1. 1Chemistry 2C Lecture 5: April 7 th, 2010 I.Why cell potentials don’t add like balancing reactions II.Transition Metal Nomenclature Lecture 5: Electrochemistry & Transition Metals

2. 2Chemistry 2C Lecture 5: April 7 th, 2010 Electrochemistry ThermodynamicsEquilibria Kinetics The Chemistry Diamond

3. 3Chemistry 2C Lecture 5: April 7 th, 2010 Zn 2+ + 2e - -> Zn Cu 2+ (aq) + 2e - -> Cu (s) Write the half reactions like they were reduction reactions The Half Cell Potential is a Standard Reduction Potential OIL RIG Zn (s) | Zn 2+ (1.0M) || Cu 2+ (1.0M) | Cu (s) BUT, the actual reaction is Zn -> Zn 2+ + 2e - Cu 2+ (aq) + 2e - -> Cu (s) You can flip the equation and flip the SRP, but you do not want to for E of the cell

4. 4Chemistry 2C Lecture 5: April 7 th, 2010 Calculating Cell Potential Redux The cell potential can be constructed as the difference of two SRP. Zn (s) | Zn 2+ (1.0M) || Cu 2+ (1.0M) | Cu (s) The total cell potential (under standard conditions): Zn 2+ + 2e - -> Zn Cu 2+ (aq) + 2e - -> Cu (s) Write the half reactions like they were reduction reactions

5. 5Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) Given: E o =0.097V If the reactants and products are in their standard state except [H + ], will the reaction occur spontaneously as written? If PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 Reaction: a) [H + ] = 6.0M b) [H + ] = 1.2 M c) pH = 4.25 In a spontaneous reaction,  G 0! So we need to calculate the cell potential of the reaction under non-standard conditions!

6. 6Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) We are given that E o =0.097V PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 This is a non-standard conditions problem, we need the Nerst equation! In a spontaneous reaction,  G 0! So we need to calculate the cell potential of the reaction under non-standard conditions! And we need to know Q, E o, n before using the Nernst equation: By inspection of the balanced equation, n=2. If you cannot see this, then balance the redox equation step-by-step. But what is Q?

7. 7Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) Substituting into the Nernst Equation: PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 The reaction quotient for this reaction is (ignoring the solid PBO 2 and water) Side question: What is the reaction Quotient for a reaction at Standard condition? We can simplify this equation, since every species other than H + is at standard conditions

8. 8Chemistry 2C Lecture 5: April 7 th, 2010 Q: Under what concentration of [H + ] will the system not be spontaneous? Logarithm properties Using Standard Conditions Plug and Chug for each case! PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 Nernst Eq. Logarithm properties Arithmetic

9. 9Chemistry 2C Lecture 5: April 7 th, 2010 Extensive vs. Intensive Properties An extensive property is a property that varies as the amount of the material changes. For example? When adding, and subtracting half-reactions and multiplying by factors to make the equations balance, do NOT multiply the cell potentials like you would for enthalpies or free energies. An intensive property is a property that does not vary as the amount of the material changes. For example? Mass, Volume, Enthalpies, … Temperature, Pressure, Average Speed of Molecules, …

10. 10Chemistry 2C Lecture 5: April 7 th, 2010 Al (s) | Al 3+ (0.36M) || Sn 4+ (0.086M) | Sn 2+ (0.54 M) Al (s)-> Al 3+ (aq) + 3e - Sn 4+ (aq) + 2e - -> Sn 2+ (aq) Oxidation Reduction 2Al (s) + 3Sn 4+ (aq) -> 2Al 3+ (aq) + 3Sn 2+ (aq) Balanced eq. The number of electrons transferred, n, is an extensive property Half reactions The cell potential, E, is an intensive property Extensive vs. Intensive Properties For example: x3 x2 So n=3x2=6 But E o does not depend on multiplicative factors

11. 11Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) Given: E o =0.097V If the reactants and products are in their standard state except [H + ], will the reaction occur spontaneously as written? If PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 Reaction: a) [H + ] = 6.0M b) [H + ] = 1.2 M c) pH = 4.25 In a spontaneous reaction,  G 0! So we need to calculate the cell potential of the reaction under non-standard conditions!

