Exam%203%20F2008%20(PRACTICE) - Name (PRINT):

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Unformatted text preview: Name (PRINT): ________________________________ Chem 102A Examination 3 Fall, 2008 I pledge on my honor that I have neither given nor received improper aid on this examination. ______________________________ (Signature) Directions: 1. No books or notes may be using during this test. There should be no data stored on your calculator for use on the exam. This includes but is not limited to periodic tables and chemical equations and formulas. 2. There are a total of 9 sections (some with several parts) for a sum of 100 points. 3. Be certain that, wherever applicable, you show your work. Problems in Sections 4, 6 and 9 that are answered without showing detailed work will not receive credit. 4. Please note: there are no intentionally misleading questions on this test. Each problem should be taken at its face value. You must, however, be certain to read each problem carefully! 5. If you have any questions please ask your professor. DO NOT REMOVE THIS PAGE FROM YOUR EXAM!!!!! Please be sure that you are officially enrolled in the section named below: Section 2– Hanusa Exam Score________ 2 SECTION 1 (12 PTS. TOTAL; 2 PTS. EACH) Circle either True or False for each answer. (1) According to the VSEPR theory, BCl3 and AlCl3 molecules are expected to have the same molecular geometry. True False (2) Of the species NO2, NO and N2, only N2 obeys to the octet rule. True False (3) Atomic radii decrease across a row from left to right in the peri- True odic table because the effective nuclear charge decreases. False (4) Compound MgSO4 contains both ionic and nonpolar covalent bonds. True False (5) In their gas phase, it requires less energy to remove a single electron from oxygen atom than from nitrogen atom. True False (6) Molecules containing a central atom with sp3 hybridization will all have tetrahedral molecular geometry. True False SECTION 2 (27 PTS. TOTAL; 3 PTS. EACH) Pick the one best answer; if you circle more than one, you will receive NO credit! (1) Which of the following species is paramagnetic? a) Ca b) O2− c) V3+ d) Kr e) None of these (2) Which of the following pairs is isoelectronic? a) Mn2+ and Ar b) Zn2+ and Cu2+ c) Na+ and K+ d) Rb+ and Br− e) None of these (3) Arrange the following ions in order of increasing ionic radius: K+, P3− , S2− , Cl− . a) K+ < Cl− < S2− < P3− b) K+ < P3− < S2− < Cl− c) Cl− < S2− < P3− < K+ d) Cl− < S2− < K+ < P3− e) None of these 3 (4) Which of the following molecules has an atom with an expanded octet? a) CCl4 b) TeCl4 c) IBr d) PCl3 e) None of these (5) Which of the following ionic solids would have the largest magnitude of lattice energy? a) NaF b) CaBr2 c) CsI d) CaCl2 e) None of these (6) The correct Lewis dot structure of AsCl3 will show a) three single bonds and one lone pair of electrons on the central atom b) two single bonds, one double bond, and 9 lone pairs of electrons on the outer atoms c) one single bond, two double bonds, and 8 lone pairs of electrons on the outer atoms d) three single bonds and two lone pair of electrons on the central atom e) None of these (7) Which of the following should have a dipole moment? a) CS2 b) GaCl3 c) XeF2 d) All of these e) None of these (8) Which of the following Lewis structures can be drawn as two or more resonance forms? 2 a) CO 3 (C is the central atom) b) c) d) e) NO2 (N is the central atom) O3 (one O is the central atom) All of these None of these (9) Using Lewis dot structures and formal charges, which of the following ions is expected to be most stable? a) OCN− (carbon is the central atom) b) ONC− (nitrogen is the central atom) c) NOC− (oxygen is the central atom) d) None of these ions are stable e) All of these ions are equally stable 4 SECTION 3 (10 PTS. TOTAL; 2 PTS. EACH) In the center column, write “>”, “<”, “=”, or “?” (can’t tell) to describe the relationship between the following quantities. The first one has been done for you. Number of Chem 102a instructors (1) The maximum number of electrons that an s orbital can hold. < Number of Chem 102a students The maximum number of electrons that a hybrid sp3 orbital can hold. (2) The angle between two equatorial bonds in a trigonal bipyramidal molecule. The smallest angle between two adjacent bonds in an octahedral molecule. (3) The number of lone pairs on the central atom in BrF4+. (Br is the central atom) The number of lone pairs on the central atom in I3−. (one I is the central atom) (4) The formal charge on the sulfur atom in the Lewis structure of SCN− . (carbon is the central atom) The formal charge on the sulfur atom in the Lewis structure of SOCl2. (sulfur is the central atom, S and O are connected via double bond) (5) The length of