Materials Included In Kit

Additional Materials Required

Safety Precautions

Power supplies and spectrum tubes operate at very high voltages and can produce a large electric shock. Do not touch the ends of the tube when the power supply is on. Do not touch the contacts on the transformer when the power is on. Always turn off the power supply before inserting, removing or adjusting the position of the spectrum tube. Spectrum tubes typically emit ultraviolet radiation, which is damaging to the eyes. Wear safety glasses or goggles that offer UV protection by filtering UV radiation. To extend the life of the tubes, do not leave the tubes on for more than 30–45 sec at a time. Turn the power on and off as needed to complete the observations. Spectrum tubes may get very hot. Never touch a spectrum tube when the power is on. After turning off the power, allow the tube to cool before removing it from the power supply.

Disposal

None required. Carefully store all equipment for future use.

Lab Hints

• Using a spectroscope is an interesting and fun activity. Instructors may be discouraged from doing an “atomic spectra” lab because of the expense involved in purchasing power supplies. Many inexpensive alternatives to spectrum tubes are readily available. “Neon” novelty lamps (available at many party stores and discount stores) are good sources of bright line emission spectra of mercury (blue) and neon (red).
• Check with the physics department for power supplies, spectrum tubes and spectroscopes.
• Most of the major equipment needed for this lab may be shared by several groups of students. For example, if three power supplies and advanced spectroscopes are available, set up one each with a hydrogen, mercury, and neon discharge tube. Let students rotate through a variety of stations (incandescent light, fluorescent light, novelty lamps and gas discharge tubes).
• Having one advanced or handheld spectroscope per student group is ideal but not necessary for this activity. If practical, allow students to take handheld spectroscopes home to investigate a variety of lights in their homes and communities. The best handheld spectroscopes contain a wavelength guide or scale to estimate the wavelengths of light. Note: The terms spectroscope and spectrometer may be used interchangeably.
• The “Spectrum Analysis Chart” available from Flinn Scientific (Catalog No. AP8676) is a poster-size, full-color chart that shows the bright line emission spectra of ten elements.
• Take advantage of Internet resources to locate full-color reference spectra for different gases. The following websites (accessed July 2014) show a wide selection of atomic spectra:
  http://astro.u-strasbg.fr/~koppen/discharge
  http://hyperphysics-phy-astr.gsu.edu/hbase/hframe.html (see Quantum Physics section).
• The color palette on a computer may be used to design and print simulated bright-line emission spectra of the elements observed in this exercise. The color palette may be manipulated to give an infinite variety of color shades in a range of brightness to match almost any observation!
• Depending on the quality and precision of spectroscopes available for student use, instructors may ask students to calibrate the spectroscopes using mercury spectrum lines as a reference. The precise wavelengths for the mercury lines are:
  • Violet, 404.7 nm
  • Blue, 435.8 nm
  • Green, 546.1 nm
  • Yellow, 579.0 nm
• This experiment may be coupled with a flame test lab to compare the characteristic observed colors of different metal ion (salt) flames and their emission spectra. Vibrant color flames are obtained with lithium (red), strontium (orange), sodium (yellow) and copper (green) salts. We recommend using the metal chlorides. The following procedure gives excellent results and is very safe. Soak wooden splints halfway in distilled water for several hours. Dip the soaked end of a wooden splint in a solid metal salt placed in a weighing dish and then place the splint in a laboratory burner flame.
• Using a spectroscope and a digital camera, you can also take pictures of emission spectra right in your classroom or lab. Set up an advanced spectroscope, spectrum tube and a lamp and check to be sure you can see the spectrum through the eyepiece. Then put the camera up to the eyepiece, and make sure the flash is off. Hold it steady and take the picture. It’s that easy!
• A modern advanced spectroscope such as the Flinn model AP5716 has two main “arms” consisting of a long tube with a slit through which light is focused on a lens and then passes through a prism. The resulting diffracted light is focused by another lens and viewed through a telescope eyepiece. A third shorter “arm” consists of an illuminated wavelength scale (see Figure 2).
  {14026_Hints_Figure_2_Spectroscope design and function}

Answers to Prelab Questions

1. Read the Procedure and the Safety Precautions. What hazards are associated with the use of spectrum tubes? What precautions must be followed to avoid these hazards?