12. 12Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) We are given that E o =0.097V PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 This is a non-standard conditions problem, we need the Nerst equation! In a spontaneous reaction,  G 0! So we need to calculate the cell potential of the reaction under non-standard conditions! And we need to know Q, E o, n before using the Nernst equation: By inspection of the balanced equation, n=2. If you cannot see this, then balance the redox equation step-by-step. But what is Q?

13. 13Chemistry 2C Lecture 5: April 7 th, 2010 Problem example (21.46) Substituting into the Nernst Equation: PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 The reaction quotient for this reaction is (ignoring the solid PBO 2 and water) Side question: What is the reaction Quotient for a reaction at Standard condition? We can simplify this equation, since every species other than H + is at standard conditions

14. 14Chemistry 2C Lecture 5: April 7 th, 2010 Q: Under what concentration of [H + ] will the system not be spontaneous? Logarithm properties Using Standard Conditions Plug and Chug for each case! PbO 2 (s) + 2Cl - (aq) + 4H + (aq) -> Cl 2 (g) + Pb 2+ (aq) + 2H 2 0 Nernst Eq. Logarithm properties Arithmetic

15. 15Chemistry 2C Lecture 5: April 7 th, 2010 Chapter 24: Transition Metals Transition Metals f-block d-block Unlike the main group elements, the transition metals can adopt several different oxidation states depending on how many e - are removed.

16. 16Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature Transition Metals form complexes with other species in solution (and also as precipitates). These are also called coordination compounds. Often such complexes take the structure of a central transition metal ion that is surrounded by molecules (“ligands”) or ions to form a “complex.” Such complexes can be neutral, or charged (either positive or negative). Such complexes can be often thought of as an entity that is held together (some more strongly than others). Ligands can be firmly attached to the metal, or can be weakly attached (“labile”). http://en.wikipedia.org/wiki/Ligand

17. 17Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature Before simple chemical formulas were used to represent complexes, e.g. CoCl 3 · 6NH 3 CoCl 3 · 5NH 3 Now we know that these really are (spatially): CoCl 3 NH 3 H3NH3N H3NH3N CoCl 2 Cl NH 3 H3NH3N H3NH3N Note the bonding between the central Co metal and the Nitrogen of the amine (ammonia) group.

18. 18Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature Cobalt is closely associated with the NH 3 or the Cl 2 as drawn. Finding six ligands on a transition metal is very common. This geometry is called the octahedron. And such complexes are referred to as octahedral, from which you know the geometry. The NH 3 and one Cl - ligands are said to be in the “inner sphere.” That is they are attached to the metal via a covalent bond and not ionic. CoCl 3 NH 3 H3NH3N H3NH3N CoCl 2 Cl NH 3 H3NH3N H3NH3N Hint for the future: If the oxidation state of the metal changes, the geometry may also change.

19. 19Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature We write these complexes like this to designate “inner sphere” ligands. CoCl 3 · 6NH 3 If these are dissolved in solution CoCl 3 · 5NH 3 Uninformative [Co( NH 3 ) 6 ] Cl 3 [CoCl( NH 3 ) 5 ] Cl 2 More useful [Co( NH 3 ) 6 ] 3+ [CoCl( NH 3 ) 5 ] 2+ & Cl -, Cl -, Cl - & Cl -, Cl - The oxidation state of the transition metal. +3 Because Cl - is -1 and NH 3 is 0 (neutral)

20. 20Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature What do we call these complexes? hexaamminecobalt (III) chloride [Co( NH 3 ) 6 ] Cl 3 [CoCl( NH 3 ) 5 ] Cl 2 pentaamminechlorocobalt (III) chloride Six amine groups five amine groups one chloro group oxidation state of metal metal typecounter ion

