exam 3 revised chem 141 fall 2015(3)

Chemistry 141 Name Martin Larter Exam 3A November 7. 2015 Multiple Choice (30 points) Page 5 (14 points) Page 6 (21 points) Page 7 (22 points) Page 8 (15 points) Page 9 (12 points) Total (114 points) Chemistry Formulas and Constants Kinetic energy = ½ mv 2 w = -P  V E = hν Δ x Δmv≥ h 4 π Rate 1 Rate 2 = √ MW 2 MW 1  G o = -nFE o  G = - RTlnK w=dxF E = mc 2 P total = P 1 +P 2 +P 3 +… u = (3RT/MW) ½ H  =E  M 1 V 1 = M 2 V 2 u = (3RT/MW) ½  G =  H - T  S P total = P 1 + P 2 + P 3 + … PV = nRT Rate  (MW) -½ ( P + n 2 a V 2 ) ( V − nb ) = nRT c = λν λ deBroglie = h mv  = iMRT P total = P 1 + P 2 + P 3 + … ln ( [ A ] [ A ] o ) =− kt ln ( k 2 k 1 ) = ( E a R ) [ ( 1 T 1 ) − ( 1 T 2 ) ] E = E o − ( 0.0592 n ) log Q ln ( P 2 P 1 ) = ( ΔH vap R ) [ ( 1 T 1 ) − ( 1 T 2 ) ] ν = c λ 1 λ =− R ( 1 n i 2 − 1 n f 2 ) λ deBroglie = h mv ΔxΔ mv ≥ h 4 π Constants F = 9.65 x 10 4 C h = 6.626 x 10 -34 J sec c= 2.9979 x 10 8 m/sec mass electron = 9.109 x 10 -31 kg R E = 2.18 x 10 -18 J e = 1.602 x 10 -19 C R = 0.0821 L atm/mol K = 62.4 L torr/mol K = 8.31 kJ/mol K 1

Grossmont College Periodic Table IA VIIA NOBLE GASES 1 H 1.008 IIA IIIA IVA VA VIA 1 H 1.008 2 He 4.002 3 Li 6.941 4 Be 9.012 5 B 10.81 6 C 12.01 7 N 14.01 8 O 16.00 9 F 19.00 10 Ne 20.18 11 Na 23.00 12 Mg 24.30 IIIB IVB VB VIB VIIB VIII VIII VIII IB IIB 13 Al 27.00 14 Si 28.09 15 P 30.97 16 S 32.06 17 Cl 35.45 18 Ar 39.95 19 K 39.10 20 Ca 40.08 21 Sc 44.96 22 Ti 47.90 23 V 50.94 24 Cr 52.00 25 Mn 54.94 26 Fe 55.85 27 Co 58.93 28 Ni 58.70 29 Cu 63.55 30 Zn 65.38 31 Ga 69.72 32 Ge 72.59 33 As 74.92 34 Se 78.96 35 Br 79.90 36 Kr 83.80 37 Rb 85.47 38 Sr 87.62 39 Y 88.91 40 Zr 91.22 41 Nb 92.91 42 Mo 95.94 43 Tc (99) 44 Ru 101.1 45 Rh 102.9 46 Pd 106.4 47 Ag 107.9 48 Cd 112.4 49 In 114.8 50 Sn 118.7 51 Sb 121.8 52 Te 127.6 53 I 126.9 54 Xe 131.3 55 Cs 132.9 56 Ba 137.3 57 La 138.9 72 Hf 178.5 73 Ta 180.9 74 W 183.9 75 Re 186.2 76 Os 190.2 77 Ir 192.2 78 Pt 195.1 79 Au 197.0 80 Hg 200.6 81 Tl 204.4 82 Pb 207.2 83 Bi 209.0 84 Po (209) 85 At (210) 86 Rn (222) 87 Fr (223) 88 Ra 226.0 89 Ac 227.0 104 Rf (261) 105 Db (262) 106 Sg (263) 107 Bh (262) 108 Hs (265) 109 Mt (266) 110 ?? (269) 58 Ce 140.1 59 Pr 140.9 60 Nd 144.2 61 Pm (147) 62 Sm 150.4 63 Eu 152.0 64 Gd 157.3 65 Tb 158.9 66 Dy 162.5 67 Ho 164.9 68 Er 167.3 69 Tm 168.9 70 Yb 173.0 71 Lu 175.0 90 Th 232.0 91 Pa 231.0 92 U 238.0 93 Np (237) 94 Pu (244) 95 Am (243) 96 Cm (247) 97 Bk (247) 98 Cf (251) 99 Es (252) 100 Fm (257) 101 Md (258) 102 No (259) 103 Lr (260) Lanthanide series 2

