Summary

This activity is designed for students to build a scientific argument about the relationship between energy and spectral lines by exploring how light interacts with atoms. In the process, students will examine proposed models of the hydrogen atom and use collected data to analyze the proposed models. They will then select one of the models and write a scientific argument to support their choice. Students will then review additional data to support and/or refute their selection. Based on their analysis, students will revise their selected model and construct a new argument to support their revisions.

Grade Level

High School

NGSS Alignment

This activity will help prepare your students to meet the performance expectations in the following standards:

• HS-PS4-4: Evaluate the validity and reliability of claims in published materials of the effects that different frequencies of electromagnetic radiation have when absorbed by matter.
• HS-PS4-5: Communicate technical information about how some technological devices use the principles of wave behavior and wave interactions with matter to transmit and capture information and energy.
• HS-ESS1-2: Construct an explanation of the Big Bang theory based on astronomical evidence of light spectra, motion of distant galaxies, and composition of matter in the universe.
• Science & Engineering Practices
  • Developing and using models
  • Analyzing and interpreting data
  • Constructing explanations (for science) and designing solutions (for engineering)
  • Engaging in argument from evidence
  • Obtaining, evaluating, and communicating information
• Crosscutting Concepts
  • Patterns
  • Cause and effect: Mechanism and explanation.
  • Energy and matter: Flows, cycles, and conservation.

AP Chemistry Curriculum Framework

This lesson plan supports the following units, topic(s) and learning objectives:

• Unit 1: Atomic Structure and Properties
  • Topic 1.5: Atomic Structure and Electron Configuration
    • SAP-1.A: Represent the electron configuration of an element or ions of an element using the Aufbau principle.
• Unit 3: Intermolecular Forces and Properties
  • Topic 3.11: Spectroscopy and the Electromagnetic Spectrum
    • SAP-8.A: Explain the relationship between a region of the electromagnetic spectrum and the types of molecular or electronic transitions associated with that region.
  • Topic 3.12: Photoelectric Effect
    • SAP-8.B: Explain the properties of an absorbed or emitted photon in relationship to an electronic transition in an atom or molecule.

Common Core State Standards

• RST .11-12.7: Integrate and evaluate multiple sources of information presented in diverse formats and media (e.g., quantitative data, video, multimedia) in order to address a question or solve a problem.
• RST .11-12.8: Evaluate the hypotheses, data, analysis, and conclusions in a science or technical text, verifying the data when possible and corroborating or challenging conclusions with other sources of information.

Objectives

By the end of this lesson, students should be able to

• Analyze data from the emission spectra of a hydrogen atom to refine their model of electrons within the atom.
• Analyze student models of electrons within the atom to create and revise a scientific argument of the location of electrons within the atom.

Chemistry Topics

This lesson supports students’ understanding of

• Energy
• Atomic Emission Spectrum
• Atomic Structure
• Spectroscopy
• Argumentation
• Model Based Reasoning
• Inferences

Time

Teacher Preparation: 20-30 minutes

Lesson: 100 minutes

Materials

• Hydrogen discharge tube
• Power source
• Diffraction gratings or spectroscopes
• Colored pencils

Safety

• Always wear safety goggles when handling chemicals in the lab.
• Students should wash their hands thoroughly before leaving the lab.
• Follow the teacher’s instructions for cleanup and disposal of materials.

Teacher Notes

• Background Information for Teachers:  Since Democritus described the first idea of “atoms” in Ancient Greece, scientists have been investigating the underlying structure of the atom. This task is particularly challenging due to the incredibly small size of atoms and subatomic particles.  Since direct observations cannot be made, scientists must make inferences about the properties and structure of the atom based upon observations of energy.  Thomson, for example, used his observations of his famous cathode ray experiments to make the connection that the negative charge of the cathode ray was caused by the negatively charged electrons within the atom.  The next step of the puzzle was found through Rutherford’s alpha radiation experiments.  His results indicated the electrons were found on the outside of the atom, contradicting the previous “plum-pudding model.”  At this point, Niels Bohr used his observations of atomic light emissions to further revise atomic models.
  
  In his experiments, Bohr knew that light acts both as a wave and a particle.  When observations are made of continuous light spectra, like natural sunlight or incandescent light bulbs, scientists can see every wavelength of visible light.  Discontinuous spectra contain only spectral lines of specific wavelengths.  Such spectra are produced by elements when sufficient energy is added. Each element releases a characteristic spectrum; these spectral lines observed from an elements emitting light is called an emission spectrum. Since these spectra are unique to each element, scientists can use the observed spectral lines to identify an unknown sample, even to identify the elemental composition of distant stars. Bohr applied his knowledge of spectral lines to establish his model of the atom. When Bohr analyzed the hydrogen atom, which contains only a single electron, he observed only four spectral lines of visible light. Since each color is determined by a specific wavelength of light, he calculated the exact energies being emitted by the hydrogen atom.  Bohr then inferred that the electrons surrounding the nucleus within an atom can only be in specific energy levels.  Since only four specific colors are visible with hydrogen, instead of random colors or the entire visible light spectrum, Bohr argued there must be some underlying structure which determines the colors emitted. In his model, the electrons were held at these energy levels due to the electrostatic attraction between the positively charged nucleus and the negatively charged electrons. If sufficient energy is added to the atom, the electrons can absorb the additional energy. This extra energy allows the electron to overcome the attractive electrostatic force of the nucleus and move further away to a higher energy level. At these higher energy levels, electrons are said to be “excited.” This arrangement is only temporary, however, and the electron eventually falls back to a lower energy level. Which level it falls to determines which type of electromagnetic radiation is released. When electrons release radiation of high levels of energy, it is considered ultraviolet radiation which is not visible by the human eye. Likewise small emissions of energy release infrared radiation which has too little energy to be seen by the human eye. Only when electrons fall between particular energy levels can visible light be released.
  
  Bohr’s model was the first to place electrons in specific places outside the nucleus.  Additional research has further revised his model. Instead of electrons moving around specific “orbits” around the nucleus, subsequent research implies the electrons are found within particular regions around the nucleus in which electrons have a high probability of being found.

• Student Prerequisite Information:
  • Parts of the atom (proton, electrons, neutrons) and their properties.
  • Properties of a wave: wavelength, amplitude, crest, trough (students should already have an understanding of the inverse relationship between wavelength and frequency).
  • Conservation of energy.
  • Experimental design and control of variables.
  • A basic understanding of direct and inverse relationships (although the terminology is not necessary).
  • Light behaves like a particle and a wave.
  • Historical models of the atom: billiard model, plum pudding model, and planetary model.
  • Experience using and revising scientific models.