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CHEM 1411
PRACTICE EXAM III (Chapters 7, 8, 9): 25 questions. Zumdahl 9th edition.
Multiple Choices: Select one best answer.
1. Which of the following statements concerning four quantum numbers is correct?
(a) The angular (i.e. azimuthal) quantum number, l, determines the energy of an
electron in a hydrogen atom and in a many‐electron atom.
(b) The principle quantum number, n, determines the shape of an orbital.
(c) The electron spin quantum number, ms, determines the orientation of an orbital
in space.
(d) The magnetic quantum number, ml, determines the size of an orbital.
(e) The principle quantum number, n, determines the size of an atom.
Hint: Definitions. Section 7.6: p. 313.
2. Which of the following four quantum numbers (n, l, ml, ms) does not represent
an electron in a 3d orbital?
(a) (3, 2, ‐1, +½)
(b) (3, 2, ‐2, ‐½)
(c) (3, 2, 0, ‐½)
(d) (3, 2, +1, +½)
(e) (3, 1, +1, +½)
Hint: Applications of Definitions. See p. 345: End-of-Chapter-Exercise 96.
3. Which of the following sets of quantum numbers in an atom is acceptable or
correct?
(a) (1, 0, +½, +½)
(b) (3, 2, ‐2, ‐½)
(c) (3, 3, 0, ‐½)
(d) (4, 2, +3, +½)
(e) (3, 0, +1, +½)
Hint: Applications of Definitions. See p. 344: End-of-Chapter-Exercises 73 & 74.
4. Which of the following is a correct ground‐state electron configuration?
(a) Al: 1s22s22p43s23p3
(b) B: 1s22s22p5
(c) F: 1s22s22p6
(d) Cu: 1s22s22p63s23p64s23d9
(e) Cr: 1s22s22p63s23p64s13d5
Hint: See section 7.11. Ground‐state electron configuration must obey Pauli Exclusion
Principle and the Hund’s Rule. Memorize the orbital energy shown in Fig 7.23 and the
chart given by the instructor. Exceptions for Cr and Cu: see p.p. 325 - 326. See p. 346
End-of-Chapter-Exercise 100.
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5. How many unpaired (i.e. spin-unpaired) electrons are present in Mn?
(a) 1
(b) 2
(c) 3
(d) 4
(e) 5
Hint: Draw the orbital diagram shown in p.p. 323-324. Mn valence electron
configuration is shown in p. 326. Core electrons are always spin-paired. So we just
need to examine the valence electrons only: 4s23d5. Remember Hund’s rule must be
applied to the degenerate orbitals.
6. What is the maximum number of electrons in an atom that has the quantum
numbers as n=2, l=0, ms= +½?
(a) 1
(b) 2
(c) 3
(d) 4
(e) 5
Hint: See p. 345 End-of-Chapter-Exercises 81 & 82.
7. What is the maximum number of electrons in an atom that has the quantum
numbers as n=4, ml= +1?
(a) 2
(b) 4
(c) 6
(d) 8
(e) 10
Hint: See p. 345 End-of-Chapter-Exercises 81 & 82, especially 81(b).
8. Which of the following violates the Pauli Exclusion Principle?
(a) ↑ ↑ ↑↑ (p orbitals)
(b) ↑ ↑↓ ↑ (p orbitals)
(c) ↑ ↑↓ ↓ (p orbitals)
(d) ↑↓ ___ ↑ ↑ ↑ (d orbitals)
(e) ↑ ↑ ↑ ↓ ↑↓ (d orbitals)
Hint: Memorize the definition. Here, (a), (b) and (c) refer to the p orbitals, while (d)
and (e) refer to the d orbitals. Pauli Exclusion Principle focuses on the spin directions
of two electrons in the same orbital. Do not answer anything related to the Hund’s
Rule.
8. Which of the following violates the Hund’s rule?
(a) ↑ ↑ ↑↑ (p orbitals)
(b) ↑ ↑↓ ↑ (p orbitals)
(c) ↑ ↑↓ ↓ (p orbitals)
(d) ↑↓ __ ↑ ↑ ↑ (d orbitals)
(e) ↑ ↑ ↑ ↓ ↑↓ (d orbitals)
Hint: Memorize the definition. Here, (a), (b) and (c) refer to the p orbitals, while (d)
and (e) refer to the d orbitals. Hund’s Rule focuses on the degenerate orbitals like p,
d, and f. Each of the orbitals must have one electron in it before putting the second
electron in it. It does not matter whether the two electrons are in opposite directions
or not. Do not answer anything related to the Pauli Exclusion Principle.