   Spectrum tubes operate at high voltages and can produce electric shock. Turn off the power before inserting or removing a spectrum tube. Wear UV-protective safety glasses to filter UV radiation.

2. What aspect of Bohr’s original model of electron structure is still present in the currently accepted theory of electron structure?

   The current theory of the electron structure of atoms retains the idea of quantized energy levels originally proposed by Niels Bohr to account for the line spectrum of hydrogen.

3. What aspect of Bohr’s original model of electron structure is no longer considered valid in the currently accepted theory of electron structure?

   The current theory of the electron structure of atoms rejects the idea that the path of the electron is restricted to certain specific orbits around the nucleus. Modern theory states that it is impossible to describe the specific location of an electron as a particle at any given time.

4. Assume that a certain atom has a total of four possible energy levels and that an electron can “jump” up or down between any of these energy levels. Draw a model of these energy levels similar to Figure 1 and use it to predict how many different spectral lines should be observed in the emission spectrum of the element.
   {14026_PreLabAnswers_Figure_3}
   Six electron transitions are possible among these four energy levels. Six spectral lines should be observed in the emission spectrum.
5. Calculate the energy of red and green light having approximate wavelengths of 640 nm and 495 nm, respectively. Which light has the higher energy?

   Using Planck’s law, E = hc/λ, the energy of red light is

   {14026_PreLabAnswers_Equation_2}
   The energy of green light is 4.01 x 10^–19 J. Green light has higher energy.

Sample Data

Spectrum Table*

*Colored lines to be added by students using the results shown in the data table.

Answers to Questions

1. According to Equation 1 in the Background section, the energy of light (ΔE) is inversely proportional to its wavelength (λ)—as the wavelength increases, its energy decreases. Based on the spectrum observed for incandescent white light, rank the colors in the visible spectrum from highest energy to lowest energy.

   Highest energy to lowest energy: Violet > blue > green > yellow > orange > red.

2. Do all of the colors of light in the visible spectrum span about the same wavelength range—that is, do the bands of color appear equally wide or narrow?

   Different colors of light have different wavelength widths. Red and violet span the largest range of wavelengths while the yellow band appears quite narrow.

3. What color of light in the visible spectrum appears brightest? Does this mean that it is the highest energy light?

   Yellow appears the brightest. The brightness of the light is not related to the energy of the light. Yellow is the most intense light because that is where the incandescent lightbulb has its peak intensity. Yellow is also where our eyes have the greatest sensitivity.

4. Using Equation 1, calculate the energy (ΔE) corresponding to each line in the observed atomic emission spectrum of hydrogen.
   {14026_Answers_Equation_3}
   Note: These calculations are for the observed lines in the data table. The actual wavelengths of light for the hydrgen spectrum are 410, 434, 486 and 656 nm.
5. As shown in Figure 1, the visible emission spectrum of hydrogen is due to transitions from excited energy levels down to the second principal energy level (n = 2). Thus, the highest energy violet line is due to the transition from n = 6 to n = 2, and the lowest energy red line is due to the transition from n = 3 to n = 2. Enter the energy values from Question 4 from highest to lowest in the following table and fill in the missing entries.
   {14026_Answers_Table_3}
6. Plot the energy of each line versus
   {14026_Answers_Equation_4}
   on the following graph and draw a trendline through the points. What does the shape of the trendline tell you?
   {14026_Answers_Figure_4}
   The graph suggests that the energy of a principal energy level is inversely proportional to n².
7. What is unique about the spectrum obtained for a fluorescent light? What element is used in fluorescent light fixtures? The fluorescent light exhibited a bright line spectrum superimposed on the continuous visible spectrum. The bright line spectrum corresponded closely to that of mercury, suggesting that mercury gas is used in fluorescent light fixtures.
8. (Optional) Discuss any interesting or unique features of other types of light sources that were examined. Is it possible to identify the gases used in other light sources based on their emission spectra? Answers will vary. It was interesting to discover that a blue-colored novelty lamp advertised as a “neon” lamp did not contain neon. The bright line spectrum of the blue “neon” lamp matched that of mercury. (Many blue “neon” lights also contain argon.) Only red neon signs actually contain neon! Other gases, such as mercury and argon, are used to create other colors of neon lights. Many communities have switched to low-pressure sodium lamps for streetlights. The streetlight that was examined had many lines in common with the atomic spectrum of sodium. The streetlight gave a bright line spectrum as opposed to a continuous spectrum.

