H 2 Mg 0 Mg +2 H +1 H 0 (oxidation) (Reduction) Oxidation is losing electrons Mg 0 Mg e - Reduction is gaining electrons H +1 + e - H 0 19

20 Mg + HCl MgCl 2 + H 2 Reactant that is oxidized = Mg Reactant that is reduced = H+ The oxidizing agent = H + The reducing Agent = Mg 20

21 Standard Reduction Potentials The more positive the E for a half reaction, the greater the tendency for that reaction to occur as written. Therefore, the half reactions with the most positive reduction potentials are the most powerful oxidizing agents, while the half reactions with the most negative oxidation potentials are the most powerful reducing agents. 21

22 22

23 Mg + HCl MgCl 2 + H 2 Reactant that is oxidized = Mg Reactant that is reduced = H+ The oxidizing agent = H + The reducing Agent = Mg 23

24 Mg + HCl MgCl 2 + H 2 Will this reaction occur? Yes You must combine a strong oxidizing agent with a strong reducing agent. The element being reduced must be a stronger oxidizing agent than the element being oxidized. 24

25 Li + Zn +2 Li + + Zn Li 0 Li +1 Zn +2 Zn 0 (oxidation) (Reduction) Oxidation is losing electrons Li 0 Li e - Reduction is gaining electrons Zn e - Zn 0 25

26 Li + Zn +2 Li + + Zn Reactant that is oxidized = Li Reactant that is reduced = Zn +2 The oxidizing agent = Zn +2 The reducing Agent = Li 26

27 Li + Zn +2 Li + + Zn Will this reaction occur? Yes The element being reduced must be a stronger oxidizing agent than the element being oxidized. The element being oxidized must be a stronger reducing agent than the element being reduced. 27

28 Now do Redox Ws #3 Oxidizing and Reducing Agents 28

29 Balancing Redox Reactions 1. Assign oxidation numbers to all species 2. Write the oxidation half reaction 3. Write the reduction half-reaction 4. Balance for charge 5. Re-write equation with factors from step #4 6. Balance for mass. 29

30 Balancing Redox Reactions H 3 PO 2 + H + + Cr Cr +3 + P + H 2 O Step 1: Assign Oxidation numbers to everything H 3 PO 2 + H + + Cr Cr +3 + P + H 2 O 30

31 Balancing Redox Reactions H 3 PO 2 + H + + Cr Cr +3 + P + H 2 O Step 2 and 3: Write oxidation and reduction half reactions Cr Cr e - P e- P Step 4: Balance for Charge. Oxidation half-reaction Reduction half-reaction Cr Cr e - Oxidation half-reaction 3(P e- P) Reduction half-reaction 31

32 Balancing Redox Reactions H 3 PO 2 + H + + Cr Cr +3 + P + H 2 O Cr Cr e - Oxidation half-reaction 3(P +1 + e- P) Reduction half-reaction Step 4: Re-write the original equation with the factor from balancing for charge. 3H 3 PO 2 + H + + Cr Cr P + H 2 O 32

33 Balancing Redox Reactions 3H 3 PO 2 + H + + Cr Cr P + H 2 O Step 6: Balance for mass. 3H 3 PO H + + Cr Cr P + 6H 2 O 33

34 Now do Redox Ws #5 Balancing Redox Reactions 34

35 Why Study Electrochemistry? Batteries Corrosion Industrial production of chemicals such as Cl 2, NaOH, F 2 and Al Biological redox reactions 35

36 Types of Cells Electrolytic cells are those in which electrical energy from an external source causes nonspontaneous chemical reactions to occur. Voltaic cells are those in which spontaneous chemical reactions produce electricity and supply it to an external circuit. 36

37 Oxidation-Reduction Reactions Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag + (aq) --> Cu 2+ (aq) + 2 Ag(s) 37

38 Electrodes Electrodes are surfaces upon which oxidation and reduction half reactions occur. The cathode is the electrode at which reduction occurs. The anode is the electrode at which oxidation occurs. 38

39 Voltaic Cells Voltaic cells are those in which spontaneous chemical reactions produce electricity and supply it to an external circuit. 39

40 Construction of Voltaic Cells In a voltaic cell, the two half reactions are contained in separate vessels; each vessel and its contents are referred to as half cells. 40

Construction of Voltaic Cells Zn -> Zn 2+ + 2e

41 Construction of Voltaic Cells Zn -> Zn e Oxidation - anode Cu e -> Cu Reduction - cathode 41

Voltaic Cell Zn --> Zn 2+ + 2e Cu 2+ + 2e --> Cu Zn +

42 Voltaic Cell Zn --> Zn e Cu e --> Cu Zn + Cu 2+ --> Zn 2+ + Cu Zn Zn 2+ (1.0 M) Cu 2+ (1.0 M) Cu 42

43 Voltaic Cell anode: oxidation cathode: reduction electrons: anode to cathode through wire anions move to anode through salt bridge cations move to cathode through salt bridge 43

44 Standard Conditions Under standard conditions, the voltage is commonly referred to as the standard emf or standard cell potential, E cell. The standard conditions of 25 C, 1 atm pressure, and 1 M solutions are assumed unless noted otherwise. 44

45 Electrochemical Cells An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. Batteries are voltaic cells 45

46 ELECTRIC CURRENT Cu 2+ ions Zn metal With time, Cu plates out onto Zn metal strip, and Zn strip disappears. Zn is oxidized and is the reducing agent Zn(s) Zn 2+ (aq) + 2e- Cu 2+ is reduced and is the oxidizing agent Cu 2+ (aq) + 2e- Cu(s) 46