9. Which of the following atoms has the electron configuration as [Ar] 4s23d7?
(a) Sc
(b) V
(c) Fe
(d) Co
(e) Zn
Hint: See section 7.11 starting p. 322. Since the species is an atom which is electrically
neutral and thus number of electrons is same as the number of protons. Thus the
total electrons shown here = 18+2+7 = 27 = total number of protons = atomic
number. Thus from the Periodic Table, the atom is Co.
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10. Which of the following electron configurations do represent similar chemical
properties of their atoms?
(i) 1s22s22p63s2
(ii) 1s22s22p3
(iii) 1s22s22p63s23p64s23d104p6
(iv) 1s22s2
(v) 1s22s22p6
(vi) 1s22s22p63s23p3
(a) (i), (v)
(b) (ii), (vi)
(c) (iii), (vi)
(d) (i), (iii)
(e) (ii), (iii)
Hint: See section 7.11. Look for the valence electron configuration by converting the
configuration to the one with noble gas core. Atoms with same number of valence
electrons are in the same group and thus will have similar chemical properties (see
chapter 2). Another method is by adding the electrons altogether to get the total
electron number which is same as the atomic number or say proton number in an
atom. Then locate its symbol and location in the periodic table.
11. What is the ion with +3 charges that has the electron configuration as [Ar]3d3?
(a) Sc3+
(b) Cr3+
(c) Co3+
(d) Rh3+
(e) Ir3+
Hint: We need to figure out the number of protons (i.e. atomic number). Since the
the complete electron configuration given is [Ar]3d3indicating there are 18+3 = 21
electrons for +3 charged ion. Thus, its number of protons will be 18+3+3 = 24.
From the periodic table, it shows that Cr has 24 protons. So its ionic form is Cr3+
12. A metal ion with a net +3 charge has five electrons in the 3d subshell. What is
this metal?
(a) Cr
(b) Mn
(c) Fe
(d) Co
(e) Ni
Hint: This question shows only the valence electrons for the +3 charged ion (i.e. it has
five electrons in the 3d subshell). To have electrons in the 3d orbitals, the ion must
have electrons completely filled in the lower energy orbitals (i.e. all the core orbitals
(i.e. 1s, 2s, 2p, 3s, and 3p) are completely filled 18 electrons in the core orbitals. Thus,
the total number of electrons for this +3 ion will be 18+5 = 23 and the number of
protons is 23+3 = 26. So this metal ion is Fe3+ and the metal (electrical neutral) is Fe.
Note: Transition metal (with d electrons) loses its ns electrons prior to its (n-1)d
electrons. See section 8.4 and p. 406 End-of-Chapter-Exercises 45 & 46.
13. Which of the following pairs are not isoelectronic?
(a) K+ and Cl‐
(b) Ca2+ and S2‐
(c) Al3+ and N3‐
(d) P3‐ and Sc3+
(e) Rh3+ and Ir3+
Hint: See p. 363 section 8.4 for definition. Isoelectronic means having same total
electron number.
Note: If this question changes to: Which of the following pairs are not
isoelectronic with Ar?
(a) K+ and Cl‐
(b) Ca2+ and S2‐
(c) Al3+ and N3‐
(d) P3‐ and Sc3+
(e) Rh3+ and Ir3+
Hint: The condition “isoelectronic with Ar” indicating that the ions must have 18
electrons.
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14. Which of the following is correct?
(a) Atomic radii: Ar < P < Na < Ca < Cs
(b) For electron affinity: F < Na < O < Al < Mg
(c) For the first ionization energy: K < Ca < P < F < Na
(d) ionic radii: F‐ < O2‐ < N3‐ < Na+ < Mg2+ < Al3+
Hint: (a) Go with Periodic Trend for Atomic Radii (p. 334 section 7.12).
