Page 141 (#15, 16)
15. f = 1.50 x 1013 Hz
λ=?
Use the wave equation c = λxf
λ = 2.00 x 10-5m; this wavelength is longer than red light
16.
λ = 5.00 cx 10-8 m
f=?
6.00 x 1015 s-1 ; it is in the UV range
Page 152 (#45-48)
#45. Use a diagram to illustrate each of the following terms:
a. wavelength
b. amplitude
c. cycle
46. Frequency is the number of wave cycles that pass a given
point per unit of time. Frequency units are cycles/second or s-1
or Hertz. Wavelength and frequency are inversely related as can
be if we rearrange the wave equation to get λ=c÷f
47. Arrange the regions of the EMS in order of decreasing wave
length:
a. v. radio waves ,vi. microwaves ,iv. infrared ,iii. visible , i.
ultraviolet, ii. X-ray
b. Arrange in order of increasing wavelength
ii. Xray, i. ultraviolet, iii. Visible, iv. Infrared, vi. Microwaves,
v. radio waves
48. The visible colors in order of increasing wavelength are:
ROYGBIV is decreasing wavelength therefore the order in
increasing wavelength is violet, indigo, blue, green, yellow,
orange, red
Page 152 (27-29 and 49-51)
#27. Rutherford’s model was known as the planetary model
because he proposed electrons move around the nucleus like the
planets move around the sun.
28. Bohr assumed that the electrons traveled in circular paths
around the nucleus.
29. In Rutherford’s model, negatively charged electrons orbit a
positively charged nucleus. In Bohr’s model, electrons orbit the
nucleus in defined concentric circular paths of fixed energy.
49. Planck showed mathematically that the amount of radiant
energy (E) of a single quantum absorbed or emitted by a body is
proportional to the frequency of radiation.
E = hf
50. A photon is a packet or quanta of light; a quantum is a
discrete amount of energy.
51. UV light has higher energy than the photon of IR radiation.
Page 153 (#52-55)
52. What is the energy of a photon with f = 5.80 x 1014 Hz?
Use the equation E = hxf
E= 6.626 X 10-34 J s(5.80 x 1014 Hz)
E=3.84 x 10-19 J
53. Classical physics views energy as continuous whereas in the
quantum model the energy changes occur in discrete packets
called quanta.
54. When a hydrogen atom absorbs a quantum of energy the
electron gets excited and jumps out to a higher energy level.
55. Transitions that occur from energy levels down to n= 2 are
known as Balmer series and these are wavelengths in the
visible portion of the EMS.
Page 152 (30-34)
#30. In the electron cloud model, the boundary represents the
outer edge of the area where an electron may be found 90% of
the time.
31. An atomic orbital is a region in space around the nucleus in
which there is a high probability of finding an electron.
32. 1s is spherical (see Fig 5.4)
2s is spherical with radius greater than the 1s
2p is dumbbell shaped and there are three located at right
angles to each other
33. There are 3 orbitals in the 2p sublevel.
34. In each of the following principal energy levels the number
of sublevels is:
a. n = 1 has 1 (1s)
b. n = 2 has 2 (2s and 2p)
c. n = 3 has 3 (3s, sp, and 3d)
d. ne – 4 has 4 (4s, 4p, 4d and 4f)