Answers to Questions & Problems in Textbook - Chapter 11

1. Rutherford directed a beam of alpha particles (fairly heavy, fast-moving, positively charged particles) at a very thin piece of gold foil. The plum-pudding model of the atom would suggest that the alpha particles would blast through with very little deflection. However, scientists were shocked when a few alpha particles were occasionally bounced backward. This led Rutherford to explain that the deflection occurred when the alpha particle struck a very small but dense positively charged nucleus.
2. See Figure 11.1 on pg 323 for a sketch of Rutheford's atom. Rutheford could not answer the question: "Why aren't the negative electrons attracted into the positive nucleus, causing the atom to collapse?"
3. Electromagnetic radiation, EMR, is radiant energy that travels through space with wavelike behavior. There are many examples of electromagnetic radiation: visible light, ultraviolet light, radio/television signals, microwaves, infrared heat waves, X-rays, cosmic rays, gammma rays.
4. in your text
5. The wavelength represents the distance between two corresponding points (peaks, troughs, etc.) on successive cycles of a wave. The picture below illustrates a typical wave. Note that the wavelength measured is the same between any two corresponding points on the wave. The frequency of electromagnetic radiation represents how many complete cycles of the wave pass a given point per second. Since the speed of EMR is constant, the wavelength and frequency are inversely related, however, since obviously a longer wave will take more time to pass a given point in space than a shorter length wave.
6. in your text. The speed of EMR is constant for all types.
7. photon
8. in your text. Depending on the nature of the experiment, EMR will exhibit its wave or its particle properties.
9. in your text
10. Higher energy photons are released with Cu2+ because green light is emitted (red light is emitted from Li+. Green light is of higher energy than red light.
11. An atom is said to be in an excited state when it possesses more than its minimum energy (its ground state). An atom is promoted from its ground state to an excited state by the absorption of energy, and returns to its ground state by the emission of the excess energy often in the form of visible light.
12. in your text
13. The emission of light (electromagnetic radiation, EMR) by excited atoms has been a key interconnection between the macroscopic world we can observe and measure, and with what is happening on a nanoscopic basis within an atom. Excited atoms emit light (which we can measure) because of changes in the nanoscopic structure of the atom. By studying the emissions of atoms we can hypothesize about what may have happened inside the atom.
14. in your text
15. Hydrogen always emits light at exactly the same wavelengths (which we see as particular colored lines on the emission spectra), corresponding to the transitions of the electron between the fixed energy states within the atom.
16. in your text. This is the state in which the atom would most prefer to
17. Bohr pictured electrons moving in circular orbits corresponding to the various allowed energy levels. He suggested that the electron could jump up to a different orbit by absorbing energy and subsequently jump down to a different orbit and emit a photon of light with exactly the correct energy content corresponding to the difxference in energy between the orbits.
18. in your text
19. Bohr suggested that the electron could jump up to a different orbit by absorbing energy and subsequently jump down to a different orbital and emit a photon of light with exactly the correct energy content corresponding to the difference in energy between the orbits. Since the energy levels of a given atom were fixed and definite, then the atom should always emit energy at the same discrete wavelengths thereby always showing the same exact bright line emission spectrum.
20. in your text
21. Schrodinger and de Broglie reasoned that, since light seems to have both wave and particle characteristics (it behaves simultaneously as a wave and as if it were a stream of particles), that perhaps the electron might exhibit both of these characteristics. That is, although the electron behaves as a discrete particle, perhaps the properties of the electron in the atom could be treated as if they were wavelike.
22. in your text
23. Any experiment which sought to measure the exact location of an electron (such as shooting a beam of light at it) would cause the electron to move. Thus any measurement made would necessitate the application or removal of energy, which would disturb the electron from where it had been before the measurement.
24. Chemists arbitrarily defined an orbital to represent a 90% probability of finding the electron within this region. Orbitals represent mathematical functions and have no distinct outer edge (so they are drawn to appear "fuzzy": they are not hard-edged capsules enclosing the electron, but merely represent the most likely region where an electron may be found.)
25. in your text
26. in your text
27. The higher the principal energy level (n), the farther from the nucleus, on average, the electron will be.
28. in your text
29. The other orbitals serve as the excited states of the hydrogen atom. When an energy source of an appropriate frequency is applied to the hydrogen atom, the electron can move from its normal orbital (ground state) to one of the other orbitals (excited states). When the energy source is removed, the electron can move back to its normal orbital and release the absorbed energy as light.
30. in your text
31. The higher the value of the principal quantum number, n, the higher the energy of the principal level.
32. in your text
33.
  1. correct (the n=2 shell contains s and p subshells)
  2. incorrect (the n=1 shell consists only of 1s orbitals, no d orbitals)
  3. incorrect (the n=3 shell contains 3s, 3p, and 3d orbitals, no f orbitals)
  4. correct (the n=4 shell contains s, p, d, and f subshells)
34. The 1s orbital is closest to the nucleus and lowest in energy, so it is always filled first.
35. in your text
36. The elements in a given vertical column of the periodic table have the same valence electron configuration. Having the same valence electron configuration causes the elements in a given group to have similar chemical properties.
37.