Multiple Choice (30 points) 1. Which statement is always true according to VSEPR theory? a. The shape of a molecule is determined by the polarity of its bonds. b. The shape of a molecule is determined by the repulsions among all electron groups on the central atom (or interior atoms if there is more than one). c. The shape of a molecule is determined only by repulsions among nonbonding electron groups. d. The shape of a molecule is determined only be repulsions among bonding electron groups. e. none of the above 2. Arrange the ions P 3- , S 2- , Ca 2+ , K + , and Cl - in order of increasing ionic radius, starting with the smallest first. a. Ca 2+ , K + , Cl - , S 2- , P 3- b. P 3- , S 2- , Ca 2+ , Cl - , K + c. P 3- , Ca 2+ , S 2- , K + , Cl - d. K + , Ca 2+ , Cl - , S 2- , P 3- e. P 3- , S 2- , Cl - , K + , Ca 2+ 3. Which one of the following ground-state orbital diagrams only violates Hund's rule ? a. b. c. d. e. 4. Which bond should have the longest length? a. N=N b. N-N c. N≡N d. All three bond lengths should be about the same. e. Impossible to determine from the data given 5. Which of these statements correctly describes the use of formal charge in choosing between possible Lewis structures to describe a molecule? I. Formal charge on all the atoms should be zero, or at least the smallest possible value. II. The sum of the formal charges on all the atoms should equal the charge on the molecule or ion. III. The formal charge on an atom is not affected by the electronegativity of the atom. IV. In an anion containing nitrogen and oxygen, the Lewis structures with a negative charge on the nitrogen will contribute the most to the description of the bonding. a. III and IV only b. I and III only c. II and III only d. II and IV only e. I and II only 6. In general, resonance ________ electrons and ________ molecules. a. delocalizes; destabilizes b. delocalizes; stabilizes c. destabilizes; destabilizes d. localizes; stabilizes e. localizes; destabilizes 7. Which type of molecular orbital is used to describe electron density building up above and below the internuclear axis to form a bond? a. π b.  * c.  * d.  e. s 4

8. The reason Fe 3+ has a smaller ionic radius than Fe 2+ is because a. Fe 2+ has a low electron affinity b. n is smaller for Fe 2+ c. Fe 3+ contains more protons d. Fe 3+ has a higher Z eff e. Fe 3+ has a higher ionization energy 9. List the elements Cs, Ca, Ne, Na, Ar in order of increasing first ionization energy. a. Ne > Na > Cs > Ca > Ar b. Ne > Ar > Na > Cs > Ca c. Ar > Ca > Cs > Na > Ne d. Ne > Ar > Ca > Na > Cs e. None of the above 10. Which of the following atoms is diamagnetic? a. B b. Mg c. C d. Na e. O 11. The greater the electronegativity difference between two bonded atoms, the a. Greater the bond order. b. Greater the ionic character of the bond. c. More unstable the bond. d. Greater the covalent character of the bond. e. None of the above 12. Which reaction below represents the second electron affinity of S? a. S 2 ⁻ (g) → S ⁻ (g) + e ⁻ b. S(g) + e ⁻ → S ⁻ (g) c. S ⁻ (g) + e ⁻ → S 2 ⁻ (g) d. S(g) → S ⁺ (g) + e ⁻ e. S ⁻ (g) → S(g) + e ⁻ 13. Based on the indicated electronegativity’s, arrange the following in order of increasing ionic character: CsBr, LaBr 3 , PBr 3 , MgBr 2 . a. CsBr, LaBr 3 , MgBr 2 , PBr 3 b. CsBr, MgBr 2 , PBr 3 , LaBr 3 c. PBr 3 , MgBr 2 , LaBr 3 , CsBr d. PBr 3 , LaBr 3 , MgBr 2 , CsBr e. Not enough information provided 14. Which of the following is not true? a. A sp 3 hybrid orbital may form a sigma bond by overlap with an orbital on another atom. b. A sp 3 hybrid orbital may hold a lone pair of electrons. c. A sp 3 hybrid orbital may form a pi bond by overlap with an orbital on another atom. d. The sp 3 hybrid orbitals are degenerate. e. All of the above are true. 15. A molecule containing a central atom with sp 3 d 2 hybridization has a(n) ________ electron geometry. a. tetrahedral b. bent c. trigonal bipyramidal d. square planar e. octahedral 16. In quantum mechanics, an atomic orbital ________ 5 element electronegativity Br 2.8 P 2.1 Mg 1.2 La 1.0 Cs 0.7

