1. Chapter six Electrochemistry

2. Oxidation NumbersOxidation Numbers Oxidation-reduction reactionOxidation-reduction reaction Oxidizing agent and reducing agentOxidizing agent and reducing agent Redox coupleRedox couple 6-1 Oxidation-reduction Concepts

3. ● Oxidation Numbers ─ The charge on an atom or a monatomic ion ─ The charge that an atom in a substance would have if the shared pair of electrons belonged to the more electronegative atom in the bond

4. Rules for Assigning an Oxidation Number (Ox#) 1. For an atom in its elemental form Ox # = 0: (O 2,Cl 2, H 2, ) (not combine with different element) (O 2,Cl 2, H 2, ) (not combine with different element). For a monatomic ion: Ox # = ion charge 2. For a monatomic ion: Ox # = ion charge Ca 2+ +2; Br -, -1; Ca 2+, +2; Br -, -1;  Ox. # = charge of molecule or ion. Sum of oxidation states = 0 in neutral compounds General rules Sum of oxidation states = charge of the ion H2SO4, Cr2O72-

5. (1) For oxygen: Ox# = -2 in most compounds Ox# = -1 in peroxides (H 2 O 2 ) Ox# = -1 in peroxides (H 2 O 2 ) 4. Rules for specific atoms or periodic table groups (2) For hydrogen: Ox# = +1 in combination with nonmetals Ox# = -1 in combination with metals Ox# = -1 in combination with metals with electropositive element (i.e., Na, K) H =-1 with electropositive element (i.e., Na, K) H =-1 (3) For fluorine: Ox# = -1 in all compounds Except with F, K OF 2 ; KO 2 (4) For Group 1A: Ox# = +1 in all compounds (5) For Group 2A: Ox# = +2 in all compounds (6) For Group 7A: Ox# = -1 in most compounds

6. x = +3 (for N) x + 2(-2) = -1 NO 2 - x = +5 (for N) X + 3(-2) = -1 NO 3 - x = +3 (for P) x + 3(-2) = -3 PO 3 3- x = +5 (for P) x + 4(-2) = -3 PO 4 3- x = +4 (for S) x + 3(-2) = -2 SO 3 2- x = +6 (for S) x + 4(-2) = -2 SO 4 2- x = +4 (for C) +1 + x + 3(-2) = -1 HCO 3 - Calculation of Oxidation Numbers:

7. Example 1 Al 2 S 3 Al is a monatomic ion with a 3+ charge, so its oxidation state is +3 (Rule 2). When combined with metals in binary compounds, S is a monatomic ion with a 2- charge, so its oxidation state is -2 (Rule 2). 2 Al at +3 each = +6 3 S at -2 each = -6 sum = 0 +3-2 Reminder: Nonmetals (like sulfur) as well as metals not in group I or II can have many oxidation states, so they must be carefully analyzed.

8. Example 2 Na 2 CO 3 Oxygen has a -2 oxidation state (Rule 1). Na is a monatomic ion with a 1+ charge, so its oxidation state is +1 (Rule 4). 2 Na at +1 each = +2 3 O at -2 each = -6 sum = -4 The overall sum must be 0 for a compound (Rule 2), so carbon must have a +4 oxidation state. +1-2+4

9. ● Oxidation-reduction (redox) reaction: Redox reaction: A reaction in which one or more electrons are transferred from one atom to another. Redox reaction: A reaction in which oxidation numbers of some elements are changed in a reaction process.

10. 2FeCl 3 + SnCl 2 = 2FeCl 2 + SnCl 4 2Fe 3+ + Sn 2+ = 2Fe 2+ + Sn 4+ Redox reaction 2Mg + O 2 = 2MgO Electron transfer Cu(s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag(s) L oss of E lectrons = O XIDATION ( LEO ) G ain of E lectrons = R EDUCTION ( GER ) - 2 e - 2 x +1 e -

11. Redox reaction Reduction reaction Oxidation reaction Oxidation reaction ─ loss of electrons ─ oxidation number increases A g →A g + + e - ─ gain in electrons ─ decrease in oxidation number Fe 2+ + 2e - →Fe Redox process always occurs together. In redox process, one can’t occur without the other. Half- reaction

13. Oxidation Reduction

14. Oxidizing agent ● Oxidizing agent and reducing agent ─ electron acceptor; species is reduced. Reducing agent ─ electron donor, species is oxidized.

15. oxidizing agent reducing agent Loss of 2 e -1 oxidation Gain of 2 e -1 reduction –An oxidizing agent is a species that oxidizes another species; it is itself reduced. –A reducing agent is a species that reduces another species; it is itself oxidized.