  1. in your text
  2. 1s2 2s2 2p6 3s2 3p6 4s2 3d10
  3. in your text
  4. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
38.
  1. in your text
  2. 1s22s22p63s23p64s1
  3. in your text
  4. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
39.
  1. in your text
  2. Phosphorus 15 electrons:
  3. in your text
  4. Argon 18 electrons:
40.
  1. in your text
  2. two (4s2)
  3. in your text
  4. five (2s2, 2p3)
41. This belief is based on the experimental evidence of K and Ca. The physical and chemical properties of K are like those of the other group 1 elements, and Ca's properties are similar to the other group 2 elements.
42. in your text
43.
  1. in your text
  2. [Rn] 7s1
  3. in your text
  4. [Xe] 6s2 4f2
44.
  1. in your text
  2. [Ne] 3s2 3p5
  3. in your text
  4. [Ar] 4s2 3d10
45.
  1. in your text
  2. two
  3. in your text
  4. ten
46.
  1. in your text
  2. 5f
  3. in your text
  4. 6p
47.
  1. in your text
  2. [Ar] 4s2 3d5
  3. in your text
  4. [Rn] 7s1
48. in your text
49. The Group 1 metals are all highly reactive, and all form 1+ ions almost exclusively when they react. Physically, these metals are soft (they can be cut with a knife) and very low in density. Because of their high reactivity, these metals tend to be found with a coating of the metal oxide which hides their metallic luster (which can be seen, however, if a fresh surface of the metal is exposed).
50. All exist as diatomic molecules (F2, Cl2, Br2, I2); all are nonmetals; all have relatively high electronegativities; all form 1- ions in reacting with metallic elements.
51. The nonmetallic elements are clustered at the upper right side of the periodic table. These elements are effective at pulling electrons from metallic elements for several reasons. First, these elements have little tendency to lose electrons themselves (they have high ionization energies). Secondly, the atoms of these elements tend to be small in size, which means that electrons can be pulled in strongly since they can get closer to the nucleus. Finally, if these atoms gain electrons, they can approach the electronic configuration of the following noble gas elements (see chapter 12 for why the electronic configuration of the noble gases are desirable for other atoms to attain).
52. in your text
53. All atoms within a given group have the same number of valence electrons, and these valence electrons occur in the same type of subshell. However, at the bottom of a group, the valence electrons are in a higher energy level making them further away from the nucleus.
54. in your text
55.
  1. in your text
  2. Be (the less reactive metals are further up in a group because metals are trying to lose electrons and its harder to lose electrons that are closer to the positively charged nucleus.)
  3. in your text
  4. The (the less reactive nonmetals are at the bottom of a group because nonmetals are trying to gain electrons and the pull on any incoming electrons is less, when they are being pulled into orbitals further away from the nucleus.)
56.
  1. in your text
  2. Ca because it is the smallest of the three and therefore harder to remove an electron from its outermost orbital.
  3. in your text
  4. S because it is the smallest of the three and therefore harder to remove an electron from its outermost orbital.
57.
  1. in your text
  2. He
  3. Ar
58. wavelength
59. in your text
60. quantized
61. in your text
62. You should use your periodic table for this problem.
  1. in your text
  2. Titanium: condensed version: [Ar] 4s23d2
  3. in your text
  4. Iron: condensed version: [Ar] 4s23d6
  5. Zinc: condensed version: [Ar] 4s23d10
63.
  1. in your text
  2. ns2 np5
  3. in your text
  4. ns1
  5. ns2 np4
64.
  1. in your text
  2. seven
  3. in your text
  4. two (the d electrons are not counted as valence electrons)
65. in your text
66. The three 2p orbitals of carbon are of the same energy; by occupying different orbitals of the same energy, repulsion between electrons is minimized.
67.
  1. in your text
  2. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
  3. in your text
  4. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
68. metals, low; nonmetals, high
69.
  1. in your text. (For the same reason as part c below.)
  2. P, because it is further to the right on the chart yet in the same period (row). This means that P's electrons are in the same energy level, but with more protons in its nucleus, it can pull those electrons in closer, making P smaller in size than the Si or Al.
  3. K, because it is further up on the periodic chart meaning it has less energy levels than the other two atoms in the same column. Less energy levels of course will be a smaller atom.
70. skip this problem
71.
  1. iv, because it has 6 valence electrons and would love to gain two more to become 8 valence electrons.
  2. ii, because of its 1 valence electron that is most willing to leave.
  3. ii, because of its 1 valence electron that is most willing to leave. The leaving process is its reactivity.
  4. v, because the 2p3 electrons are in 3 separate orbitals all unpaired because with within a subset (type of orbital) the electrons will spread out as much as possible.
  5. i, because its two valence electrons will gladly transfer to oxygen which wants to take on two valence electrons.
  6. Na3N
  7. Mg3N2
  8. ii < i < vi < v < iv < iii
72.
  1. F
  2. Yb (Lu)
  3. Po
  4. Fr
  5. At
  6. Xe
  7. No (Lr)
73. Element 120 would be a solid and a metal. It would be expected to be highly reactive with water and have a tendency to form ions with a 2+ charge. Essentially it would have the same chemical properties as the elements in group 2A, the alkaline earth metals.
74. skip this question