Atomic spectra and the Bohr model

Students view continuous spectra from incandescent and fluorescent lights and line spectra of selected elements. Students relate energy to frequency of light seen in the spectra. The presence of only certain lines in atomic spectra is related to Bohr's model of the atom. In a second experiment, students determine electron energies in the hydrogen atom.

A lesson plan for grades 9–12 Science

By Lisa Bacon

Learn more

• The Bohr Model Diagrams and information on the Bohr model.

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Related topics

• Learn more about Bohr model, atomic structure, chemistry, electromagnetic spectrum, electron energy, energy, hands-on, hydrogen, labs, and science.

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Learning outcomes

• Students will contrast continuous spectra with line spectra.
• Students will explain why the emission spectra of elements are line spectra rather than continuous spectra.
• Students will match colors of light in a spectrum with the energy transitions occurring in the atom (for the hydrogen atom only).

Teacher planning

Time required for lesson

2 hours

Materials/resources

Physical resources

• Chemistry Reference Tables (electromagnetic spectrum and Bohr model), PDF from the North Carolina Department of Public Instruction, one for each student
• Card-mounted diffraction gratings (we used Science Kit and Boreal 65681-00) and/or spectroscopes (we used Science Kit and Boreal 16525-00)
• Spectrum tubes of various elements (including hydrogen and mercury)
• Spectrum tube power supply
• Incandescent light bulb
• Fluorescent light bulb
• Hot plate (optional)

Classroom environment

• Lab or classroom space with sufficient outlets for spectrum tube power supplies
• Room that can be darkened for viewing line spectra

Technology resources

• A computer with a color monitor and access to the internet
• Acces to websites that provide some background information listed in the On the Web section in the sidebar

Pre-activities

Students should summarize the development of atomic models (Dalton’s solid spheres, Thomson’s “plum pudding” or “chocolate chip cookie” model, Rutherford’s nuclear model). At this point, the model of the atom includes the location of electrons outside the nucleus but not the idea that electrons are found only at certain distances from the nucleus.

The teacher should understand the difference between blackbody radiation by a hot object (which emits radiation with a mix of frequencies related to the object’s temperature) and the emission spectra of atoms (which are related to movement of electrons from higher energy levels to lower energy levels). It is not necessary to go into detail with students. Simply explain to students that energy emitted by a hot object is not emitted because of electron transitions.

Activities

Note: This plan was implemented on the block schedule with ninety minutes per class period.

Day One

Activity One

Twenty minutes at the beginning of the period was used to review history of atomic theory (early philosophies of matter of Aristotle and Democritus, Dalton, Thomson, Millikan, Rutherford). Emphasis was on how atomic models change as more experimental evidence is gathered.

Activity Two

• Introduce the idea that energy of radiation is related to frequency of radiation (about five minutes to show, discuss electromagnetic spectrum, and discuss what would be done in lab).
• Either turn on a hot plate or ask students to imagine a hot plate (or electric heating element on a stove). Lead students through a discussion of how they can sense heat before the heating element glows, then the element glows red.
• Lead students to the idea that red light is more energetic than the infrared radiation they feel as heat. Point out the infrared portion of the electromagnetic spectrum and then the visible portion. Infrared radiation is less energetic and lower frequency than red light, which has a lower frequency and less energy than orange light, etc.

Activity Three

• Demonstrate continuous spectra (about five–six minutes).
• Hand out the diffraction gratings and Spectra of Elements Lab: Fluorescent lights versus incandescent lights and have students look at the room lights (fluorescent light bulbs). Remind students that white light is a mix of frequencies of light (the visible portion of the electromagnetic spectrum) and, therefore, a mix of energies.
• Have students record whether the spectrum contains more red or more blue/violet light. Students will look at the spectrum of an incandescent light bulb and record whether that spectrum contains more red or more blue/violet light.
• Students will also look at two fabric swatches: one looks either red or maroon depending on the type of light; the other looks either blue or purple.

Activity Four

• Display line spectra of several elements using spectrum tubes (about forty minutes because students swap samples between groups).