16. Oxidizing agent: Reducing agent: Losses one or more electrons Causes reduction Undergoes oxidation Becomes more positive or less negative Gains one or more electrons Causes oxidation Undergoes reduction Becomes more negative or less positive

17. Example CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2H 2 O (l) CHOCHOCHOCHO -4+10 +4+1-2 Which species is oxidized ? (lost electrons/Ox state became more positive) Which species is reduced ? (gained electrons/Ox state became more negative)

18. ● Redox couple (Pair of Electrons Oxidation and Reduction) 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Oxidation: 5Fe 2+  5Fe 3+ + 5e  The overall reaction is split into two half-reactions, one involving oxidation and one reduction.

19. 2Fe 3+ + Sn 2+ = 2Fe 2+ + Sn 4+ Reduction: Fe 3+ + e - →Fe 2+ Oxidation: Sn 2+ - 2e - →Sn 4+ Redox couples Oxidized species / reduced species Sn 4+ / Sn 2+ ; Fe 3+ / Fe 2+ The oxidized and reduced states of each substance taking part in a half-reaction form a redox couple. Notation of redox couple

20. Ox + n e - Red Oxidized state + n e - Reduced state Redox couple half-reactions:

21. 6-2 Voltaic Cells (Primary Cells) Electron transfer Zn(s) + CuSO 4 (aq)  ZnSO 4 (aq) + Cu(s) - 2 e - +2 e - L oss of E lectrons = O XIDATION ( LEO ) G ain of E lectrons = R EDUCTION ( GER )

22. What could happen when we put a piece of zinc metal into the solution of copper sulfate?

23. Copper is deposited on the zinc The copper plates out onto the The copper plates out onto the surface of the zinc metal surface of the zinc metal The blue copper ( Ⅱ ) ions are gradually replaced by colorless zinc ions Chemical energy (reaction enthalpy) is released as heat

24. CHEMICAL CHANGE  ELECTRIC CURRENT 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 - 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) Cu 2+ is reduced and is the oxidizing agent Cu 2+ (aq) + 2e -  Cu(s)

25. Oxidation: Zn(s)  Zn 2+ (a q) + 2e - Reduction: Cu 2+ (a q) + 2e -  Cu(s) ---------------------------------------------------- Zn(s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu(s) Electrons are transferred from Zn to Cu 2+, but there is no useful electric current. CHEMICAL CHANGE  ELECTRIC CURRENT (2) 2 e -

26. To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. CHEMICAL CHANGE  ELECTRIC CURRENT (2) This is accomplished in a VOLTAIC cell. (also called GALVANIC cell) A group of such cells is called a battery.

27. Voltaic Cell A device in which chemical energy is changed to electrical energy. energy. 6-2.1 Cu- Zn Primary Cell

28. (-) (+)

29. Electrons travel thru external wire. Electrons travel thru external wire. Salt bridge allows anions and cations to Salt bridge allows anions and cations to move between electrode compartments. move between electrode compartments. This maintains electrical neutrality. This maintains electrical neutrality. ANODE (-) CATHODE (+) Negative electrode generates electron Oxidation Occur Positive electrode accepts electron Reduction Occur Zn half-cellCu half-cellZn 2+ / ZnCu 2+ / Cu

30. Voltaic Cells A voltaic cell consists of two half-cells. Each half-cell is a portion of the electrochemical cell in which a half-reaction takes place. A simple half-cell can be made from a metal strip dipped into a solution of its metal ion. For example, the zinc-zinc ion half cell consists of a zinc strip dipped into a solution of a zinc salt.

31. Another simple half-cell consists of a copper strip dipped into a solution of a copper salt. In a voltaic cell, two half-cells are connected in such a way that electrons flow from one metal electrode to the other through an external circuit.

32. As long as there is an external circuit, electrons can flow through it from one electrode to the other. Because zinc has a greater tendency to lose electrons than copper, zinc atoms in the zinc electrode lose electrons to form zinc ions. The electrons flow through the external circuit to the copper electrode where copper ions gain the electrons to become copper metal.

33. The two half-cells must also be connected internally to allow ions to flow between them. Without this internal connection, too much positive charge builds up in the zinc half-cell (and too much negative charge in the copper half-cell) causing the reaction to stop. Figure A and B show the two half-cells of a voltaic cell connected by salt bridge. A salt bridge is a U shape tube of an electrolyte in a gel that is connected to the two half-cells of a voltaic cell.A salt bridge is a U shape tube of an electrolyte in a gel that is connected to the two half-cells of a voltaic cell.

34. The salt bridge allows the flow of ions but prevents the mixing of the different solutions that would allow direct reaction of the cell reactants.

35. 6-2.2 Cell Reaction The two half-cell reactions, as noted earlier, are: oxidation half-reaction reduction half-reaction

36. Note that the sum of the two half-reactions Note that electrons are given up at the anode and thus flow from it to the cathode where reduction occurs. is the net reaction that occurs in the voltaic cell; it is called the cell reaction.

