1. Chapter 18 Electrochemistry

2. Definition The study of the interchange of chemical and electrical energy. Involve oxidation-reduction reactions

3. Half-Reaction Method Used to balance redox rxns in aqueous solution One of the rxns is the oxidation rxn Other is the reduction rxn The half reactions are balanced separately and then added Differs in acidic vs. basic solution

4. In Acidic Solution 1.Separate the oxid. and red. parts of the rxn. 2.For each: a. Balance all elements but H and O b. Balance O using H 2 O c. Balance H using H + d. Balance charge using electrons 3.If needed, mult. One or both half-rxns. by a whole # to make electrons equal for both 4.Add half-rxns and cancel same on both sides 5.Make sure elements/charges are balanced

5. Acronym for Balancing in Acid Big White Hogs Eat Bagels and Cream Cheese B = balance all but H & O W = add water for O’s H = Add H E = Add electrons B = Make sure electrons match CC = Combine and cancel

6. Example in Acid Balance the following equation in acid solution using the half-reaction method. Cu (s) + HNO 3(aq) -> Cu 2+ (aq) + NO (g) Answer: 3Cu + 2HNO 3 + 6H + -> 3Cu 2+ + 2NO + 4H 2 O Cr 2 O 7 2- (aq) + NO (g) -> Cr 3+ (aq) + NO 3 - (aq) Answer: Cr 2 O 7 2- (aq) + 2NO (g) + 6H + (aq) -> 2Cr 3+ (aq) + 2NO 3 - (aq) + 3H 2 O (l)

7. In Basic Solution Follow the steps used in acidic solution THEN… Add OH - to match H + on same side. Add to other side as well Form H 2 O by adding OH - and H + Cancel H 2 O’s on both sides Check for element/charge balance

8. Example in Base Balance the following equation in basic solution: P 4 -> PH 3 + HPO 3 2- Answer P 4 + 2H 2 O + 4OH - -> 2PH 3 + 2HPO 3 2-

9. Galvanic Cell Concept Separating the oxidation and reduction half-reactions will make it so the energy in the electron transfer can be harnessed. A salt bridge or porous disk is needed to connect the half-reactions so ions can flow and electrons don’t build up on one side of the reaction (one beaker) **VIDEO?

10. Galvanic Cell Definition Device which chemical energy is changed to electrical energy. Oxidation occurs at the ANODE Reduction occurs at the CATHODE (cat gets fat = cathode gains electrons a.k.a. reduction) –An ox and a red cat (anode/oxidation, reduction/cathode)

11. Electrodes If there is an element (not ion) in either half-reaction, it is what that particular electrode is made of ***comes up later! When all reactants/products are in solution (aq) Pt or graphite can be used

12. Galvanic Cell Picture (Parts)

13. Cell Potential/Electromotive Force (EMF) Represented by E° cell Unit = volt (V) = 1 joule/coulomb Measured with a voltmeter (not completely accurate b/c of heat). A potentiometer is used instead where the maximum cell potential can be measured.

14. Standard Reduction Potentials If we can find the potential for each half- reaction (Table 18.1 pg. 829), we can determine the cell potential (E° cell ) Half-reaction manipulations (DO NOT MANIPULATE VOLTAGE): –One must be reversed (oxidation)…can reverse E so you have -E = -voltage… –Electrons lost must = electrons gained, so multiplication of reaction may be needed (DO NOT MULTIPLY VOLTAGE BY THIS NUMBER!) EQUATION (MEMORIZE) E° cell = E°(cathode) - E°(anode)

15. Table 18.1 Better oxidizing agents: easily reduced, LEFT side rxn. = largest, most positive standard reduction potential Better reducing agents: easily oxidized, RIGHT side = most negative standard reduction potential (aka most positive standard oxidizing potential) Example: Which is best reducing agent Cu +, F -, H -, H 2 O, I 2, K (find in products) Answer: K (-2.92) > H - (-2.23) > Cu + (0.16) > I 2 (1.20) > H 2 O (1.23) > F - (2.92)

16. Galvanic Cell Example Calculate the emf values (E° cell ) for the following Mg (s) + 2H + (aq) -> Mg 2+ (aq) +H 2(g) Answer: E° cell = +2.37V

17. In Galvanic Cells… When E° cell is positive, the reaction will run spontaneously. (last slide) If negative, it will run in the opposite direction (will NOT run spontaneously as written).

