Redox reactions & electrochemistry

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Terence Goodman
6 years ago
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1 Redox reactions & electrochemistry

2 Electrochemistry Electrical energy ; Chemical energy oxidation/reduction = redox reactions

3 Electrochemistry Zn + Cu 2+ º Zn 2+ + Cu Oxidation-reduction reactions always involve transfer of electrons from one species to another. Species losing electrons is oxidized Species gaining electrons is reduced

4 Electrochemistry Zn + Cu 2+ º Zn 2+ + Cu OXIDATION: loss of electrons REDUCTION: gain of electrons Also change in oxidation number

5 Electrochemistry Zn + Cu 2+ º Zn 2+ + Cu Oxidizing agent: oxidizes another species; it is itself reduced. Reducing agent: reduces another species; it is itself oxidized.

6 Electrochemistry loss of 2 e - Oxidation Zn + Cu 2+ º Zn 2+ + Cu gain of 2 e - Reduction Zn reducing agent Cu 2+ oxidizing agent

7 Electrochemistry Mg + 2 HCl º MgCl 2 + H 2 Mg 0 + 2H + Cl - º Mg 2+ Cl H 0 2 Mg oxidized to Mg 2+ H + reduced to H 2

8 Electrochemistry Chemical energy 6 Electrical energy

9 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Write unbalanced ionic equation 2. Separate into half-reactions

10 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Write unbalanced ionic equation 2. Separate into half-reactions Oxidation

11 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Write unbalanced ionic equation 2. Separate into half-reactions Reduction

12 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Balance where possible Fe 2+ º Fe 3+ oxidation Cr 2 O 7 2- º 2 Cr 3+ reduction

13 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Add water to balance oxygen Cr 2 O 7 2- Cr 2 O 7 2- º 2 Cr 3+ º 2 Cr H O 2

14 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Add H + to balance hydrogen 14 H + + Cr 2 O 7 2- º 2 Cr H 2 O 5. Check charges

15 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Add H + to balance hydrogen 14 H + + Cr 2 O e - º 2 Cr H 2 O 5. Check charges

16 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Compare half reactions 14 H + + Cr 2 O e - º 2 Cr H 2 O Fe 2+ º Fe 3+ + e -

17 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Compare half reactions 14 H + + Cr 2 O e - º 2 Cr H 2 O 6 Fe 2+ º 6 Fe e -

18 Balancing redox equations Fe 2+ + Cr 2 O 7 2- acid º Fe 3+ + Cr Add half reactions X 14 H + + Cr 2 O e - º 2 Cr H 2 O X 6 Fe 2+ º 6 Fe e - 14 H + + Cr 2 O Fe 2+ º 2 Cr H 2 O + 6 Fe 3+

19 Balancing redox equations In basic solutions repeat steps 1-3 Add OH - both sides, 1 for each H + Combine OH - & H + 6 H 2 O (eg 19.1)

20 Types of electrochemical cells Oxidation reduction reactions separated into two half-reactions. Electrochemical cell: electrodes dip into an electrolyte in which a chemical reaction either uses or generates an electric current.

21 Types of electrochemical cells The force with which electrons travel from the oxidation halfreaction to the reduction halfreaction is measured as voltage.

22 Types of electrochemical cells 1. Galvanic or Voltaic Spontaneous reactions Produces electrical energy

23 Types of electrochemical cells 2. Electrolytic Non-spontaneous reactions Requires electrical energy

24 Types of electrochemical cells 2. Electrolytic Many reactive metals obtained by electrolysis of a molten salt Li, Mg, and Ca metals obtained by the electrolysis of chlorides

25 A simple example Cu + Ag + 9 Cu 2+ + Ag oxidant? reductant?

26 A simple example Cu 2+ + Zn 6 Cu + Zn 2+

27 Electrochemical cells To get better control over the system, each half-reaction placed in separate cell Connected by salt bridge

28 Electrochemical cells oxidation at anode Zn 6 Zn e -

29 Electrochemical cells reduction at cathode Cu e - 6 Cu

30 Electrochemical cells Zn*Zn 2+ (1M)2Cu 2+ (1M)*Cu

31 Standard hydrogen electrode Reference electrode Compare to other half-cells Pt*H 2 (1atm), 1M2

32 Some theory Galvanic cells push electrons through circuit. Magnitude of this ability called potential. Potential or electromotive force (EMF) given in volts (V). Volt defined as amount of energy (J) per unit of charge (coulombs)