37. 6-2.3 Notation for Voltaic Cells It is convenient to have a shorthand way of designating particular voltaic cells. The anode (oxidation half-cell) is written on the left. The cathode (reduction half- cell) is written on the right. The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written anodecathode （－）（＋）

38. A boundary between different phases (e.g., an electrode and a solution) is represented by a single vertical line (│) The boundary between half-cell compartments, usually a salt, is represented by a double vertical line ( ) The cathode in a voltaic cell has a positive sign The anode in a voltaic cell has a negative sign because electrons flow from it.

39. Notation for Voltaic Cells The two electrodes are connected by a salt bridge, denoted by two vertical bars. The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written anodecathode salt bridge

40. Notation for Voltaic Cells The cell terminals are at the extreme ends in the cell notation. The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written anodecathodesalt bridge

41. Notation for Voltaic Cells A single vertical bar indicates a phase boundary, such as between a solid terminal and the electrode solution. The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written anodecathodesalt bridge

42. 2Fe 3+ (c 1 ) + Sn 2+ (c 2 ) = 2Fe 2+ (c 3 ) + Sn 4+ (c 4 ) Half reaction Fe 3+ + eFe 2+ Sn 2+ - 2eSn 4+ Sn 4+ / Sn 2+ Fe 3+ / Fe 2+ Pt│Sn 4+ (c 4 ), Sn 2+ (c 2 ) Pt│Fe 3+ (c 1 ), Fe 2+ (c 3 ) Cell notation Pt│Sn 4+ (c 4 ), Sn 2+ (c 2 )Fe 3+ (c 1 ), Fe 2+ (c 3 ) │Pt

43. When the half-reaction involves a gas, an inert material such as platinum serves as a terminal and an electrode surface on which the reaction occurs. hydrogen electrode The cathode half-reaction is

44. The notation for the hydrogen electrode, written as a cathode, is To write such an electrode as an anode, you simply reverse the notation.

45. To fully specify a voltaic cell, it is necessary to give the concentrations of solutions and the pressure of gases. In the cell notation, these are written in parentheses ( ). For example,

46. Line Notation solid  Aqueous  Aqueous  solid solid  Aqueous  Aqueous  solid Anode on the left  Cathode on the right Single line different phases. Double line salt bridge. If all the substances on one side are aqueous, a platinum electrode is indicated. Cu(s)  Cu +2 (aq)  Fe +2 (aq),Fe +3 (aq)  Pt(s)

47. 6-2.4 Electromotive Force The maximum potential difference between the electrodes of a voltaic cell is referred to as the electromotive force (emf) of the cell, denoted EThe maximum potential difference between the electrodes of a voltaic cell is referred to as the electromotive force (emf) of the cell, denoted E E = φ cathode – φ anode = φ + -φ -E = φ cathode – φ anode = φ + -φ - E is a positive number.E is a positive number.

48. The standard emf, E o, is the emf of a cell operating under standard conditions of concentration (1 M), pressure (1atm), of concentration (1 M), pressure (1atm), and temperature (25 o C). and temperature (25 o C).

49. Standard Notation for Electrochemical Cells ANODE Zn / Zn 2+ // Cu 2+ / Cu CATHODE OXIDATION Anode electrode Active electrolyte in oxidation half-reaction Cathode electrode Active electrolyte in reduction half-reaction Salt bridge Phase boundary REDUCTION

50. Anode and Cathode OXIDATION occurs at the ANODE. REDUCTION occurs at the CATHODE. Mnemonic: O and A are vowels; R and C are consonants

51. 6-2.5 Types of Electrodes (a) metal/metal ion electrode (b) metal/ insoluble salt electrode (c) gas electrode (d) redox electrode

52. Types of Electrode (continued)

53. l l l l l

54. The types of electrode The types of electrode Metal-metal ion electrode Zn(s) ∣ Zn 2+ ( aq ); Cu(s ) ∣ Cu 2+ ( aq ) Zn(s) ∣ Zn 2+ ( aq ); Cu(s ) ∣ Cu 2+ ( aq ) M ∣ M n+ : M n+ + ne == M M ∣ M n+ : M n+ + ne == M Gas electrode Pt ∣ H 2 (p) ∣ H + (c) 2H + + 2e == H 2 Pt ∣ H 2 (p) ∣ H + (c) 2H + + 2e == H 2 Metal-insoluble salt electrodes Pt ∣ Hg(l) ∣ Hg 2 Cl 2 (s) ∣ Cl - (c) Pt ∣ Hg(l) ∣ Hg 2 Cl 2 (s) ∣ Cl - (c) Hg 2 Cl 2 (s)+ 2e == 2Hg + 2Cl - Hg 2 Cl 2 (s)+ 2e == 2Hg + 2Cl - Oxidation-reduction electrodes Fe 3+ + e == Fe 2+ Pt ∣ Fe 3+ (c 1 ),Fe 2+ (c 2 ) Fe 3+ + e == Fe 2+ Pt ∣ Fe 3+ (c 1 ),Fe 2+ (c 2 )