Introduction

Sunlight passing through a prism produces a rainbow of colors—the visible spectrum. The separation of white light into its component colors occurs when light waves of different wavelengths are bent by different amounts. When a pure atomic gas such as hydrogen or helium is subjected to a high-voltage electrical discharge, light is produced and the gas glows. When this light is passed through a diffraction grating, however, the spectrum it produces is different. Instead of giving the familiar rainbow of colors, the light emitted by the gas gives a series of bright, colored lines. The series of bright lines is called an atomic emission spectrum and is unique to each element.

Concepts

• Atomic emission spectrum
• Quantization of energy
• Planck’s law
• Electron energy levels
• Electron transitions
• Spectroscopy

Background

The phenomenon of atomic spectra has been known since the mid-1800s. Their cause, however, remained unexplained until the structure of the atom and, in particular, its electronic structure, was solved. Rutherford’s discovery of the nucleus of the atom in 1911 answered many questions concerning the structure of the atom. It also raised new questions, including, where are the electrons?

In 1913 Niels Bohr proposed a model of electron structure that would explain the phenomenon of atomic spectra. According to Bohr’s model, an electron is restricted to certain specific orbits around the nucleus of the atom. These orbits differ in their distance from the nucleus and in their energy levels. Electrons that are closer to the nucleus are lower in energy than electrons that are farther away from the nucleus. This idea is called the quantization of energy—electrons can only occupy specific energy levels. They may not have intermediate energy levels between these allowed states. The picture that is often used to describe this idea is the rungs on a ladder. An electron must always be on one of the energy rungs, not between them. An electron may be “excited” or promoted from a lower energy level to a higher energy level by absorbing energy of the appropriate wavelength. Conversely, an electron may be “relaxed” down to a lower energy level from a higher energy level by emitting energy of the appropriate wavelength in the form of a photon.

Bohr’s theory successfully predicted the atomic spectrum of hydrogen. When electrical energy is supplied to hydrogen atoms in a gas discharge tube, also called a spectrum tube, the atoms absorb energy and the electrons are promoted to excited energy levels. Once excited, however, the electrons have a natural tendency to drop back down to a lower energy level by emitting light of the appropriate wavelength and energy. The emitted light for a given transition is observed through a diffraction grating as a bright line in the emission spectrum of hydrogen.

The relationship between the energy of light and its wavelength is given by Planck’s law:

ΔE is the difference in energy between the two energy levels in joules, h is Planck’s constant (h = 6.626 x 10^–34 J • sec), c is the speed of light (c = 2.998 x 10⁸ m/sec), and λ (lambda) is the wavelength of light in meters.

When Bohr calculated the allowed energy levels for the electron in the hydrogen atom, he found that the results correctly predicted the wavelengths of visible light observed in the emission spectrum of hydrogen, which exhibits four bright lines in the visible region (see Figure 1). The Bohr model of electron structure was found to be inadequate for atoms containing more than one electron. However, the idea that only certain stable electron energy levels are allowed has endured. The development of quantum mechanics in the 1920s built on the idea of quantized energy levels and introduced the idea of the wave nature of matter to describe the properties of electrons. According to quantum mechanics, the location of an electron is not restricted to specific orbits but can only be defined in terms of the probability of finding an electron. A system of atomic orbitals was introduced to account for the arrangement of electrons around the nucleus of an atom. An atomic orbital is a region in space where an electron may be found. Atomic orbitals differ in their size, shape, and orientation in space, and also in their energy. The characteristic atomic emission spectrum of an element can be interpreted based on the unique arrangement of atomic orbital energy levels for its atoms.

In this experiment, we will use a special instrument, called a spectroscope, to view the “bright line” emission spectra of different elements and to determine their wavelengths. A spectroscope contains a diffraction grating that separates light into its component wavelengths.

Experiment Overview

The purpose of this experiment is to recognize continuous versus line emission spectra for various sources of light using a spectroscope. The spectroscope will also be used to observe the atomic spectra of different elements in spectrum tubes and to identify the elements that may be present in fluorescent lights, street lamps, novelty “neon” lamps, etc.