(b) See p.p. 332-333 for Electron Affinity including Fig 7.33 and Table 7.7. See the
Periodic Trend given on the white board in class.
(c) See section 7.12 for Periodic Trend for Ionization Energy: p.p. 329‐331 including
Fig 7.31.
(d) See section 8.4 for Ionic Radii: p.p. 363-365 and the Summaries given on the
white board in class.
15. Which of the following is a correct order of lattice energy?
(a) NaCl < AlCl3 < MgCl2
(c) Na2O < MgO < Al2O3
(b) LiF < LiCl < LiBr
(e) LiCl < NaCl < KCl
(d) Ga2O3 < CaO < K2O
Hint: See section 8.5: be sure to use the lattice energy formula. Consider charge factor
(Q+Q‐) first as it is the most important. If the lattice energy cannot be determined by the
charge factor, then consider the distance factor (1/r) which is less important.
16. Which of the following is a polar covalent bond?
(a) The HN bond in NH3.
(b) The SiSi bond in Cl3SiSiCl3.
(c) The CaF bond in CaF2.
(d) The OObond in O2.
(e) The NN bond in H2NNH2.
Hint: See sections 8.3: Bond Polarity and 8.6: Partial Ionic Character of Covalent Bonds
(i.e. polar covalent bonds). Covalent bond is formed by electron-sharing between two
nonmetal atoms and the “partial” means electrons have preference to stay closer to a
specific atom whose electronegativity (p.p. 356-358, especially Figure 8.3: p. 369) is
greater (Figure 8.13). In CHEM 1411, we call a bond is nonpolar only when the
electronegativity difference (either s zero or positive) of bonded atoms is ZERO. Here
(c) is wrong because Ca is a metal and thus Ca-F is an ionic bond.
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17 & 18. Draw the Lewis structures for (a) I3 , (b) BF3, (c) CO2, (d) CN , and (e)
NO3 -.
Hint: See section 8.10: p. 379 for CO2 and CN , p. 380 for I3 , p. 381 for BF3, and p. 385 for
NO3 - .
19. Which of the following does not have resonance or delocalized structure (see section
8.12: if not familiar, draw Lewis structure)?
(a) NO3(b) CO3-2
(c) C6H6
(d) NO2
(e) PF5
20. Which of the following obeys the octet rule?
(a) BF3
(b) SF4
(c) NO2
(d) PF5
(e) NO3Hint: See section 8.11. If not familiar, draw Lewis Structure. To obey the Octet Rule, the
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Central atom must have exact 8 electrons.
21. What is the ∆Hrxn (kJ) for the reaction H2(g) + I2(g) → 2HI(g)?
The bond enthalpies for H‐H, I‐I and H‐I are 436.4, 151.0 and 298.3 kJ/mol, respectively.
(a) ‐9.2
(b) +9.2
(c) ‐289.1
(d) +289.1
(e) 0.0
Hint: See section 8.8: p.p. 373-376 especially formula in p. 375.
22. Which of the following is a polar molecule?
(a) SF6
(b) CHCl3
(c) BCl3 (d) PF5
(e) XeF4
Hint: See section 8.3: p. 358-361 and End-of-Chapter Exercise 125.
23. Which of the following has the bond angle as 120o?
(a) C2H2
(b) SO42(c) AsF3
(d) SF6
(e) HCOH (formaldehyde)
Hint: Draw the Lewis structure and apply section 8.13: Memorize Table 8.6 and 8.7.
24. What is the hybrid orbital for Xe in XeF4?
(a) sp
(b) sp2
(c) sp3
(d) sp3d
(e) sp3d2
Hint: Draw the Lewis structure and apply section 9.1: Memorize Figure 9.24.
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25. Which of the following has the same bond order (p.p. 430-431) as He2 ?
(a) H2+
(b) H2
(c) He2
(d) Li2
(e) None of the above
Hint: See section 9.2 and use the formula in p. 430. Bond order is same as bond. Lewis
structure does not tell geometry (or shape). If the total electron number is even, you can
draw the Lewis structure to figure out the bond order. Bond order is either zero (i.e. no
bond) or positive (either whole number or fraction). If the bond order is zero, the
molecule or ion does not exist; if the bond order is 0.5, it is the half bond. If the bond
order is two, it is the double bond.
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