21. 21Chemistry 2C Lecture 5: April 7 th, 2010 Transition Metal Nomenclature Question: What is the oxidation number of Rh in this complex? [RhCl( NH 3 ) 5 ] (NO 3 ) 2 NH 3 is neutral Cl - is -1 NO 3 - is -1 (formal charge of polyatomic ion) 02x(-1) The charge of the complex (inner sphere) is +2 since there is a -1 charge for each Nitrate ion Thus by adding charges: +2=-1+Rh -> Rh=+3 Name: Pentaamminechlororhodium (III) nitrate

22. 22Chemistry 2C Lecture 5: April 7 th, 2010 What is the complex is an anion? PtCl 4 2- Tetrachloroplatinate (II) Which can form a solid with potassium K 2 (PtCl 4 ) Potassium Tetrachloroplatinate (II)

23. 23Chemistry 2C Lecture 5: April 7 th, 2010 Common Ligands Azide, N 3 - Bromide, Br- Chloride, Cl - Cyanide, CN - Fluoride, F - Hydroxide, OH - Carbonate, CO 3 2- Oxalate, C 2 O 4 2- Oxide, O 2 - Name in complexes Neutral Ligands Azido Bromo Chloro Cyano Fluoro Hydroxo Carbonato Oxalato Oxo Name by itself Ammonia, NH 3 Caron monoxide, CO Water, H 2 0 Ammine Carbonyl Aqua Rules are given in lab practical!

24. 24Chemistry 2C Lecture 5: April 7 th, 2010 I.Batteries II.Corrosion III.Electrolysis Lecture 7: Electrochemistry & Transition Metals

25. 25Chemistry 2C Lecture 5: April 7 th, 2010 Electrochemistry revisited Batteries: 1.Self Contained, portable electrochemical power sources (galvanic cells) 2. Comprised of one or more voltaic cells, often connected in series to get higher voltage (=electromotive force ) Remember Galvani battery which was a “pile” of Galvanic cells, all connected via salt bridges 1.5V Sum=3.0V

26. 26Chemistry 2C Lecture 5: April 7 th, 2010 Electrochemistry revisited Batteries: 3. Convention: The anode/cathode of the battery is the same definition as given previously with isolated Galvanic cells The lifetime of the battery depends on the quantities of the reactants (they wear down) Cathode (+) 4. Usually includes some kind of porous barrier and a conducting medium, called an electrolyte (the salt used in a salt bridge) Anode (-)

27. 27Chemistry 2C Lecture 5: April 7 th, 2010 Electrochemistry revisited Common batteries have two classes: A. Primary – discarded when EMF -> 0. They are not rechargeable (e.g. normal alkaline flashlight battery) B. Secondary – Rechargeable from external electrical sources (e.g. Ni-Cd, Li-ion, batteries)

28. 28Chemistry 2C Lecture 5: April 7 th, 2010 Electrochemistry revisited The 3 rd class of batteries: C. Fuel Cells – Not true battery because not self-contained, but often added as a 3 rd category. CH 4 (g) + 2O 2 -> CO 2 (g) + 2H 2 0 (l) (nat. gas cell) 2H 2 (g) + O 2 (g) -> 2H 2 O (l) (H 2 cell) More efficient than the best internal combustion engines in producing electricity, but so far also more expensive. Also called, “flow” batteries” because must supply constant source of reactants.

