1. Calculate the following…
a. The frequency of light of 510.0 nm.
b. The wavelength of light with a frequency of 1.5 x1013 Hz. Is this wavelength
longer or shorter than that of red light (700 nm)?
c. The frequency of radiation with a wavelength of 5.0x10-8 m. Where is this on the
electromagnetic spectrum (i.e., what kind of radiation is it)?
d. The wavelength of light of 6.0x104 Hz.
e. The frequency of gamma rays with a wavelength of 5.0x10-13 m.
2. Calculate the frequency of a photon of light that has a wavelength equal to:
a. 486 nm
b. 7.50x10-5 cm
c. 6200 Å
3. Find….
a. The energy of light of 0.0012 Hz.
b. The energy associated with purple light of 385 nm.
c. The energy associated with UV waves of 3.00x1016 Hz.
d. The energy associated with X-rays of 3.00x1018 Hz.
e. The energy associated with gamma rays of 3.00x1020 Hz.
f. The energy of light of 15.0 cm.
4. What is a quantum?
5. What are photons?
6. Briefly, what is the photoelectric effect?
7. What is meant by light having a “dual nature”. Give examples of each.
8. What is the relationship between energy and the wavelength of light?
9. Find the following:
a. The energy associated with radio waves of 80.0 m.
b. The energy associated with infrared waves of 0.0000112 m.
c. The energy associated with red light of 650 nm.
d. The energy associated with blue light of 465 nm.
Use ∆E = hc/λ and/or E = -2.178 x 10-18 J (Z2/n2) for the following:
10. A certain spectral line in the Balmer series (electrons fall to n=2) has a wavelength equal
to 656.3 nm. What was the initial principal quantum state of the electron?
11. A mole of hydrogen atoms undergoes a transition from N=7 to N=1. What is the
wavelength of the photons produced? Can you see these photons? What part of the
electromagnetic spectrum are they in?
12. Chlorophyll absorbs blue light (λ = 460 nm) and emits red light (λ = 660 nm). Calculate
the net energy change in chlorophyll when a single photon of 460 nm radiation is
absorbed and one of 660 nm is emitted.
13. Calculate the wavelength of the photons produced when a mole of hydrogen electrons
undergoes a transition from N=6 to N=2. Can you see this photon?
14. What wavelength of light will cause a transition from N=1 to N=5 in the hydrogen atom?
What part of the electromagnetic spectrum is this light in?
15. If a mole of hydrogen atoms undergoes a transition from N=8 to N=3, what wavelength
will the emitted photons have?
16. If a mole of hydrogen atoms undergoes a transition from N=8 to N=2, will the photons
produced be visible? What region of the electromagnetic spectrum will the photons be
in?
17. A mole of hydrogen atoms undergoes a transition from N=∞ to N=2. What is the
wavelength of the photons produced?
18. A student viewing an excited sodium atom observes a strong spectral line at 589 nm. The
color is bright yellow. What is the energy that would correspond to this wavelength?
19. A sample of hydrogen atoms already in the ground state is illuminated with white light.
The white light causes a transition to the N=8 state. Calculate the wavelength of the light
that caused this transition to occur. Is the spectrum produced an emission or absorption
spectrum?
20. A particular AM radio station broadcasts at a frequency of 1120 kHz. What is the
wavelength, in nm, and energy per photon of this radiation?
21. When a mole of hydrogen atoms undergoes a transition from N=8 to a lower stationary
state, it emits photons with a wavelength of 9.5x10-5 cm. What is the final principle
quantum state of the electrons?
22. Calculate the wavelength of light needed for ionization of an electron in the N=2
principle quantum state of an atom of hydrogen.