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Chemistry 400: General Chemistry Miller Fall 2015 Final Exam Part Deux December 14, 2015 Name: _____________________________________ Approximately 150 points Please answer each of the following questions to the best of your ability. If you wish to receive partial credit, please show your work. For all ionic species, please show the charge on each ion to receive full credit. Good luck! All Lewis Structures must include all valence electrons. Element Electronegativity F 4.0 O 3.5 Cl 3.0 N 3.0 S 2.8 Br 2.8 C 2.5 H 2.1 I. Nomenclature (4 points each, spelling counts) If the name is given, please give the formula. If the formula is given, please give the name. A. NaNO2 B. CH3COOH C. Fe2S3 D. N2O3 E. CuSO4•5H2O II. Drawing 1. For CF4: (i) Draw the correct Lewis structure. (4 points) (ii) Draw the correct electron geometry and (iii) Draw dipoles. (6 points) (iv) What is the bond angle around the oxygen atom? Use a <, >, or = symbol as part of your answer. (2 points) (v) Is this molecule polar or nonpolar? (4 points) 1 2. If MgSO4·xH2O, is heated to 250°C, all the waters of hydration evaporate. On heating a 0.4000 g sample of the hydrate MgSO4·xH2O, 0.1953 g of MgSO4 remains. How many molecules of water occur per formula unit of MgSO4? In other words, solve for x. For this problem, your show your work if you'd like to receive any credit. (10 points) 3. Will the pH of 0.100 M HCl or 0.100 M CH3COOH be lower? Explain your choice with specific reference to the types of acids involved and their percents of ionization. (8 points) 2 4. There is a relationship between intermolecular forces, temperature, and the temperatures at which the solid, liquid, and gas phases exist. Each blank will be filled in with one of the words in bold. You may use each word more than once. (1 point per blank) A. ______________________ is proportional to the average kinetic energy. B. At low temperatures, materials are in the ______________________ phase. In this phase, ______________________ are strong enough to keep the particles in the material (atoms, ions, or molecules) close enough such that the particles have 3dimensional order. C. At high temperatures, materials are in the ______________________ phase. In this phase, no matter how strong the ______________________ are, the energy supplied by the high temperatures is enough to overcome the ______________________ . D. At intermediate temperatures, materials are in the ______________________ phase. The exact temperatures of the phase transitions depend upon the strength of ______________________ . 3 5. Calculate the change in the enthalpy of reaction, ΔHrxn, for the following combustion reaction two ways: C2H6(g) + 3.5 O2(g) → 2 CO2(g) + 3 H2O(g) A. Using bond dissociation energies. (14 points) B. Using standard changes in the enthalpy of formation, ΔHf°. (6 points) ΔHf°(C2H6) = -84 kJ/mol 4 6. In the experiment below, the electron shooter shoots one electron at a time through the double slits at the screen. Draw the pattern that the electrons make on the screen after a long period of time in which many electrons have been shot at the screen but the