Assigning Oxidation Numbers:

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Phyllis Lloyd
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1 Assigning Oxidation Numbers: 1. Oxidation number of a free element or diatomic molecule is zero. Ex: Na(s), Cu(s), H 2 (g), F 2 (g) 2. In most cases the oxidation number of hydrogen is +1, oxygen is -2, and fluorine is -1 when combined with another element. 3. The sum of the oxidation numbers of each of the elements in a molecule or ion must equal the charge.

2 Reduction-Oxidation Reactions (REDOX): Oxidation- Process in which oxidation state of an element increases. Species loses electrons. Reduction- Process in which oxidation state of an element decreases. Species gains electrons.

3 Using Oxidation Numbers: Ex: Zn(s) + Cu 2+ Zn 2+ + Cu(s) Zn(s): oxidized(lost electrons). Cu 2+ (aq): reduced(gained electrons). Ex: 2H 2 (g) + O 2 (g) 2H 2 O(l) H 2 (g): oxidized(lost electrons). O 2 (g): reduced(gained electrons).

4 REDOX cont : OXIDATION Zn(s) Zn e - REDUCTION Cu e - Cu(s) REDOX Zn(s) + Cu 2+ Zn 2+ + Cu(s) Zn(s): oxidized/reducing agent. Cu 2+ (aq): reduced/oxidizing agent.

5 Writing Balanced Redox Reactions: Oxidation and reduction reactions occur together. Occur in acidic or basic medium. Ex: (acidic) SO MnO 4 - SO Mn 2+ STEP 1: Identify the oxidized and reduced species and write the corresponding half reactions.

6 Writing Balanced Redox Reactions cont : STEP 2: Balance each of the half reactions. First atoms other than H and O. Balance O atoms by adding H O molecules and then 2 balance H atoms by adding H + ions. STEP 3: Balance the number of electrons. STEP 4: Add both half reactions and simplify.

7 Writing Balanced Redox Reactions cont Balance the following redox reaction which occurs in a basic medium. CrO S 2- Cr(OH) 3 (s) + S(s) NOTE: In basic medium add an equal number of OH - ions to both sides to neutralize H + ions. OH - + H + H 2 O

8 Electrochemical Cells: Consider, Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Zn(s) Zn 2+ (aq) + 2e - Cu 2+ (aq) + 2e - Cu(s)

9 Electrochemical Cells: Electrode- Strip of metal. Half cell-strip of metal in contact with its ion. Salt bridge- Allows passage of charge but not reactants. Anode: Zn(s) Zn 2+ (aq) + 2e - Cathode: Cu 2+ (aq) + 2e - Cu(s)

10 Electrochemical Cells cont : - +

11 Electrode Potentials: The voltage recorded by an electrochemical cell is referred to as the electromotive force(emf) and given the following symbol. E cell

12 Describing Electrochemical Cells (Cell Diagram): ANODE(OXIDATION) on left. CATHODE(REDUCTION) on right. indicates phase change. indicates salt bridge. M(s) M n+ (aq) M n+ (aq) M(s) (anode) (cathode)

13 Standard Electrode Potentials: Reaction occurs under standard conditions: 25 C and all substances are at unit concentration (1 M for all ions and 1 atm for all gases). Measured potential given the following symbol. E o cell The cell potential can be broken up into two components or half cells. E o cell E o oxidation E o reduction

14 Standard Hydrogen Electrode(SHE): 2H + (aq)(1 M) + 2e - H 2 (g)(1 atm) E o = 0 V

15 SHE/Cu o E cell V

16 Zn/SHE o E cell V

17 Standard Reduction Potentials: Can now determine the potential for each half cell. Tabulated and are referred to as standard reduction potentials. Ex:Calculate the standard cell potential for the following: Cd(s) + Cu 2+ (aq) Cd 2+ (aq) + Cu(s) Given: Cu 2+ (aq) + 2e - Cd 2+ (aq) + 2e - E reduc (V) Cu(s) V Cd(s) V

18 Standard Reduction Potentials cont Ex2: Calculate the standard cell potential for the following: 2Al(s) + 3Cu 2+ (aq) 3Cu(s) + 2Al 3+ Given: E (V) red Cu 2+ (aq) + 2e - Cu(s) V Al 3+ (aq) + 3e - Al(s) V

19 Predicting Spontaneous Redox Reactions: w electrical = nf E cell G = -nfe cell w = electrical work electrical n = # moles of electrons transferred F = Faraday constant(96485 C/mole) E = Voltage of cell. NOTE: 1 J = 1 C V cell G Gibbs Free Energy G<0 E cell = + # process process is is spontaneous G>0 E cell = - # process process is not is not spontaneous

20 Cu 2+ more likely to be reduced. Zn(s) + Cu 2+ Zn 2+ + Cu(s) o E cell 1.10 V spontaneous Reverse reaction Zn 2+ + Cu(s) Zn(s) + Cu 2+ o E cell 1.10 V not spontaneous Consider: Cu 2+ (aq) + 2e - Cu(s) Zn 2+ (aq) + 2e - Zn(s) E red (V) V V

21 Cell Potential as a Function of Concentration: Nernst Equation: E cell E cell Q V LogQ n Where if aa + bb cc + dd [ C] [ A] c a [ D] [ B] d b

22 Cell Potential as a Function of Concentration cont : Ex: Calculate E cell for the following voltaic cell. Zn(s) Zn 2+ (0.15 M) Cu 2+ (0.20 M) Cu(s) Ex:2 Calculate E cell for the following voltaic cell. Pt(s) H 2 (g)(0.5 atm) H + (0.50 M) Ag + (0.20 M) Ag(s)

23 Electrolysis: Electricity is used to cause a nonspontaneous redox reaction to occur. Ex: Electrolysis of molten sodium chloride. 2Na + (aq) + 2Cl - (aq) 2Na(s) + Cl 2 (g) E cell = V NOT SPONTANEOUS!!!

24 Electrolysis of Water: Ex: 2H 2 O(l) 2H 2 (g) + O 2 (g) E cell = -1.23V

25 Quantitative Electrolysis: Direct relationship between the amount of electricity used and the amount of products obtained from an electrolysis. E = voltage R: resistance(in ohms) E = IR I: current(in amps) 1 mole of electrons = 1 F(faraday) 1 F = C 1A = 1 C s -1

26 Quantitative Electrolysis Examples: Ex: 2Na + + 2Cl - 2Na(s) + Cl 2 (g) In the electrolysis of NaCl, how much Na(s) and Cl 2 (g) is produced by a current of A in 10.0 minutes.

27 Quantitative Electrolysis Examples: Ex: 2 Cu 2+ + H O O (g) + Cu(s) 2 2 Balance the following redox reaction and calculate the mass of O (g) and Cu(s) 2 produced in 30 sec by a current of 1.50 A. Ex: 3 Zn(s) + Cu 2+ Zn 2+ + Cu(s) If a current of 1.00 A is used, how many minutes does it take to electrolyze g of copper?

28 Corrosion: Deteriation of metals. Rust forms when Fe(s) is oxidized to Fe 2 O 3 and O 2 (g) is reduced to H 2 O. 4Fe(s) + 3O 2 (g) 2Fe 2 O 3 (brittle) E = + # cell occurs spontaneously Ag(s) Ag S 2 Cu(s) CuCO (patina) 3

17.1 Redox Chemistry Revisited

Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6