Materials

Prelab Questions

1. Read the Procedure and the Safety Precautions. What hazards are associated with the use of spectrum tubes? What precautions must be followed to avoid these hazards?
2. What aspect of Bohr’s original model of electron structure is still present in the currently accepted theory of electron structure?
3. What aspect of Bohr’s original model of electron structure is no longer considered valid in the currently accepted theory of electron structure?
4. Assume that a certain atom has a total of four possible energy levels and that an electron can “jump” up or down between any of these energy levels. Draw a model of these energy levels similar to Figure 1 and use it to predict how many different spectral lines should be observed in the emission spectrum of the element.
5. Calculate the energy of red and green light having approximate wavelengths of 640 nm and 495 nm, respectively. Which light has the higher energy?

Safety Precautions

Power supplies and spectrum tubes operate at very high voltages and can produce a large electric shock. Do not touch the ends of the tube when the power supply is on. Do not touch the contacts on the transformer when the power is on. Always turn off the power supply before inserting, removing or adjusting the position of the spectrum tube. Spectrum tubes typically emit ultraviolet radiation, which is damaging to the eyes. Wear safety glasses or goggles that offer UV protection by filtering UV radiation. To extend the life of the tubes, do not leave the tubes on for more than 30–45 sec at a time. Cycle the power on and off as needed to complete the observations. Spectrum tubes may get very hot. Never touch a spectrum tube when the power is on. After turning off the power, allow the tube to cool before removing it from the power supply.

Procedure

Follow the instructor’s or manufacturer’s directions for use of a specific spectroscope or spectrometer model.

1. Using a spectroscope, handheld spectrometer or diffraction grating, observe the continuous “rainbow” spectrum from an incandescent lightbulb.
2. Observe the colors of light in the visible spectrum and the wavelength range for each color band. Sketch the spectrum of white light using colored pencils in the appropriate wavelength boxes in the spectrum table. Note that the units of wavelength on the spectroscope are nanometers (1 nm = 10^–9 m).
3. (Optional) For optimum viewing of the emission spectra of gas discharge tubes using a handheld spectroscope, stabilize the spectroscope on a ring stand. Set up a ring stand in front of the power supply and attach one ring clamp. Place the spectroscope on the ring clamp and adjust the height of the ring clamp so that the eyepiece on the spectroscope is approximately level with the middle of the gas discharge tube. Attach a second ring clamp on top of the spectroscope so that it will be held firmly in position without moving.
4. With the power OFF, insert the hydrogen spectrum tube between the contacts on the power supply.
5. Move the spectroscope so that it is about 3–5 cm away from the spectrum tube.
6. Turn on the power supply, and observe the atomic emission spectrum of hydrogen. Work with a partner to note the principal features in the hydrogen spectrum.
7. Turn OFF the power supply. Record the following information in the data table for the emission spectrum of hydrogen: the number of lines, their colors and their approximate wavelengths.
8. Using colored pencils, sketch the atomic spectrum of hydrogen in the wavelength boxes in the spectrum table. Turn the power supply on and off, as necessary, to complete the observations in step 7.
9. Check to make sure the power supply is off. When the spectrum tube is cool, remove it from the power supply and insert a mercury spectrum tube.
10. Turn on the power and observe the atomic emission spectrum of mercury.
11. Turn OFF the power supply. Record the emission spectrum of mercury: the number of lines, their colors and their approximate wavelengths.
12. Using colored pencils, sketch the atomic spectrum of mercury in the wavelength boxes in the spectrum table. Turn the power supply on and off, as necessary, to complete the observations in step 11.
13. Repeat steps 9–12 for any other gas spectrum tubes that are available.
14. Using a spectroscope, observe the spectrum of visible light obtained from a fluorescent light. What kind of spectrum is produced? If any bright lines are present, record the number of lines, their colors and their approximate wavelengths.
15. (Optional) Using a handheld spectroscope, observe the emission spectrum of other light sources (e.g., neon signs, streetlights, headlights, novelty lamps). What kind of spectrum is produced? If any bright lines are present, record the number of lines, their colors and their approximate wavelengths.

Student Worksheet PDF

14026_Student1.pdf