29. 29Chemistry 2C Lecture 5: April 7 th, 2010 Primary Batteries: Alkaline batteries: Improvement over the old flashlight batteries which had much shorter life due to discharge (short circuiting) even when not being used the Zn metal that made up the anode was also the battery container (can), and it got dissolved away (HN 3 was produced) Electrochemistry revisited Zn + KOH (Zn is powdered in a gel with KOH) At the anode, Zn is oxidized, since basic conditions: MnO 2 + graphite separated by a porous fabric Anode: Cathode: Today’s batteries: In a steel can… Zn (s) -> Zn 2+ (aq) + 2e - Zn 2+ (aq) + 2OH - (aq) -> Zn(OH) 2 (s)

30. 30Chemistry 2C Lecture 5: April 7 th, 2010 Primary Batteries: Electrochemistry revisited Zn + KOH (Zn is powdered in a gel with KOH) At the cathode, Mn is reduced: MnO 2 + graphite separated by a porous fabric Anode: Cathode: Today’s batteries: In a steel can… 2MnO 2 (s) + H 2 0 (l) + 2e - -> Mn 2 O 3 (s) + 2OH - (aq) Once Mn or Zn is depleted, the cell doesn’t give any voltage. (The cell is at equilibrium)

31. 31Chemistry 2C Lecture 5: April 7 th, 2010 Secondary Batteries: Lead-Acid (storage) battery is one example. You see under the hood of your car Electrochemistry revisited At the anode, Pb is oxidized, since basic conditions: Lead in three oxidation states: Pb, Pb 2+, PbO3 http://www.landiss.com/battery.htm 35% H 2 SO 4 by mass Contains Pb (s) + HSO 4 - -> PbSO 4 (s) + H+ (aq) + 2e- PbO 2 (s) + 3H+ (aq) + HSO 4- (aq) + 2e - -> PbSO 4 (s) + 2H 2 O(l) Anode: Cathode:

32. 32Chemistry 2C Lecture 5: April 7 th, 2010 Net: Electrochemistry revisited When the car is running, the battery is kept recharged by the alternator, a recharge (using external sources of electricity) can prolong the life of the battery, but eventually will fail because the PbSO4 drops off the electrodes and the acid becomes diluted. http://www.landiss.com/battery.htm What is the cell potential? PbO 2 (s) + Pb (s) + 2H+ (aq) + 2HSO 4- (aq) -> 2PbSO 4 (s) + 2H 2 O(l) Recharge: The direction of the reaction is reversed and E supplied > E cell

33. 33Chemistry 2C Lecture 5: April 7 th, 2010 Fuel Cells www.ballard.com Fuel cells generate electricity, just like batteries, but are not self- contained

34. 34Chemistry 2C Lecture 5: April 7 th, 2010 Unwanted oxidation of material Corrosion Oxidation http://www.landiss.com/battery.htm Cost of $200 a year to fix, especially for steel (alloy of Fe) O 2 is the “oxidizing agent”, figure that! Fe (s) -> Fe(OH) 2 (ferrous iron) Fe(OH) 2 -> Fe (OH) 3 (ferric iron) -> Fe 2 O 3 · H 2 0 (complex that is rust) O 2 (g) + 2H 2 O + 4e - -> 4OH - Reduction

35. 35Chemistry 2C Lecture 5: April 7 th, 2010 How to avoid corrosion: Corrosion Oxidation http://www.landiss.com/battery.htm Fe (s) -> Fe(OH) 2 (ferrous iron) Fe(OH) 2 -> Fe (OH) 3 (ferric iron) -> Fe 2 O 3 · H 2 0 (complex that is rust) O 2 (g) + 2H 2 O + 4e - -> 4OH - Reduction 1) Tin coating (“tin” cans)- add a thin coat of Sn (s) by dipping the can into liquid Sn 2) Cathodic Protection: A chunk of another metal is attached to the steel object and it is more “active” and preferentially oxidized. It then supplied e - to the Fe which acts as an inert electrode. This protects the Fe as long as the other metal exists (e.g. Mg, Zn, Al)

36. 36Chemistry 2C Lecture 5: April 7 th, 2010 How to avoid corrosion: Corrosion Oxidation Fe (s) -> Fe(OH) 2 (ferrous iron) Fe(OH) 2 -> Fe (OH) 3 (ferric iron) -> Fe 2 O 3 · H 2 0 (complex that is rust) O 2 (g) + 2H 2 O + 4e - -> 4OH - Reduction 1) Tin coating (“tin” cans)- add a thin coat of Sn (s) by dipping the can into liquid Sn. (Also galvanizing) Separate metal from O 2 and H 2 0 by adding a protection layer!

