Download Chapter 4 Chapter Review Question Answers A. Radio waves (long

Chapter 4 Chapter Review Question Answers
1.
A. Radio waves (long and short/AM & FM), microwaves, infrared light, visible light, ultraviolet light, X rays, gamma rays
b. 3.00 × 108
2.
3.
4.
5.
6.
𝑚
𝑠
Light’s wave properties are noticed more when light is traveling and includes measurable characteristics like wavelength
and frequency, as well as interference, refraction and diffraction. Light’s particle properties are noticed more when light is
interacting with matter, and includes the emission of light by hot objects, the photoelectric effect, and the absorption and
emission spectra lines in atoms/elements.
The frequency range is
4.29 × 1014 𝐻𝑧
𝑡𝑜
7.50 × 1014 𝐻𝑧
𝑜𝑟
429 𝑇𝐻𝑧 𝑡𝑜 750 𝑇𝐻𝑧
The wavelength range is
400 nm to 750 nm
Red, Orange, Yellow, Green, Blue, Violet
The wave theory could not explain the photoelectric effect and the line-emission spectra of elements.
A. 𝑐 = 𝜆𝜈 c = velocity (speed of light), λ = wavelength, and ν = frequency
b. 𝐸 = ℎ𝜈 E = Energy of photon, h = Planck’s constant, and ν = frequency
c. 𝐸 =
ℎ𝑐
𝜆
7.
8.
a. Wave Theory
b. Particle Theory
c. Particle Theory
The ground state of an atom is the atom’s lowest energy state; all electrons are in lowest energy positions. An excited state
of an atom is any energy state that is higher in energy than the ground state; one or more electrons are excited and moved
to a higher energy position.
9. Bohr states that a line-emission spectrum is produced when an electron drops from a higher-energy orbit releasing a
photon. The energy released by the photon is equal to the difference in energy between the two levels.
10. 7.05 × 1016 Hz
11. 2.35 × 10−16 J
12. 𝐸 =
ℎ𝑐
𝜆
13. 267 𝑠
14. 1.99 × 10−13 J
15. Bohr’s atomic model could not explain chemical properties of atoms, and its mathematics only works for the hydrogen
atom.
16. A. quantum number used to the main energy level of an atom’s electrons
b. n
c. all the orbitals within the same energy level
d. 2n2 is the calculation for the number of electrons in an energy level.
17. A. The angular momentum quantum number indicates the type of orbital, and therefore, the shape of the orbital.
b. A sublevel or subshells are divisions of orbitals that share the same l value.
18. A. 1, s
b. 2, s & p
c. 3, s, p, & d
d. 4, s, p, d, & f
e. 7
19. A. The magnetic quantum number gives the orbital orientation around the nucleus.
b. s = 1, p = 3, d = 5, f = 7
c. The different p orbitals are distinguished using subscripts to identify which axis around the nucleus the orbitals are
aligned. For example: px py & pz
20. A. The relationship between n and the number of orbitals in an energy level is given by the equation: n 2
b. 3rd energy level has 9 orbitals (1 s, 3 p, 5 d), the 5th energy level has 25 orbitals (1 s, 3 p, 5 d, 7 f, 9 g)
21. A. The spin quantum number indicates the spin state and resulting magnetic field of an electron in an orbital.
b. +½ & −½
22. A. 2
b. 18 c. 32 d. 72 e. 98
23. S orbital is spherical in shape, p orbitals are dumbbell shaped
24. 2s orbital is higher in energy for the electrons than the 1s orbital. 2s orbital is further from the nucleus. 2s orbital is larger.
25. They would be at right angles to each other.
26. A. Electrons occupy the lowest energy orbital they can.
b. Multi-electron atoms will have electrons filling the lowest energy orbitals first. When those are filled the electrons will
begin to fill the next lowest energy orbitals, and so on until all the electrons are located.
27. a. Hund’s rule states that when electrons enter orbitals of equal energy, they enter each orbital one at a time, with the
same spin. When each orbital has one electron with the same spin, then other electrons will fill the orbital by pairing up.
b. Having single electrons separated as much as possible by placing them in different oriented orbitals, minimizes electronelectron repulsion, and the atom ends up with a lower energy arrangement. This is confirmed by experimental evidence.
28. a. No two electrons in the same atom can have the same set of quantum numbers.
b. The spin quantum number establishes that two electrons in the same orbit must have opposite spins or there would be
higher energy required for the atom due to electron-electron repulsion.
29. a. The highest occupied energy level in an atom in the ground state is the electron-containing energy level that has the
highest principal quantum number.
b. Inner-shell electrons are electrons that are not in the highest occupied energy level.