18. Cell Potential, Electrical Work, and Free Energy Spontaneous IF: –Positive cell potential –Negative Gibbs Free Energy (we will learn more about Gibbs Free Energy in a later chapter)

19. Line Notation Not required for AP Exam A double vertical line separates the anode on left and cathode on right –Represents a salt bridge or porous disk A single vertical line separates different phases Ex: anode Cd -> Cd 2+ + 2e - cathode Hg 2+ + 2e - -> Hg Line notation: Cd (s) I Cd 2+ (aq) II Hg 2+ (aq) I Hg

20. Galvanic Cell: Complete Description In half-reaction descriptions…FOUR items needed: 1.The cell potential - positive when E° cell = E°(cathode) - E°(anode) and the balanced cell reaction 2.The direction of electron flow, obtained by inspecting the half-reactions and using the direction that gives a positive E° cell. 3.Designation of the anode and cathode. 4.The nature of each electrode and the ions present in each compartment. A chemically inert conductor is required if none of the substances participating in the half-reaction is a conducting solid.

21. Example: Complete Description Describe a galvanic cell based on the two half- reactions below. Cu 2+ + 2e - -> CuE° = 0.34 V Cr 2 O 7 2- + 14H + + 6e - -> 2Cr 3+ + 7H 2 O E° = 1.33V 1.Balanced cell rxn: 3Cu (s) + Cr 2 O 7 2- (aq) + 14H + (aq) -> 3Cu 2+ (aq) + 2Cr 3+ (aq) + 7H 2 O (l) E° cell = 0.99V (needs to be positive) 2.1.33 (cathode) -0.34 (anode) means Cu needs to be reversed. Cu will be giving off e - which will travel from Cu (anode) to cathode (platinum electrode).

22. Continued… 3. Anode (copper metal electrode), cathode (platinum electrode) 4. The copper metal electrode (anode) will be in the Cu/Cu 2+ compartment while the platinum electrode (cathode) will be in the Cr 2 O 7 2- /Cr 3+ compartment

23. Changing Concentration In chemical reactions… Standard conditions = 1M for all When reactants are >1M, it will increase product concentration and will increase the cell potential When products are >1M, it will decrease the product concentration (oppose forward reaction), decreases cell potential Mg (s) + 2H + (aq) -> Mg 2+ (aq) +H 2(g)

24. A galvanic cell that has the same component on each side but at different concentrations Causes a cell potential Voltages are typically small Ex: a cell has on its left side a 0.20 M Cu 2+ solution and a 0.050 M Cu 2+ solution on the right side Concentration Cell

25. Reaction Quotient Represented by Q Q = concentration of products x [reactants] y *Solids and liquids are not included… Multiply concentrations of products/reactants by one another X and Y represent coefficients of each reactant and product respectively

26. H 2(g) + I 2(g) 2HI (g) [H 2 ] o = 0.81 M, [I 2 ] o = 0.44 M, [HI] o = 0.58 M Q = __[HI] 2 _ = (0.58) 2 = 0.94 [H 2 ] 1 [I 2 ] 1 (0.81)(0.44) Reaction Quotient Example (13.5)

27. Nernst Equation At 25°C…to find actual cell potential E cell = E° cell - 0.0591 log(Q) n n is number of moles of electrons transferred Q is the reaction quotient (see previous slide) E° cell is calculated using standard reduction potentials (learned earlier this chapter)

28. Nernst Equation Example Calculate E cell for a galvanic cell based on the following half-reactions at 25C. (Eq 1) FeO 4 2- + 8H + + 3e - -> Fe 3+ + 4H 2 O E° = +2.20 V (Eq 2) O 2 + 4H + + 4e - -> 2H 2 O E° = +1.23 V [FeO 4 2- ] = 2.0 X 10 -3 M, [Fe 3+ ] = 1.0 X 10 -3 M, [O 2 ] = 1.0 X 10 -5 M, [H + ] = 6.31 X 10 -6 M (see pg. 399 in study guide) E = 0.54 V