Caution: spectral tube power supplies contain transformers and involve high voltages. Students should be certain power supplies are off when samples are being inserted or removed. The glass tubes will also get very hot if the sample is on for several minutes.

• Students will record the spectral lines on their lab papers. Emphasize that the fact that only certain lines appear in an element’s spectrum is related to only certain energies being involved.
• If a spectrum tube of mercury vapor is available, students will also be able to use a spectroscope to look at a fluorescent light bulb and see brighter lines that match mercury’s line spectrum. (This part took an additional ten minutes because students had to share the spectroscopes and practice aiming the slit in the spectroscope properly.)

Activity Five

Summarize lab results, emphasizing the point that the line spectra support Bohr’s model of the atom. Use the mercury spectrum and fluorescent light as an example of spectra serving as “atomic fingerprints” (about eight–ten minutes).

Day Two

Activity Six

Calculate energies of transition for the hydrogen atom.

• Hand out the Electron Energies in the Hydrogen Atom lab.
• Students will use spectroscopes to measure the wavelengths of light present in the spectrum of the hydrogen atom. If spectroscopes are not available, these values may be found elsewhere.
• The frequencies and energies associated with the lines in the spectrum will be calculated and related to energy levels in the hydrogen atom (about thirty minutes). If quantitative spectroscopes are not available, students may use data from the previous day if hydrogen was one of the samples. Recording the spectrum should take about five minutes.
• Calculations and comparing the spectrum with the Bohr model on state reference tables should take about twenty-five minutes.

Activity Seven

Assess student learning. Some assessment will be based on lab results. Students will also answer four multiple-choice questions (about five minutes).

Assessment

The assessment includes questions in the conclusions of the labs and multiple-choice questions. The multiple-choice questions were answered immediately after the lab on measuring electron energies in the hydrogen atom.

Supplemental information

Bibliography for background information for teachers:

Glencoe Chemistry: Concepts and Applications, Section 2.1

Note: I used online material from the NOW workshop on Light and Visual Phenomena, but this website is not available to persons not enrolled in the course. For those who lack access to this, the links provide content on similar topics.

Comments

In the lab Spectra of Elements, a revised copy is available that does not include the part on contrasting the spectra of incandescent and fluorescent lights, nor does it include looking at fabric swatches under the two types of light. I found this part to distract from the main point of the lab: how line spectra support Bohr’s model of the atom.

Students found it easy to associate higher frequency with higher energy. They found more difficult and only partially mastered the inverse relationship between wavelength and energy. Students had little difficulty in understanding that line spectra support the idea of specific energy levels, but most were unable to answer this lab question: “Notice the energy change for an electron moving in two energy levels (from energy level 4 to energy level 2) compared to the energy change for an electron moving in only one energy level (from energy level 3 to energy level 1). Why doesn’t the electron lose twice as much energy when it moves in two levels as it loses when it moves only one energy level?”

Only some students were comfortable calculating frequency from wavelength many did not remember the equation c = wf (wavelength × frequency) from physical science, but all students were able to calculate energy from frequency using the equation E = hv.

Students enjoyed viewing the atomic spectra with diffraction gratings. They liked the colors seen both with the unaided eye and the lines in the spectra. Some students had physical difficulty aiming the spectroscopes properly and were frustrated when trying to identify the gaseous element present in fluorescent light bulbs. A few students preferred to have one correct answer and were distracted over whether a spectral line was red or orange, blue or violet, etc. Students were less enthusiastic about the second activity (calculating energy of electron transitions from spectral data) because it involved a small amount of hands-on activity for a relatively large amount of calculations.

Viewing atomic spectra to support the Bohr model worked well. Students enjoyed viewing the excited gas samples and the spectra because they liked the colors. The first lab on atomic spectra has been modified to leave out viewing fabric swatches under fluorescent and incandescent lights. It was difficult to find appropriate swatches, and a few students saw no difference in hue between any of the samples (“both look blue” or “both look red” were sample comments). In addition, that part of the activity is not related to atomic spectra and the Bohr model, the core concept of the lab. On the second lab (Atomic Spectra And Electron Energies) I would have students estimate wavelength using the spectrometers rather than giving the wavelengths in the lab handout.