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Atomic Physics Worksheet
- Which of the gas samples shows an emission line with a wavelength
between 4000 and 5000 Angstroms (400-500 nanometers): __________________
between 6500 and 7500 Angstroms (650-750 nanometers):__________________
(There may be more than one.)
n = ∞ (^) E = 0 eV
n = 5
n = 6 E = ñ0.54 eV
E = ñ0.38 eV
n = 4 (^) E = ñ0.85 eV
n = 3 (^) E = ñ1.51 eV
n = 2 (^) E = ñ3.40 eV
n = 1 (^) E = ñ13.6 eV
In the diagram above we have represented some allowed energy levels of the electron in a hydrogen atom. Ordinarily, the electron is in the ìground stateî ( n = 1) and has energy of ñ13.6 eV. ( eV = 1.6 × 10 ñ19^ J.) This means that you would have to add 13.6 eV of energy to the electron in order to break it away from the hydrogen atom. (If the electron just barely escaped from the atom ñ and had no energy left over to go anywhere ñ we would say that it has ìzeroî total energy, or E = 0 eV.) If an electron in the ground state ( n = 1) gains energy its energy level would rise, for instance to n = 2, n =3, or some higher value of n. The larger the value of n , the greater is the energy of the electron. The maximum possible energy for an electron in the hydrogen atom (E ≈ 0 eV) would correspond to n = ∞.
When fast-moving electrons are sent flowing through a tube of hydrogen gas, some of them collide with the hydrogen atoms and increase the energy of the electrons in the atoms. In this way, the hydrogen electrons may briefly acquire some of the higher energy levels shown in this diagram. Usually, after a brief moment, these electrons will abruptly lose at least some of their extra energy. When they do this, we say they ìdropî in energy level. In this process, they can emit one single electromagnetic photon ñ sort of a ìparticleî composed of e-m waves. The energy of the emitted photon must exactly equal the amount of energy that is lost by the electron as it changes its energy level.
When observing hydrogen gas in an electrical discharge tube (in which high-energy electrons are passed through the gas), the hydrogen glows with a colorful light. When the light is examined through a specially designed filter (a spectroscopic ìgratingî), it is possible to identify the wavelengths of the individual electromagnetic waves that make up that glowing light. These include the following wavelengths: 656.3 nm; 486.1 nm; 434.1 nm; 410.2 nm. (1 nm = 1 × 10 ñ9^ m). Letís try to identify which electronic transitions in the hydrogen atom correspond to these different e-m photons. (Here weíll need to recall that the energy of a photon of frequency f is given by E = hf; Planckís constant h has the value h = 6.63 × 10 ñ34^ J-s = 4.14 × 10 ñ15^ eV-s.)
- In the following table, enter for each of these four different wavelengths of light the corresponding frequency of the e-m wave, as well as the color of light of that wavelength. Also find the minimum ìenergy package sizeî for that e-m wave. That would be the lowest energy with which that e-m wavelength could possibly be observed, and it corresponds to the energy of a ìphotonî of that wavelength.
symbol wavelength (nm)
frequency (Hz)
color photon energy (eV)
Hα 656. Hβ 486. Hγ 434. Hδ 410.
- Suppose an electron drops from the n = 2 energy level down to the n = 1 level; in this process it would have to emit an e-m photon. a) Could that photon correspond to one of the four shown here? If so, which one is it?
b) If not, state whether or not the e-m wave corresponding to that photon should be visible. Explain your reasoning. If yes, what color would it be?
- Photons produced by electrons dropping directly into the n = 1 level are said to be part of the ìLymanî series.
a) How many different photons in the Lyman series could be emitted as a result of electronic transitions among the six levels represented on this diagram? (There exist many others besides these.
- What frequency of light could be absorbed if an electron is raised from the n = 1 energy level to the n = 6 energy level? Could this wavelength of e-m wave be observed as visible light?
- If a gamma ray photon struck the hydrogen atom and was absorbed in the process, what do you think would happen to the electron in the hydrogen atom? Would it still be bound to the nucleus of the atom? Explain. (Hint: Refer to the discussion at the bottom of page 1.)
Questions #10-15 refer to this energy level diagram of some unknown atom. The distances between the horizontal lines are proportional to the energy differences between the levels.
n = 4
I J
G H
E F
B
C D
A
n = 3
n = 2
n = 1
- Suppose that transition B corresponds to the energy of a photon of blue light.
Then could transition D represent:
A) emission of infrared light B) emission of ultraviolet light C) absorption of infrared light D) absorption of ultraviolet light Explain your answer:
- Which transition corresponds to the longest wavelength absorption?
- Which transition corresponds to the highest frequency emission?
- Weíll use the notation ∆E 12 to represent the absolute value of (En = 2 ñ En = 1); this is the energy change of an electron that makes a transition between level 1 and level 2. Suppose we are given that ∆E 23 ÷ h = [frequency of yellow light]. Which transition could represent the emission of infrared waves?
- Rank all transitions in order of increasing frequency of emitted photon (but only include transitions corresponding to emission):
lowest frequency _________________________________________highest frequency
- Rank all transitions in order of increasing wavelength of absorbed photon (but only include transitions corresponding to absorption):
smallest wavelength _________________________________________longest wavelength
Explain how you found the answer to this question.