29. Cell potential is related to concentrations Electrodes can be used that are sensitive to specific ions They measure concentrations of specific ions which can have an effect on the cell potential Ion-Selective Electrodes

30. Batteries Groups of galvanic cells connected which add together to give the total battery potential Different types…

31. Lead Storage Battery Used in cars Lead = anode Lead coated with lead dioxide = cathode Electrodes are dipped in sulfuric acid

32. Dry Cell Batteries Calculators, electronic games, digital watches, portable audio players, etc. Acid version, alkaline version, silver cell, mercury cell, etc.

33. Fuel Cells Reactant constantly supplied to a galvanic cell Reaction is: 2H 2(g) + O 2(g)  2H 2 O (l)

34. Corrosion Metals corrode because they oxidize easily (spontaneously) –The metals’ reduction potentials are less positive than oxygen gas –When half reactions are combined, there is a positive E° value –Speed of oxidation varies –Some metals form a thin protective oxide coating, preventing further erosion

35. Iron Corrosion Steel corrosion must be prevented b/c it makes the backbone of most of our buildings, bridges, and cars Chemical composition of steel is not uniform –Places where iron is easily oxidized/reduced Ions and electrons migrate via moisture on steel’s surface, causing rust Moisture acts as a salt bridge between the anode/cathode regions

36. Corrosion Prevention A coating (paint/metal plating) is applied to protect the metal from oxygen and moisture –Chromium/tin used to plate steel with an oxide coating –Zinc used in galvanizing steel forms an oxide- carbonate coating. Oxygen reacts with the zinc (sacrificial coating) instead of the metal underneath Alloying: covers metals with a thin layer of stainless steel or other desirable alloy

37. More Corrosion Protection… Cathodic protection: for buried tanks/pipes attaches an active metal (like Mg) loses electrons, keeping iron from oxidizing –Ships’ hulls work this way: bars of titanium attached to the steel hull keep the hull from being oxidized

39. Electrolysis Electrolytic cells use electrical energy to produce a chemical change (requires energy) –Electrolysis forces a current through a cell to produce a chemical change for which the cell potential is negative –Causes a nonspontaneous rxn to occur Charges batteries, produces aluminum metal, plates objects with chrome Switches the anode/cathode, flows in the opposite direction as the galvanic cell

40. Stoichiometry in Electrochemistry Can measure how much chemical change occurs with the flow of a given current for a specific amount of time We add current and time to the chemical equation and can use the mole to mole ratio to do problems Important Conversion Factor for Ampere (amp): 1 A = coulomb/second 1 Faraday: 1 mole e - = 96,486 C

41. Stoichiometry Example How many grams of copper can be reduced by applying a 3.00 A current for 16.2 min to a solution containing Cu 2+ ions? Answer: 0.96 g Cu

42. Plating Order Plating means depositing neutral metal on the electrode by reducing the metal ions in solution More positive (higher) E° value will plate out quickest Ex: pg. 851

43. Electrolysis Used Commercially… Production of Aluminum Electrorefining of Metals Metal Plating Electrolysis of Sodium Chloride

44. Production of Aluminum Aluminum used to not be able to be produced by humans (very expensive) Hall-Heroult process uses molten cryolite (Na 3 AlF 6 ) as a solvent for aluminum oxide Made aluminum easy (cheap) to produce Production of Al uses 5% of the total electricity used in the U.S.

45. Electrorefining of Metals Purifies metals like Cu, Fe, Zn, Ag, Au, Pt

46. Metal Plating Metals that corrode easily can be plated with a thin coating of a metal that is less likely to corrode Ex: “tin” cans = steel cans with thin tin coating, bumpers in cars are steel coated in chrome Object needing coating serves as the cathode in ions of the coating metal

47. Electrolysis of Sodium Chloride How sodium metal is primarily produced NaCl mixed with CaCl 2 to lower mp Downs cell used for electrolysis Pg. 856