Chapter 1. Emission of Light and Atomic Models

Discussion

At the end of this activity you should be able to:

• Explain what photons are, and be able to calculate their energies given either their frequency or wavelength.
• Order the common kinds of radiation (UV, visible light, infrared) in the electromagnetic spectrum according to their wavelengths or energy.
• Calculate the wavelength of electromagnetic radiation given its frequency or its frequency given its wavelength.
• Be able to compare and contrast different models for the hydrogen atom.
• Distinguish between conceptual and physical models.
• Discuss how the emission spectrum of hydrogen gas was used to support (or refute) a given atomic model.
• Explain how electronic transitions between quantized energy levels are responsible for the emission of photons of light with different energies.
• Collect emission data and apply atomic models to explain the electronic structure of elements like hydrogen and helium.

These learning objectives are assessed in the post-lab Carmen quiz.

There is no laboratory report.

Introduction

When light from the sun or white light is passed through a prism it produces a “rainbow” of different colors. This has been known for centuries, with Isaac Newton making significant advances in the area of “Opticks” in the 1700s by investigating how light is reflected, refracted, dispersed, etc.

However, something very different occurs when light from a single element is emitted and passed through a prism. In this case only a small number of discrete lines are observed. Even more remarkably, the pattern of these lines is a defining characteristic of each element. What happens when light is emitted from an element like hydrogen? As you will see in this investigation, the answering of this question has led to profound insights into the electronic structure of the atom and, in turn, our modern understanding of chemistry.

The emission of photons (particles of light) from atoms is thought to occur in the following way: First, the atom absorbs energy and an electron moves to a higher energy level, or “excited state." When the electron eventually returns to a lower energy state energy is released. The specific amount of energy released corresponds to the difference between the energy levels. When this energy difference corresponds to visible light then our eyes can observe the photons. It is also possible the emitted photons will have energies we cannot see, such as in the infrared or ultraviolet regions of the electromagnetic spectrum.

Atoms can absorb energy in various forms (heat, electrical, radiant) and subsequently emit photons. Today, in this investigation, you will complete activities that involve radiant energy (photons of light). In the experiment, you will return to this topic and investigate the role of heat and electrical excitations. In each case the emission of photons is similar, although different energy levels and different elements may be investigated.

Total lab time~2 hours

All in-lab activities may be discussed with classmates and/or completed in a small group. The teaching assistant will help manage your time.

Laboratory Instrumentation

The instrument used to view line spectra is a spectroscope. Photons emitted from the atoms in their excited states strike the prism that separates them based on their particular wavelength. The resulting spectrum is superimposed on a scale and is viewed through the eyepiece. Note: Only the visible light portion of the spectrum is observed since the detector is the human eye.

Scientific Models

What is a scientific model? Sometimes scientists construct physical models that represent an object on a different scale. Models of an amoeba or bacterium (things very small made larger) or the solar system (something very large made smaller) are examples of physical models.

Scientists also devise conceptual models. A weather map is an example of a conceptual model. The weather map communicates some aspects of the weather, e.g., the temperature, areas of high and low pressure, precipitation, etc. It allows one to make predictions (will it rain in Columbus tomorrow?) and makes sense of a complex system. It is not, however, a physical representation.

Atomic models are conceptual. They aim to communicate different characteristics and make predictions for something that cannot be directly observed.

Part 1. Computer Simulation—Models of the Hydrogen Atom

One of the most tantalizing puzzles at the beginning of the 1900s was how to describe the atom. In Chapter 2 you learned about subatomic particles like protons, neutrons and electrons.

With this activity you will continue your own exploration. This begins by first observing a simulated light spectrum of hydrogen. You will then look at spectra predicted by different models of the atom. The models you will test include:

• John Dalton’s Billiard Ball model
• J.J. Thomson’s Plum Pudding model
• Ernest Rutherford’s Classical Solar System model
• Niels Bohr’s Shell model
• Louis de Broglie’s Electron Wave model
• Erwin Schrodinger’s Quantum Mechanical model

Procedure and Questions

The emission spectrum for hydrogen is a very important experimental observation used to support (or refute) different models of the hydrogen atom. Scientists cannot “look” inside the ?-box of hydrogen and actually see an atom. Instead, they construct theories and conceptual models to explain observations related to the atom. Models can be tested with experiments to learn if the model is useful for predicting how nature behaves

Your task in this simulation is to examine different models of the hydrogen atom. youwill 1) examine the model, 2) see how the model produces an emission spectrum forhydrogen, and 3) decide which model agrees best with the experimental data. It will be helpful to keep the experimental spectrum (the snapshot) available for comparisons.