1. (9 points) Write the electron configurations for the following atoms or ions as predicted by the periodic table a. Si (complete configuration) _________________________________ b. Platinum (Pt) (shorthand configuration) ______________________________ c. Cobalt(II) ion (shorthand configuration)______________________________ 2. (6 points) Both vanadium and its 3+ ion are paramagnetic; Use words, electron configurations and orbital diagrams to explain why this is so. 3. (8 points) As you move across a row on the periodic table, the atomic radius decreases. As you move down a column on the periodic table, the atomic radius increases. However, in both cases (across a row and down a column) the number of protons increases. Provide a complete explanation ( and define your terms ) explaining how while the number of protons increases in both cases, the change in atomic radius is different. A complete response will address both trends. 4. (4 points) Why is SF 4 a stable molecule, while the molecule OF 4 does not exist? 7

5. (18 points) Write the best Lewis Electron Dot Structures for the following molecules or ions (Central atom is listed first). Tell the orbital and molecular geometry for each molecule/ion. Show formal charges for all non-zero charges. If resonance structures exist, show them. NS 2 Cl orbital geometry_________________ molecular geometry_______________ Bond angle________________ Hybridization ______________ Polar or Nonpolar_______________ SeOF 2 orbital geometry_________________ molecular geometry_______________ Bond angle________________ Hybridization ______________ Polar or Nonpolar_______________ XeF 2 orbital geometry_________________ molecular geometry_______________ Bond angle________________ Hybridization ______________ Polar or Nonpolar_______________ 6. (4 points) Compare and contrast valence bond theory and molecular orbital theory. 8

8. (8 points) Look at the compound pictured below. Explain the bonding in terms of valence bond theory. That is show the atomic orbitals on the Br atom, describe any electron promotion and hybridization necessary, and label the orbitals involved in both sigma and pi bonding as well as the orbital holding the lone pair of electrons on Br. You to draw a 3D representation of the orbitals. O Xe O 9. (4 points) NH 3 is a polar molecule but BH 3 is not. Explain this observation in terms of the structures of the two molecules. 10

11 Molecule/ion_________________________ _____ Molecule/ion______________________ _______ Molecule/ion________________________ ______ Molecule/ion_______________________ _______

12. (6 points) Seven ice cubes having a combined mass of 185 g are taken from the freezer at -15.0 o C and are added to 591 mL (20 fl. oz.) of soda at 19.0 o C in an insulated cup. Then thermal equilibrium is reached the solution is 0.00 o C. Assuming no heat is exchanged with the cup or rest of the environment, what mass of ice remains? (Assume the density and heat capacity of soda are the same that of water) 13. (8 points) A new substance developed in a laboratory has the following properties: normal melting 83.7 C; normal ̊ boiling point, 177 C; triple point 200. Torr and 38.6 C. Sketch the approximate phase diagram and label the solid, ̊ ̊ liquid, and gaseous phases. Based on your drawing is the solid or liquid denser? 13 Specific heat of ice 2.06 J/g o C 37.1 J/mol o C Specific heat of water 4.184 J/g o C 75.4 J/mol o C Specific heat of steam 2.0 J/g o C 36 J/mol o C Heat of fusion 333 J/g 6.01 kJ/mol Heat of vaporization 2260 J/g 40.7 kJ/mol

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