the nitrogen-nitrogen bond in N2H2. The length of the nitrogen-nitrogen bond in N2H4. SECTION 4 (6 PTS) Credit will not be given if detailed step-by-step work is not shown. Given the following ionization energy (IE) and electron affinity (EA) values, calculate the energy change ( E) for the reaction: Mg2+(g) + 2Cl−(g) → Mg(g) + 2Cl(g) Atom IE1 (kJ/mol) IE2 (kJ/mol) EA (kJ/mol) Mg 738 1450 +3 Cl 1251 2300 −349 5 SECTION 5 (10 PTS. TOTAL; 2 PTS. EACH) Match the following with the best possible answer. (1) A binary compound formed between calcium and nitrogen would be expected to have the formula: _____ (2) The type of hybrid orbitals used by the central atom in XeF4 is _____. (3) The H-N-H bond angle in NH2− is closest to _____. (4) In a molecule, the ability of an atom to attract electrons to itself is called _____. (5) The molecular geometry of a molecule with 3 outer atoms and 2 lone pairs of electrons on the central atom is expected to be _____. (A) bent (M) CaN (Y) electronegativity (B) linear (N) Ca3N2 (Z) 90 ° (C) octahedral (O) Ca2N3 (AA) 109.5° (D) seesaw (P) sp (BB) 120° (E) square planar (Q) sp2 (CC) 180° (F) square pyramidal (R) sp3 (DD) oxygen (G) tetrahedral (S) sp3d (EE) silicon (H) T-shaped (T) sp3d2 (FF) chlorine (I) trigonal bipyramidal (U) bond energy (GG) gold (J) trigonal planar (V) ionization energy (HH) cadmium (K) trigonal pyramidal (W) lattice energy (II) arsenic (L) CaN2 (X) electron affinity (JJ) iron 6 SECTION 6 (12 PTS. TOTAL) Complete the following table. Chemical formula Lewis dot structure (2 pts each) Number of electron groups around the central atom (1 pt each) Electron group geometry (1 pt each) Molecular geometry (1 pt each) Does this molecule/ion have a dipole moment? (1 pt each) SeCl4 IF4− 7 SECTION 7 (6 PTS. TOTAL) Credit will not be given if detailed step-by-step work is not shown. The standard enthalpy of dissociation of H2O(g) into elemental hydrogen and oxygen gases is +241.8 kJ/mol. The bond enthalpies in H2(g) and O2(g) molecules are 436 kJ/mol and 489 kJ/mol, respectfully. Determine the average O-H bond enthalpy (in units of kJ/mol). SECTION 8 (7 PTS. TOTAL) Consider the molecule shown below for the following questions. CH3 H H C C C N C C H2 C O C H2 C N C CH2 H m olec ular geometry? (1) In each box, indicate the hybridization (sp, sp2, sp3, sp3d, or sp3d2) about each indicated carbon atom. (2) How many pi (π) bonds are present in this molecule? __________ (3) The molecular geometry about the indicated carbon atom is: ________________ 8 SECTION 9 (10 PTS.) Credit will not be given if detailed step-by-step work is not shown. Construct the Born-Haber cycle to determine the second ionization energy of calcium given the following data: Heat of sublimation of Ca(s) is 193 kJ/mol; The first ionization energy of Ca(g) is 590 kJ/mol; The bond enthalpy in O2(g) molecule is 498 kJ/mol; electron affinity of O(g) is -141 kJ/mol. Electron affinity of O−(g) is 878 kJ/mol. Hf of CaO(s) is-636 kJ/mol. Lattice energy of CaO(s) is -3414 kJ/mol. 9 1 1 H H 1.00794 3 4 Li Be 6.941 11 9.01218 Na 5 Mg 22.98977 19 Ca K 39.0983 37 40.078 85.4678 87.62 55 Sc Ti 23 V 24 Cr 44.95591 39 47.867 40 50.9415 41 51.9961 88.90585 57 95.94 56 132.90545 137.327 138.9055 87 88 38 Rb Cs Fr (223) Sr Ba Ra (226) Y La 89 Ac (227) Zr 42 91.224 Nb 92.90638 Mo 72 73 74 Hf Ta W 25 Mn A ctinide s Fe 27 Co 28 Ni 29 Cu 55.845 44 58.93320 45 58.6934 46 63.546 47 (98) 75 101.07 76 102.9055 77 106.42 107.8682 79 Tc Re Ru Os Rh Ir 178.49 180.9479 183.84 186.207 190.23 192.217 104 105 106 107 108 Pd 78 Pt 12.0107 14.0067 15.9994 18.99840 14 15 16 17 Ag Au (261) Db (262) Sg (263) Bh (262) 59 60 61 Ce Pr Nd Pm (145) 140.116 140.90765 90 91 Th Pa 232.0381 231.0359 144.24 92 U 238.0289 93 Np (237) Hs Zn 65.39 48 Cd 112.411 80 Hg 63 200.59 31 Ga 69.723 49 In 114.818 81 Tl 204.3833 P 32 33 28.0855 30.97376 Ge S 32.066 35.4527 34 35 72.61 As 74.92160 Se 50 51 52 Sn 118.710 82 Pb 207.2 Sb 78.96 Te 121.760 127.60 83 84 Bi 208.9804 Cl Po (209) Br (265) 64 94 95 157.25 96 At Rn (210) Pu Am Cm (244) (243) (247) 66 67 68 69 70 71 Dy Ho Er Tm Yb Lu Bk (247) 164.9303 167.26 98 99 100 Cf Es Fm (251) (252) (257) 168.9342 101 173.04 102 Md No (258) (259) 36 Kr 131.29 86 65 162.50 39.948 126.9045 85 I Tb 158.9253 97 18 Ar 83.80 (266) Eu Sm 151.964 Gd 150.36 10 79.904 53 Mt 62 196.96655 30 Si 109 Rf 195.078 20.1797 O 13 54.93805 43 58 Lanthanides 26 Ne N 26.98154 22 F C Al 21 1.00794 9 4.002602 8 10.811 24.3050 20 7 B Periodic T le of the Elements ab 12 6 2 He 174.967 103 Lr (262) 54 Xe (222) ...
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This note was uploaded on 01/10/2010 for the course CHEM 102a taught by Professor Hanusa during the Spring '06 term at Vanderbilt.

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