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Element identification from emission spectra

Learn how to use the light analysis feature by performing an emission spectra experiment.

Scientists use models to make predictions, solve problems, and test claims. As technology improves, new evidence often either confirms or contradicts previous models. Today's atomic model is the product of centuries of refinement. In the early 1900s, Niels Bohr proposed an atomic model based on the work of many other scientists, including Ernest Rutherford and Max Planck. Bohr was familiar with hydrogen's line emission spectrum and realized that the energy of emitted light must be related to an atomic structure that included energy levels. Bohr hypothesized that emitted light was the result of an electron jumping from one energy level to another. He proposed that the energy of an emitted photon is equal to the energy difference between the ground state and excited state. The distinct pattern of spectral lines unique to each atom could be related to energy levels.

Bohr's model was crucial in the development of modern atomic theory because it provided the idea that electrons are organized according to energy levels outside of an atom's nucleus. Unfortunately, the model failed to fully explain the emission spectra of elements beyond hydrogen. The value of studying Bohr's model lies in how it conveys a complex idea in a simple representation.

A spectrometer is a versatile tool that allows the observer to study the line emission spectrum of atoms with energized electrons. Spectrometers allow scientists to identify elements present billions of light years away from Earth, as well as provide evidence for the quantum model of the atom.

Materials required

Safety

During the experiment, follow these important safety precautions in addition to your regular classroom procedures:

  • Wear gloves when handling the spectrum tube.
  • Do NOT touch the high-voltage spectrum tube or power supply while they are active.
  • Use caution around the high-voltage spectrum tube power supply at all times.

Setup

  1. Start Chemvue, then connect the Wireless Spectrometer to the program. The Spectrometry Module should open automatically.
  2. Place the rectangular end of the Fiber Optic Cable into the cuvette opening of the spectrometer. Make sure the arrows on the side of the housing are aligned so that they point in the direction shown in the illustration above.
  3. Use a multi clamp to secure the probe (rounded end) of the Fiber Optic Cable on a ring stand, as shown above.
  4. Place the end of the probe a distance of 2 cm or less from the gas tube. Adjust the probe to point towards the tube. Do not allow the probe to directly touch the tube.

Collect data

  1. Switch to the light spectra page using the Navigation Bar.
  2. Connect the spectrum tube power supply to the gas spectrum tube, then turn on the power supply.
  3. Select Start to begin collecting data. If the wavelength reading is too intense or too weak, adjust the probe angle and distance from the gas spectrum tube until it is within an appropriate range. You can also use the sliders on the left of the screen to adjust the integration time, number of scans to average, or smoothing.
  4. Once you have a stable reading, select Stop and turn off the power supply.

Analyze data

  1. Select Scale to Fit to ensure all data is clearly visible on the display.

  2. Select Coordinates , then drag the coordinates data target to find the wavelengths (in nm) of each distinct peak. Record each wavelength in a table like the one below.

    Color of Photon Wavelength (nm) Energy (J)
       
       
       
       
       

  3. Use the arrows at the bottom-right of the screen to view each reference spectrum. Which element is contained in the gas tube?

  4. Use the wavelength of the identified peaks to solve for the energy of emitted photons using the following equation:

Questions to consider

  1. Which color has higher energy: red or blue?
  2. Is wavelength directly or indirectly related to energy?
  3. Is frequency directly or indirectly related to energy?

  4. Niels Bohr used mathematical relationships to predict the structure of the atom. He assumed lower energy electrons were closest to the nucleus and higher energy electrons were farther from the nucleus, as shown below.

    1. Suppose an electron in its ground state (n = 1) absorbs energy, transitions to an excited state (n > 1), and then returns to ground while releasing a photon. Which of the paths in the above image will require less energy to complete: A or B?
    2. Which light color is more likely to result from the path you identified as requiring less energy: red or blue?

  5. Astronomers use high-powered spectrometers to analyze light throughout space. How is it possible for an element to have the same line emission pattern every time it is energized, whether the element is in outer space or inside a gas spectrum tube in a classroom?