Dalton’s Billiard Ball Model

10. In Chapter 2 you learned that John Dalton proposed that an atom was simply a very tiny hard ball. To see this simulated, in the top left corner, switch from Experiment to Predict and highlight the Billiard Ball model.

Why do you think the Dalton model does not include electrons?

Thomson’s Plum Pudding Model

Rutherford’s Classical Solar System Model

The Classical Solar System Model has a significant shortcoming. You may need to run the simulation at a slow speed, or pause it, to find it. What is the shortcoming of the solar system model?

In the Bohr model, electrons can exist only at certain energy levels (also called shells), not at any energy levels between them. Shells in this energy diagram go from 1 to 6. An increase in energy level (say from 1→6) can only occur if a photon of incoming light is absorbed. A decrease in energy level (say from 2→1) is accompanied by the emission of a photon as the excited electron releases its excess energy.

What atomic model best agrees with the experimental emission spectrum for atomic hydrogen?

Part 2. Emission Spectra for Hydrogen and Helium

Materials Required

Equipment

• none

Common Equipment

• hydrogen and helium discharge tubes
• 5000 V step-up transformer
• spectroscope

Chemicals

• none

Cautions

The 5000 V step-up transformer and contacts for tubes on boards are an electrical hazard. Always wear your lab goggles, since they will block out almost all harmful UV radiation.

Procedure

Make sure the overhead lights in the balance room are off during observations.

Waste Disposal

None

Data Analysis

Calibration of Spectroscope

The scale readings for each spectrum are simply the location of photons after passing through the prism. Since you know the wavelength of each of these transitions for hydrogen you can use this information to calibrate the spectroscope and convert from scale reading to wavelength for the helium emission spectrum.

Construct a graph of the wavelength, λ (in nm) on the y-axis versus scale reading on the x-axis. Your calibration line should curve smoothly from the upper left to the lower right. Based on the calibration curve and the observed scale readings, calculate the wavelength for each line in the helium emission spectrum and enter them into the table.

Interpreting the Emission Data

The emission of light from hydrogen or helium is not a continuous spectrum. The light from the sun or a white light produces a “rainbow” spectrum, whereas the spectra emitted by a single element consists of a small number of discrete lines. In 1913 Neils Bohr provided an explanation for this phenomenon. In Bohr’s theory (which is investigated in the computer simulation) the atoms of an element emit radiation when electrons undergo a transition from a higher energy state to a lower one. The opposite occurs when light is absorbed; the energy moves the electron from a lower energy level to a higher one.

Hydrogen is the simplest element. An atom of hydrogen has only 1 proton (with a positive charge) and 1 electron (with a negative charge). The proton and electron experience an electrostatic attraction since oppositely charged particles attract each other. Finally, in Bohr’s model, the electron’s energy levels are quantized and stable; the electron does not lose energy and spiral into the nucleus.

In the Bohr model, the energy of an energy state is:

where Z is the number of protons in the nucleus of the atoms of that element, and n is a positive integer identifying the energy state (or “orbit” of the electron).

The change in energy, ΔE, for an electronic transition is therefore:

Common calculations involve converting between energy and frequency (Planck’s relationship E = hv, where h = 6.63 × 10^–34 J s) and between frequency and wavelength (c = λv, c = 3.00 × 10⁸ m/s).

In your notebook calculate the wavelength for each line in the visible region of the emission spectrum of hydrogen. In each case, n_f = 2. Choose energy values higher in energy (n_i > 2) and calculate the ΔE. Then, convert this to wavelength and express your answer in nm. Finally, check your answer by looking in Table 7.1.

Note: Questions similar to this one are frequently on General Chemistry exams!

Show your work for the calculation when n_i = 3 and n_f = 2.

Which statement(s) is/are true regarding the emission of a photon from a hydrogen atom? (Circle all correct answers.)

a. The emission of photons is an endothermic process, so ΔE is positive.

b. The emission of photons is an exothermic process, so ΔE is negative.

c. The wavelength of the n_i = 3 to n_f = 2 transition is shorter than the n_i = 4 to n_f = 2 transition.

d. The energy of the n_i = 3 to n_f = 2 transition is greater than the n_i = 4 to n_f = 2 transition.

After his initial success with the hydrogen atom, Bohr attempted to apply his theory to the helium atom, which has two electrons orbiting a nucleus consisting of 2 protons and 2 neutrons. Despite years (!) of effort and numerous modifications to his theory, Bohr was unable to extend his success beyond the hydrogen atom.

How do the emission spectra for hydrogen and helium differ from each other?

Using what you know about the Bohr model, the structure of hydrogen and helium atoms, etc., propose a model to explain why the spectra differ. There is not a single “right” answer for this question, but your response should be consistent with the available evidence. You may choose to communicate your model with words, drawings, equations, etc.

Activity Completed!