47 ELECTRIC CURRENT Zn metal Cu 2+ ions Electrons are transferred from Zn to Cu 2+, but there is no useful electric current. Oxidation: Zn(s) Zn 2+ (aq) + 2e- Reduction: Cu 2+ (aq) + 2e- Cu(s) Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) 47

48 ELECTRIC CURRENT To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. Zn Zn 2+ ions wire electrons salt bridge Cu Cu 2+ ions This is accomplished in a VOLTAIC cell. A A group of such cells is called a battery. 48

49 Zn wire electrons salt bridge Cu Zn 2+ ions Cu 2+ ions Electrons travel thru external wire. Salt bridge allows anions and cations to move between compartments. Anions move towards anode. Cations move towards cathode. Salt bridge maintains electrical neutrality. 49

50 Electrons move from anode to cathode in the wire. Anions & cations move thru the salt bridge. Electrochemical Cell 50

51 CELL POTENTIAL, E 0 Zn and Zn 2+ anode Cu and Cu 2+ cathode Electrons are driven from anode to cathode by an electromotive force. For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 C C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. 51

52 Electrochemical Cell 52

53 Al metal in AlCl 3 Pb metal in Pb(NO 3 ) 2 The stronger reducing agent is oxidized The stronger oxidizing agent is reduced Anode Oxidation Reaction Al 3e - + Al +3 Cathode Reduction Reaction Pb e - Pb 53

54 E 0 cell = E 0 cathode - E 0 anode Look up standard Reduction Potentials Anode Al 3e - + Al Volts Cathode Pb e - Pb Volts E 0 cell = E 0 cathode - E 0 anode E 0 cell = (-1.66) E 0 cell = Volts 54

55 Ni metal in NiCl 2 Cu metal in CuSO 4 The stronger reducing agent is oxidized The stronger oxidizing agent is reduced Anode Oxidation Reaction Ni 2e - + Ni +2 Cathode Reduction Reaction Cu e - Cu 55

56 E 0 cell = E 0 cathode - E 0 anode Look up standard Reduction Potentials Anode Ni 2e - + Ni Volts Cathode Cu e - Cu 0.34 Volts E 0 cell = E 0 cathode - E 0 anode E 0 cell = (-0.25) E 0 cell = +.59 Volts Overall Ni + Cu +2 Cu + Ni Volts 56

57 Cell Potential The sign of the E 0 cell tells you the direction of electron flow. A positive E 0 value means that the reaction will occur naturally in the direction written. The voltaic cell will work A negative E 0 value means the reaction will not occur naturally in the direction written. The voltaic cell will not work. 57

58 Zn - Volts Salt Bridge + H 2 Zn 2+ H + Zn Zn e- OXIDATION ANODE 2 H + + 2e- H 2 REDUCTION CATHODE Overall reaction is reduction of H + by Zn metal. Zn(s) + 2 H + (aq) Zn 2+ + H 2 (g) Therefore, E o for Zn Zn 2+ (aq) + 2e- is V. V E o = V Zn is a better reducing agent than H 2. 58

Zn/Cu Electrochemical Cell wire Anode, negative, source of electrons Zn Zn 2+ ions electrons salt bridge Cu Cu 2+ ions Cathode, positive, sink for electrons Zn(s) Zn 2+ (aq)

59 Zn/Cu Electrochemical Cell wire Anode, negative, source of electrons Zn Zn 2+ ions electrons salt bridge Cu Cu 2+ ions Cathode, positive, sink for electrons Zn(s) Zn 2+ (aq) + 2e- E o = V Cu 2+ (aq) + 2e- Cu(s) E o = V Cu 2+ (aq) + Zn(s) Zn 2+ (aq) + Cu(s) E o = V 59

60 Uses of E o Values 1. Decide on relative ability of elements to act as reducing agents (or oxidizing agents) 2. Assign a voltage to a half-reaction that reflects this ability. Zn wire electrons salt bridge Cu Zn 2+ ions Cu 2+ ions 60

61 E o for a Voltaic Cell Volts Cd Salt Bridge Fe Cd 2+ Cd --> Cd e- or Cd e- --> Cd Fe 2+ Fe --> Fe e- or Fe e- --> Fe 61

62 E o for a Voltaic Cell Cd Cd 2+ Volts Salt Bridge Fe 2+ Fe From the table, you see Fe is a better reducing agent than Cd Cd 2+ is a better oxidizing agent than Fe 2+ Overall reaction Fe + Cd 2+ Cd + Fe 2+ E o = V 62

63 Now let s do Redox Ws #7 Voltaic Cell Potentials 63

64 Electrolysis Electrolytic cells are those in which electrical energy from an external source causes nonspontaneous chemical reactions to occur. An electrochemical reaction that is nonspontaneous can be driven to proceed by applying an external electrical supply. Electrolytic Cell has a battery that gives the energy for the nonspontaneous reaction 64

65 Electrolysis of Aqueous NaCl NaCl in H 2 O Two possible reducing agents Cl- or O -2 The best reducing agent is Cl -, therefore, Cl - is oxidized. Two possible Oxidizing agents H + or Na + The best oxidizing agent is H +, Therefore H + is reduced. 65

66 Electrolysis of Aqueous NaCl The best reducing agent is Cl + Anode (+) Oxidation 2Cl - Cl 2 (g) 2e - The best oxidizing agent is H + Cathode (-) Reduction 2H + + 2e - H 2 (g) 66

67 Electrolysis electrons Anode + BATTERY Cl - Na + Cathode Electrolysis of molten NaCl. Here a battery pumps electrons from Cl - to Na +. 67

Oxidation & Reduction (Redox) Notes

Oxidation & Reduction (Redox) Notes Chemical Activity (or Chemical Reactivity) is the measure of the reactivity of elements. If an element has high activity, then it means that the element is willing to