33 Zinc -copper reaction E = reduction potential for a half reaction measures how willing a species is to gain or loss electrons Everything compared to hydrogen 2H + + 2e - 6 H 2

34 Cu e - º Cu Zn º Zn e - E o = +0.34V E o = -0.76V

36 Calculating cell potential E (cell) = E (red) - E (ox) using values on previous table Calculate E for Zn + Cu +2 6 Zn +2 + Cu OX: Zn(s) 6 Zn e - RED: Cu e - 6 Cu

37 º F º F

38 Calculating cell potential E (cell) Calculate E for Zn + Cu +2 6 Zn +2 + Cu OX: Zn e - 6 Zn(s) E = V RED: Cu e - 6 Cu E = +0.34V Have to reverse Zn equation

39 Calculating cell potential E (cell) using values on previous table Calculate E for Zn + Cu +2 6 Zn +2 + Cu OX: Zn e - 6 Zn V RED: Cu e - 6 Cu +0.34V E = V - (-0.76 V)= V

40 Batteries Portable voltaic cells D chemicals in a paste or solid W liquid solution Zinc-carbon dry cell 1.5 V

41 Batteries Zinc-carbon dry cell Zn + 2 MnO 2 + 2NH Zn 2+ + Mn 2 O 3 + 2NH 3 + H 2 O

42 Batteries Lead storage cell Large capacity and high current Rechargeable H 2 SO 4 6 H 2 O Density to 1.0 Hydrometer Cold weather: ions move slowly due to increased viscosity

43 Spontaneity of redox reactions How is õe cell related to )GE and K? )GE = -nf õe cell (1) n = moles e - transferred F = Faraday constant = 96,500 C/mol )GE (-) for spontaneous reaction 6 (+) õe cell

44 Spontaneity of redox reactions )GE = -2.3 RTlogK (2) -nf õe cell = -2.3 RTlogK õe cell = 2.3 RTlogK (3) nf At 298 K and using R, F values... õe cell = logk n Use (1), (2), (3) to find variables

45 Nernst equation Used for non-standard conditions -nf õe cell = -2.3 RTlogK õ =õe logq cell n Q is reaction quotient

46 Corrosion Metal deterioration due to electrochemical reactions Iron rusting Fe O H O Fe + O 2 + 4H Fe H 2 O õe = 1.67v

47 Corrosion Metal deterioration due to electrochemical reactions Silver tarnishing Ag 6 Ag 2 S

48 Corrosion Metal deterioration due to electrochemical reactions Cu/brass turning green Cu 6 CuCO 3

49 Corrosion 2Fe + O 2 + 4H Fe H 2 O Where does the acid come from? H 2 CO 3 Fe 2+ further oxidized by O 2 & H 2 O Rusting enhanced by salt water (nails in water) Aluminum doesn t rust Al 2 O 3

50 1. Painting Protecting metals

51 Protecting metals 2. Cu, Sn, Zn on Fe

52 Protecting metals 3. (steel)

53 Protecting metals 4. Cathodic protection connect Fe to Zn or Mg, more easily oxidized (ship hulls)

54 Electrolysis Electrical energy used to make non-spontaneous reactions occur Molten NaCl 2 NaCl (l) 6 2 Na (l) + Cl 2 (g)

55 Electrolysis Electrical energy used to make non-spontaneous reactions occur Water 2 H 2 O (l) 6 2 H 2 (g) + O 2 (g) )GE =424 Anode: 2 H 2 O 6 4H + + O 2 + 4e - Cathode: 2H + + 2e - 6 H 2 2

56 Electrolysis Aqueous NaCl What is present? Na + Cl - H O 2 2 anode reactions possible 2 H O 6 4H + + O + 4e Cl - 6 Cl + 2e - 2 At high [Cl - ]: Cl produced 2

57 Electrolysis Aqueous NaCl What is present? Na + Cl - H O 2 2 cathode reactions possible Na + + e - 6 Na 2 H O + 2e OH - + H 2 2 At high [Cl - ]: Cl produced 2

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem.

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem. 1 Electrochemistry Oxidation-Reduction Review Topics Covered Oxidation-reduction reactions Balancing oxidationreduction equations Voltaic cells Cell EMF Spontaneity of redox reactions Batteries Electrolysis