electrons have been shot at the screen one at a time. A. No laser (no observer) to determine which slit the electron goes through (4 points) electron shooter: shoots one electron at a time double slit screen B. A laser (an observer) is present to determine which slit the electron goes through (4 points) electron shooter: shoots one electron at a time double slit screen 7. Define an orbital. Make sure to use the words volume and probability in your definition (6 points) 5 8. Write a paragraph about the relationship between the 3d and 4s orbitals in terms of energy. For what elements is 3d lower in energy? For what elements is 4s lower in energy? To answer this question, use the electron configurations (with or without noble gas configurations) of a potassium atom, a potassium ion, a titanium atom, and a doubly charged titanium ion. (12 points) electron configurations: K atom K+ ion Ti atom 6 Ti2+ ion 9. A. Draw the best Lewis Structure for NH3 and the best Lewis structure for ClH3, named chlorine trihydride. All atoms will have zero formal charge. For each Lewis structure, calculate the difference in electronegativity for each type of bond and draw dipoles for each bond (if they exist). (8 points) B. Determine the dominant type of intermolecular force for each molecules above and suggest a reason why they are different (even though the differences in electronegativity of each type of bond might suggest otherwise). (8 points) 10. The nonvolatile, nonelectrolyte aspirin, C9H8O4, is soluble in diethyl ether CH3CH2OCH2CH3. How many grams of aspirin are needed to generate an osmotic pressure of 7.72 atm when dissolved in 190 ml of a diethyl ether solution at 298 K. (8 points) 7 11. Shown below is the distribution of velocities for hydrogen (H) atoms at 300 K. Place a line on this graph that approximately represents the distribution of velocities for helium (He) atoms at 300 K. (4 points) Fraction of Atoms x 10^5 100 90 80 70 60 50 40 30 20 10 0 0 1000 2000 3000 4000 5000 Velocity (m/s) 8 6000 12. From calculations in class, it was clear that the wavelength of an electron was approximately the same size as an atom of hydrogen, r = 0.05 nm. A. Calculate the wavelength of an electron traveling at 2.2 x 106 m/s. (6 points) B. Are there any other particles or atoms for which the wavelength is comparable to the size of the atom? If so, please show through calculations. If not, then please explain why. (10 points) 9 10 Chemistry 400 Conversions and Equations 1 L = 1.057 qt 453.6 g = 1 lb ºF = 1.8 º C + 32 1 m = 39.37 in 1 yd = 36 in = 3 ft Na = 6.02 x 1023 1 gal = 4 qt 1 lb = 16 oz 1 calorie = 4.184 J specific heat of water = 4.184 J/g °C q = Energy = (mass) (Csp) (ΔT) average atomic mass = PO2 = (% O2) PT λ= € E = hν 2 % yield = actual/theoretical × 100% PT = P1 + P2 + P3 + … λ= $ 1 1' E = −2.18 ×10 J Z && 2 − 2 )) € % n2 n1€ ( −18 € (mass isotope 1)(%) + (mass isotope 2)(%) + (mass isotope 3)(%) 100% P1 V1 / T1 = P2 V2 / T2 c ν 1 qt = 4 cups 1 mile = 5280 feet 1 atm = 101.3 kPa = 1.013 bar = 14.7 psi For gases: standard T = 273.15 K, P = 1 atm h KE = ½ mv2 mv h = 6.626×10–34 J•s c = 3.00×108 m/s€ 1 L•atm = 101.325 J PV=nRT ΔxmΔv = w = -P ΔV h 4π Mass of electron = 9.1×10–31 kg Kw = [H3O+] [OH–] = 1.0×10–14 [H+] [OH–] = 1.0×10–14 C 1V 1 = C 2V 2 R=0.08206 L•atm/mol•K = 8.314 J/mol•K Specific heat of ice: 2.09 J/g•°C Specific heat of water: 4.184 J/g•°C Specific heat of steam: 2.03 J/g•°C Heat of fusion of H2O = Δ Hfus = 6.02 kJ/mol Heat of vaporization of H2O = Δ Hvap = 40.7 kJ/mol € ΔT = m i K π =iMRT € € " P % −ΔHvap " 1 1 % ln$ 2 ' = $ − ' R # T2 T1 & # P1 & € € 1 1 ppm = 1 x 10-6 g/mL € 1 ppb = 1 x 10-9 g/mL Chemistry 400 Conversions and Equations H 436 C 414 347 Average Single Bond Dissociation Energy (in kJ/mol) N O F Si P S Cl Br I 389 464 565 323 322 368 431 364 297 305 360 485 301 272 339 276 213 163 222 272 200 243 159 142 190 452 335 203 201 234 159 565 490 327 253 237 278 226 293 464 310 234 201 326 184 266 253 218 243 216 208 193 175 151 H C N O F Si P S Cl Br I Comparison of Average Single, Double and Triple Bond Energies (in kJ/mol) Bond Type Single Bond Double Bond Triple Bond C–C 347 611 837 N–N 163 418 946 O–O 142 498 C–N 305 615 891 C–O 360 736* 1072 C–Cl 339 651 *For CO2, the C=O bond is 799 kJ/mol Boiling Point Elevation and Freezing Point Depression Constants Solvent Formula Kb (°C/m) Kf (°C/m) Water H 2O 0.512 –1.86 Ethanol CH3CH2OH 1.22 –1.99 Chloroform CHCl3 3.63 –4.70 Benzene C 6H 6 2.53 –5.12 Diethyl ether CH3CH2OCH2CH3 2.02 –1.79 Element F O Cl N S Br C H 2 Electronegativity 4.0 3.5 3.0 3.0 2.8 2.8 2.5 2.1 Chemistry 400 Conversions and Equations Material Ag(s) Ag+(aq) Al(s)€ Al3+(aq) Al2O3(s) AlCl3(aq) AlCl3(s) Br(g) Br2(g) Br2(l) C(g) C(s, dia) C(s, gr) C2H4(g) C2H4O(g) C2H5OH(l) C6H12O6(s) C3H8(g) CH3CH2CH2CH3(l) Ca(g) Ca(OH)2(aq) Ca(OH)2(s) Ca(s) Ca2+(aq) Ca2+(g) CaCl2(s) CaCO3(s) CaF2(s) CaO(s) CH3OH(g) CH4(g) CHCl3(l) Cl–(aq) Cl(g) Cl2(g) CO(g) CO2(g) Cu(s) Cu2+(aq) Fe2O3(s) Fe3O4(s) H(g) H+(aq) H2(g) H2O(g) H2O(l) H2O(s) H2O2(aq) Δ Hf°(kJ/mol) 0 105.79 0 –538.4 –1675.7 –1039.7 –704.2 111.9 30.9 0 716.7 1.88 0 52.4 –166.2 –277.6 –1273.3 –107.85 –147.3 177.8 –1003 –985.2 0 –542.8 1934.1 –795.4 –1207.6 –1228.0 –634.9 –201.0 –74.6 –134 –167.1 121.3 0 –110.5 –393.5 0 64.9 –824 -1118 218 0 0 –241.8 –285.8 –291.8 –191.2 3 Material H2O2(l) H3O+(aq) HBr(g) € HCl(aq) HCl(g) H2SO4(l) I2(g) I2(s) Mg2+(aq) MgCl2(aq) N(g) N2(g) N2H4(l) N2O(g) N2O4(g) Na(s) Na+(aq) Na2SO4(s) NaCl(aq) NaCl(s) NaOH(aq) NH3(aq) NH3(g) NH4+(aq) NH4Cl(aq) NH4Cl(s) NH4NO3(aq) NH4NO3(s) NI3(s) NO(g) NO2(g) O(g) O2(g) O3(g) OH–(aq) SO2Cl2(g) SO2(g) SO42–(aq) Zn(s) Zn2+(aq) HgO(s) Hg(l) Δ Hf°(kJ/mol) –187.78 –285.8 –36.3 –167.2 –92.3 –814 62.42 0 –467.0 –801.2 472.7 0 50.6 81.6 11.1 0 –240.34 –1387.1 –407.2 –411.2 –470.1 –80.29 –45.9 –133.26 –299.66 –314.43 –339.9 –365.6 192 91.3 33.2 249.2 0 142.7 –230.02 –364 –296.8 –909.3 0 –153.39 –90.8 0 Chemistry 400 Conversions and Equations Unit of Concentration Symbol Formula Mass/mass percent % (w/w) %(m/m) = grams of solute x 100 grams of solution Mass/volume percent % (w/v) % (m/v) = grams of solute x 100 mL of solution Volume/volume percent % (v/v) % (v/v) = Molarity M M = moles of solute L of solution parts per million ppm ppm = grams of solute x 106 grams of solution mL of solute x 100 mL of solution 4