Chemistry Electrolytic Cells and Electroplating Advanced Placement 2014 2 Ar Kr 10 18 36 54 86 Xe He Ne Rn (222) 83.80 4.0026 20.179 39.948 131.29 ) 3 I 9 F Cl At Br 17 35 53 85 Lr 71 Lu 10 (210 19.00 79.90 (260) 35.453 126.91 174.97 ) 8 S O Te 16 34 52 84 Se Po 70 Yb No 102 (209 16.00 32.06 78.96 (259) 127.60 173.04 7 P N Bi 15 33 51 83 As Sb 69 Md Tm 101 74.92 (258) 14.007 30.974 121.75 208.98 168.93 1 2 0 6 C Si 8.71 14 32 50 82 Sn Pb Er Ge 68 Fm 10 28.09 72.59 207. (257) 12.01 11 167.26 11 §Not yet named 5 B Tl In Al 4.82 13 31 49 81 67 99 Es Ga Ho (252) 26.98 69.72 10.8 11 204.38 164.93 2 § 30 48 80 Zn Cd Hg Cf 66 98 Dy 11 12.41 (277) 65.39 (251) 1 200.59 162.50 § 29 47 79 Ag Au Cu 65 97 Tb Bk 111 (272) 63.55 (247) 107.87 196.97 158.93 ) 0 § Ni Pt 28 46 78 Pd 64 96 Gd 11 Cm (269 (247) 58.69 106.42 195.08 157.25 2 Ir 27 45 77 Mt Co Rh 63 95 Eu 109 Am ble of the Elements (266) 58.93 192. (243) 102.91 151.97 Ta 1 2 4 26 44 76 Fe Hs Ru Os 62 94 Pu 108 Sm (265) 55.85 101. 190. (244) 150. Tc 25 43 75 61 93 Bh Re Np Mn Pm 107 (98) Periodic (262) (145) 54.938 186.21 237.05 U W Cr 24 42 74 Sg 60 92 Mo Nd 106 (263) 52.00 93.94 183.85 144.24 238.03 ) 4 1 5 V Pr Ta 23 41 73 59 91 Pa Nb Db 10 (262 50.9 92.9 180.95 140.91 231.04 ) 0 2 4 Ti Zr Hf Rf 22 40 72 58 90 Th Ce 10 (261 47.9 91.2 178.49 140.12 232.04 : : Y 21 39 57 89 La Sc Ac * † 44.96 88.91 138.91 227.03 0 8 2 4 Sr 12 20 38 56 88 Be Ba Ca Ra Mg 9.012 24.3 40.0 87.6 137.33 226.02 ) 1 †Actinide Series 1 3 K H Li Fr 11 19 37 55 87 *Lanthanide Series Cs Na Rb (223 6.94 22.99 39.10 85.47 1.0079 132.91 ADVANCEDADVANCED PLACEMENTPLACEMENT CHEMISTRYCHEMISTRY EQUATIONSEQUATIONS ANDAND CONSTANTSCONSTANTS ThroughoutThroughout thethe testtest thethe followingfollowing symbolssymbols havehave ththee definitionsdefinitions specifiedspecified unlessunless otherwiseotherwise noted.noted. L,L, mLmL == liter(s),liter(s), milliliter(s)milliliter(s) mmmm HgHg == millimetersmillimeters ofof mercurymercury gg == gram(s)gram(s) J,J, kJkJ == joule(s),joule(s), kilojoule(s)kilojoule(s) nmnm == nanometer(s)nanometer(s) VV == volt(s)volt(s) atmatm == atmosphere(s)atmosphere(s) molmol == mole(s)mole(s) ATOMICATOMIC STRUCTURESTRUCTURE EE == energyenergy EE == hhνν νν == frequencyfrequency cc == λνλν λλ == wavelengthwavelength Planck’sPlanck’s constant,constant, hh == 6.6266.626 × × 1010−−3434 JJ s s SpeedSpeed ofof light,light, cc == 2.9982.998 ×× 101088 msms−−11 Avogadro’sAvogadro’s numbernumber == 6.0226.022 ×× 10102323 molmol−−11 ElectronElectron charge,charge, e e == −−1.6021.602 ××1010−−1919 coulombcoulomb EQUILIBRIUMEQUILIBRIUM [C][C]cdcd [D] [D] KKcc == ,, wherewhere aa AA ++ bb BB RR cc CC ++ dd DD EquilibriumEquilibrium Constants Constants [A][A]abab [B] [B] cdcd KKcc (molar(molar concentrations)concentrations) ((PPPPCC )( )(DD ) ) KKpp == abab KKpp (gas(gas pressures)pressures) ((PPPPABAB )( )( ) ) KKaa (weak(weak acid)acid) [H[H+-+- ][A ][A ] ] KK == K (weak base) aa [HA][HA] Kbb (weak base) KK (water)(water) [OH-+-+ ][HB ] ww KK == [OH ][HB ] bb [B][B] ++ −− −−1414 KKww == [H[H ][OH][OH ]] == 1.01.0 ××1010 atat 2525°°CC == KKaa ×× KKbb pHpH == −−log[Hlog[H++]] , , pOHpOH == −−log[OHlog[OH−−]] 1414 == pHpH ++ pOHpOH -- pHpH == ppKK ++ loglog[A[A ] ] aa [HA][HA] ppKKaa == −−loglogKKaa,, ppKKbb == −−loglogKKbb KINETICSKINETICS ln[A]ln[A] −− ln[A]ln[A] == −−ktkt tt 00 kk == raterate constantconstant 1111 tt == timetime -- == ktkt AAAA [][]tt [][]00 tt½½ == half-lifehalf-life 0.6930.693 tt == ½½ kk GASES, LIQUIDS, AND SOLUTIONS P = pressure V = volume PV = nRT T = temperature moles A n = number of moles PA = Ptotal × XA, where XA = total moles m = mass M = molar mass Ptotal = PA + PB + PC + . D = density n = m KE = kinetic energy M Ã = velocity K = °C + 273 A = absorbance a = molarabsorptivity m D = b = path length V c = concentration 1 KE per molecule = mv2 2 Gas constant, R = 8.314 J mol--11 K Molarity, M = moles of solute per liter of solution --11 = 0.08206 L atm mol K --11 A = abc = 62.36 L torr mol K 1atm= 760 mm Hg = 760 torr STP= 0.00D C and 1.000 atm THERMOCHEMISTRY/ ELECTROCHEMISTRY q heat = m = mass q= mcD T c = specific heat capacity T temperature DSSDD= ÂÂ products- S D reactants = SD = standard entropy DDHHDD= ÂÂproducts- D HfD reactants f HD = standard enthalpy GD standard free energy DDGGDD= ÂÂf products- D GfD reactants = n = number of moles D DDGD = HDD - TS D E = standard reduction potential I = current (amperes) = -RTln K q = charge (coulombs) t = time (seconds) = -nFED q Faraday’s constant , F = 96,485 coulombs per mole I ϭ t of electrons 1 joule 1volt = 1coulomb Electrochemistry Electrolytic Cells and Electroplating What I Absolutely Have to Know to Survive the AP Exam The following might indicate the question deals with electrochemical processes: E°cell; cell potential; reduction or oxidation; anode; cathode; salt bridge; electron flow; voltage; electromotive force; galvanic/voltaic; electrode; battery; current; amps; time; grams (mass); plate/deposit; electroplating; identity of metal; coulombs of charge ELECTROCHEMICAL TERMS Electrochemistry the study of the interchange of chemical and electrical energy OIL RIG Oxidation Is Loss, Reduction Is Gain (of electrons) LEO the lion says GER Lose Electrons in Oxidation; Gain Electrons in Reduction Oxidation the loss of electrons, increase in charge Reduction the gain of electrons, reduction of charge Oxidation number the assigned charge on an atom ELECTROCHEMICAL CELLS Electrochemical Cells: A Comparison Electrodes made Battery its cell Galvanic (voltaic) spontaneous − Is separated into 2 from metals (inert Pt potential drives the cells oxidation-reduction half-cells or C if ion to ion or reaction and thus the reaction gas) e− Battery charger − non-spontaneous requires an external Usually occurs in a Usually inert Electrolytic cells oxidation-reduction energy source to single container electrodes reaction drive the reaction and e− Copyright © 2014 National Math + Science Initiative, Dallas, Texas. All Rights Reserved. Visit us online at www.nms.org Page 1 Electrochemistry – Electrolytic Cells and Electroplating Electrolysis and Non-spontaneous Cells The Electrolytic Cell: What is what and what to know An electrolytic cell is an electrochemical cell in which the reaction occurs only after adding electrical energy; thus it is not thermodynamically favorable – it requires a driving force Electrolytic reactions often occur in a single container; not 2 separate ½ cells like Voltaic cells The calculated E° for an electrolytic cell is negative (i.e. that is the amount of potential required to drive the reaction). This driving force causes the electrons to travel from the positive electrode to the negative electrode, the exact opposite of what you would expect. Remember: EPA – Electrolytic Positive Anode Anode − the electrode where oxidation occurs. Cathode − the electrode where reduction occurs. Electron Flow − From Anode To CAThode § Electrolytic reactions typically occur aqueous solutions (or in molten liquids) § If there is no water present, you have a pure molten ionic compound thus § the cation will be reduced § the anion will be oxidized § If water is present, you have an aqueous solution of the ionic compound, thus § You must decide which species is being oxidized and reduced; the ions or the water! § No alkali or alkaline earth metal can be reduced in an aqueous solution - water is more easily reduced. § Polyatomic ions are typically NOT oxidized in an aqueous solution - water is more easily oxidized. § When it comes to water, be familiar with the following § REDUCTION OF WATER: − − 2 H2O() + 2 e → H2(g) + 2 OH E° = −0.83 V § OXIDATION OF WATER: + − 2 H2O()→ O2(g) + 4 H + 4 e E° = −1.23 V Copyright © 2014 National Math + Science Initiative, Dallas, Texas. All Rights Reserved. Visit us online at www.nms.org Page 2 Electrochemistry – Electrolytic Cells and Electroplating Electrolysis: What to Do and How to Do it! An electric current is applied to a 1.0 M KI solution. The possible reduction half reactions are listed in the table. Reduction Half Reactions E° (V) − 2 H2O() + 2 e− → H2(g) + 2 OH −0.83 V + − O2(g) + 4 H + 4 e → 2 H2O() +1.23 V − − I2(g) + 2 e → 2 I +0.53 V K+ + e− → K(s) −2.92 V (A) Write the balanced half-reaction for the reaction that takes place at the anode. (B) Write the balanced overall reaction for the reaction that takes place. (C) Which reaction takes place at the cathode? + − First realize the solution contains 3 important species: K , I , and H2O. From those 3 species, you must decide which is being oxidized and which is being reduced. Where to start – Water can be both oxidized and reduced in an aqueous solution. The I− cannot be further reduced (it has a 1– charge) The K+ cannot be further oxidixed (it has a 1+ charge) − For OXIDATION: either I or H2O will be OXIDIZED… Write the oxidation half-reactions – Remember, the more positive oxidation potential will be oxidized.