37. 37Chemistry 2C Lecture 5: April 7 th, 2010 How to avoid corrosion: Corrosion Oxidation Fe (s) -> Fe(OH) 2 (ferrous iron) Fe(OH) 2 -> Fe (OH) 3 (ferric iron) -> Fe 2 O 3 · H 2 0 (complex that is rust) O 2 (g) + 2H 2 O + 4e - -> 4OH - Reduction 2) Cathodic Protection: A chunk of another metal is attached to the steel object and it is more “active” and preferentially oxidized. It then supplied e - to the Fe which acts as an inert electrode. This protects the Fe as long as the other metal exists (e.g. Mg, Zn, Al)

38. 38Chemistry 2C Lecture 5: April 7 th, 2010 Electrolysis The process by which a chemical reaction occurs as a result of passing electric current through a solution (not a “galvanic cell”, but an “electrolytic cell”) HUGE industrial applications!!! 1)Most metals are produced from their ores by electrolysis (e.g. Al via the Hall process and Na via electrolysis of NaCl) 2)Electroplating (applying a thin layer of a metal over another metal) 3)Electrolysis of liquids (water)

39. 39Chemistry 2C Lecture 5: April 7 th, 2010 Electrolysis of water The energy of an externally applied energy source and breakdown water 2H 2 0 (l) -> 2H 2 (g) + O 2 (g) Cathode (where reduction occurs) but is negative (-) Anode (where oxidation occurs) but is positive (+) E cell o =-1.23V Not spontaneous. In practice we need > 1.23 V to split the water and overcome the “overpotential”

40. 40Chemistry 2C Lecture 5: April 7 th, 2010 Electrolysis of water Water is a poor conductor of electrons. So often a little electrolye (Na 2 SO 4 ) is added to carry charge SO 4 2- ions migrate toward the anode Cathode (where reduction occurs) but is negative (-) Anode (where oxidation occurs) but is positive (+) Na + ions migrate toward the cathode

41. 41Chemistry 2C Lecture 5: April 7 th, 2010 Electroplating (turning Pb to Gold) In electroplating, a metal is deposited on the cathode material For example, electroplating Ag onto steel Cathode reaction: Ag + (aq) + e - -> Ag (s) Anode reaction: Ag (s) -> e- + Ag + (s) Or the electrolysis reaction may occur: 2H 2 O -> O 2 (g) + 4H+ + 2e- But the more net favored reaction (takes less electricity) occurs: Ag (s) @ anode -> Ag (s) @ cathode http://en.wikipedia.org/wiki/Electroplating

42. 42Chemistry 2C Lecture 5: April 7 th, 2010 Faraday discovered that metal ions of salts are deposited as metals on the cathode in electrolysis n e =It/F I=current (amps) t=time (sec) F= Faraday’s constant (9.65x10 +4 C/mole) ne =# of moles of electrons Determine the mass of cooper that is platted out when a current of 10.0A is passed for 30 min through a solution containing Cu 2+ (aq). Cu 2+ (aq) + 2e - -> Cu (s) Therefore, 2 moles of transferred electrons is 1 mole of Cu platted n e = (10.0Amp)(30.0 min)(60s/min)/96,500 C/mole) = 0.186 mol of e- transferred

43. 43Chemistry 2C Lecture 5: April 7 th, 2010 n e = (10.0Amp)(30.0 min)(60s/min)/96,500 C/mole) = 0.186 mol of e- transferred Because 2 moles of transferred electrons equals 1 mole of Cu platted, then 0.186/2=0.093 moles of Cu platted 0.093 mole * 63.55 g/m (for Cu) = 5.93 g Cu platted Thus, the anode is 5.93 g heavier!