30. a. First energy level (n = 1)
b. Second energy level (n = 2)
c. Third energy level (n = 3)
d. Fourth energy level (n = 4)
e. Fifth energy level (n = 5)
31. a. P ____ ____ ____ ____ ____ ____ ____ ____ ____
1s
2s
2p
3s
3p
b. B ____ ____ ____ ____ ____
1s
2s
2p
c. Na ____ ____ ____ ____ ____ ____
1s
2s
2p
d. O ____ ____ ____ ____ ____
1s
2s
2p
32. a. Li 1s2 2s1
b. Ne 1s2 2s2 2p6
c. O 1s2 2s2 2p4
d. Al 1s2 2s2 2p6 3s2 3p1
33. a. 8
b. 8
c. O ____ ____ ____ ____ ____
d. 2
1s
2s
2p
e. 2nd energy level (n = 2)
f. 2
g. 1s orbital
34. a. Noble gas elements are the elements in group 18 of the periodic table, He, Ne, Ar, Kr, Xe, & Rn
b. Noble gas configuration refers to an outer electron energy level filled (sort of) resulting in chemical stability. This is
usually a filled p-orbital. (Except He, which has the 1s filled)
35. a. Cl [Ne] 3s2 3p5
b. Ca [Ar] 4s2
c. Se [Ar] 3d10 4s2 4p4
36. a. This indicates that this atom has all the electrons that an atom of Neon would have and then also 2 electrons in the 3s
orbital. b. Magnesium
37. a.
Na 1s2 2s2 2p6 3s1
Na [Ne] 3s1
2
2
6
2
6
10
b.
Sr 1s 2s 2p 3s 3p 3d 4s2 4p6 5s2
Sr [Kr] 5s2
c.
P 1s2 2s2 2p6 3s2 3p3
P [Ne] 3s2 3p3
38. a. Boron (B)
b. Fluorine (F) c. Magnesium (Mg) d. Silicon (Si) e. Chlorine (Cl) f. Potassium (K) g. Iron (Fe)
39. 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
40. a. As [Ar] 3d10 4s2 4p3
b. Pb [Xe] 4f14 5d10 6s2 6p2
c. Lr [Rn] 5f14 6d1 7s2
d. Hg [Xe] 4f14 5d10 6s2
e. Sn [Kr] 4d10 5s2 5p2
f. Xe [Kr] 4d10 5s2 5p6
g. La [Xe] 5d1 6s2
41. In these examples, the electrons occupy the 3d orbitals before filling the 4s orbital. This results in lower energy
arrangement for the electrons.
42. a. Yellow light has a longer wavelength than green light
b. X-rays will have higher frequency than microwaves
c. They both travel at the same constant speed; the speed of light: 3.00 × 108 m/s
43. a. Ar 1s2 2s2 2p6 3s2 3p6
Ar [Ne] 3s2 3p6
2
2
6
2
6
10
2
b. Br 1s 2s 2p 3s 3p 3d 4s 4p5
Br [Ar] 3d10 4s2 4p5
c. Al 1s2 2s2 2p6 3s2 3p1
Al [Ne] 3s2 3p1
−5
44. 4.00 × 10 m
45. a. The electromagnetic spectrum (EM) is the range of wavelengths and frequencies of electromagnetic radiation.
b. Any length unit can be used. Short wavelengths are measured in nanometers (nm) and longer wavelengths are
measured in meters.
c. Frequencies are measured in Hertz (Hz)
46. a. 15 electrons
b. atomic number is 15
c. See 31 a. above
d. 3 unpaired electrons
e. 3rd energy level
f. 10 inner shell electrons
g. 1s, 2s, & 2p
9
47. 2.34 × 10 Hz
48. a. Hf [Xe] 4f14 5d2 6s2
b. Sc [Ar] 3d1 4s2
c. Fe [Ar] 3d6 4s2
d. At [Xe] 4f14 5d10 6s2 6p5
e. Ac [Rn] 6d1 7s2
f. Zn [Ar] 3d10 4s2
49. Bohr’s model only worked for the hydrogen atom, whereas Schrödinger’s model works for all atoms. The major difference
between the two models is the issue of certainty. Bohr used calculations with electrons as particles in defined orbits, whereas
Schrödinger applied wave properties to electrons and then described the location of electrons in terms of probability. Both models
describe an electron’s energy with its location relative to the nucleus. Finally, Schrödinger’s model includes that the electron in a
hydrogen atom is at a distance from the nucleus that is exactly equal to the value calculated by Bohr for the lowest energy orbit.
50. a. 5.09 × 1014 Hz (509 THz)
b. 5.90 × 10−7 m (590 nm)
51. a. An orbital is a 3-D space around the nucleus where there is a high probability that an electron is likely to be located.
b. Orbitals are like clouds that show the region of probable electron locations. The sizes and shapes of electron clouds
depend on the energies of the electrons that occupy them.
52. The absorption lines are at 656 nm, 486 nm, 434 nm, and 410 nm. These are exactly the same as the emission line
wavelengths for hydrogen. This is because, an electron must absorb exactly the right energy to jump to a new energy level, and
when it goes back to the lower energy level, it releases the exact same amount of energy it absorbed.
53. 7.00 × 10−19 J
54. Electrons a & b are in the same orbital, as they are identical in numbers except for their spin numbers.
55. Possible: b; all quantum numbers are possible
Not Possible: a; If n = 2, then the only l values are 0 and 1. c; If l = 0, then the only m value possible is 0.
56. a. Flame tests
b. Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Calcium (Ca), Strontium (Sr), & Barium (Ba)
c. Chloride salts are used to produce colors in fireworks.
d. Wavelengths of around 500 nm would be visible parts of Barium’s emission spectrum.