Emission and Absorption Spectroscopy

Purpose:
In this experiment emission and absorption spectra are examined to test the validity of the simple Bohr model for the Hydrogen atom emission properties, to determine the identities of two unknown gases, and to compare and contrast the absorption spectrum of aqueous solutions of compounds to those of gas phase emission spectra.

References:
Chemistry, 4th ed, by J. McMurry and R. Fay. Prentice Hall, 2004, Sections 5.2, 5.9, 5.11

Background:
    As electrons move between the energy levels of an atom they can emit or absorb light energy. If the electron is promoted, i.e. moves from a lower energy level to a higher energy level, the atom must absorb the energy by absorbing the light with energy exactly equal to the difference between the energy levels. If the electron falls from a higher energy lever to a lower energy level it will release energy by emitting light. By absorbing specific wavelengths of light, an electron moves from a lower energy level to a higher energy level.  Since every kind of atom has a different electronic configuration, the wavelengths of light absorbed or emitted by an atom are unique to that element. It is like an atomic fingerprint. The fingerprint recorded by measuring which wavelengths of light an atom absorbs is called the Atomic Absorption Spectrum. The fingerprint recorded by looking at which wavelengths of light an atom emits is called the Atomic Emission Spectrum. For gaseous atoms, these spectra will be a series of discrete "lines" of different colors (energies, wavelengths, frequencies).

    In this experiment you will be looking at the Emission Spectra of several different elements. These elements are present in sealed glass tubes which contain a pure gas of the given element, like H₂, He, Ne, etc. The tubes are placed in a special holder that passes an electrical current from one end of the tube to the other, like in a neon street sign. The electricity passing through the gas excites the electrons in the gaseous atoms and promotes them to higher energy levels. When they return to lower energy states, they emit characteristic wavelengths of light, i.e., an emission spectrum. Within every atom there are many different electronic transitions possible and we will be looking at the "visible" or colored wavelengths. 

You will also examine colored aqueous solutions with ionic compounds dissolved in them.  These compounds contain multiple atoms interacting through both ionic and covalent bonds.  The compounds themselves are also dissolved in water and interacting with the solvent too.  Therefore, the spectra are more complex in appearance.

In order to separate the different wavelengths of visible light, you will use a hand held spectroscope. These spectroscopes separate light into its component wavelengths that can be read on a wavelength scale, typically in nanometers (10^-9m).  You will carefully record the observed intensities and wavelengths to carry out your analysis of each sample considered.

Procedure:
There are six stations located around the lab. At each station you are required to make some observations using the hand held spectroscope. The stations may be done in any order so there should never be too many people at one station. Success in this laboratory exercise is strongly dependent on your ability to make and record accurate observations. At each station make sure that you enter into your lab book, in a table and, either, in words or in pictures (or both), an accurate record of what you observed.  Remember, both intensity (relative brightness) and wavelength will be important data to record.  Use the same spectroscope for the whole experiment.

Station 1: Observe the helium discharge lamp. Record the wavelength (nm) of the 7 lines that make up the He spectrum. The accepted values are: 668, 588, 502, 492, 471, 447, and 403 nm. Use Excel to make a calibration curve for your spectroscope by plotting the wavelengths you measured (observed) against the accepted (actual) values. Determine the linear equation for this line in the form: actual = slope*observed + constant ( y = mx + b). You will use this calibration curve to correct the measurements you make at stations 2-5. 

Station 2: Observe the hydrogen discharge tube. Record the three brightest lines in the visible spectrum. Use your calibration curve to correct your observed wavelengths. (Show your work!)  Determine the values of n_i and n_f for the initial (higher) and final (lower) energy level for each emission wavelength observed.  Consider all possible combinations of n_f and n_i from 1 to 10 to calculate theoretical wavelengths.  This is easily done in an Excel Spreadsheet (i.e.  n_i =10 to n_f =9, n_i =10 to n_f =8, ....... , n_i =2 to  n_f =1) where n must be a whole number and n_i > n_f.  Use the Rydberg Equation (calculate in Excel) provided below to calculate each theoretical wavelength (l).  Note that the Rydberg Equation gives 1/l and you need to find l.
                                                                          where R = 1.097 x 10^-2 nm^-1.

By comparison, assign your observed wavelengths to the closest matching, calculated, theoretical wavelength.  From your corrected wavelengths, also calculate the energy of the transition responsible for each of these three lines. 

Pick 2 stations from stations 3-5:

Station 3-5: Record several wavelengths and relative intensities, (the more the better, but at least three), of the brightest lines of each unknown. Correct your observations using the calibration curve from part one. Use the Emission Lines Excel database located on the 1102 website to determine the identity of your unknowns. The database contains the Element, Wavelength of the Line, and the relative intensity of the line. Using Excel you can sort this data by element, wavelength, or intensity. Using the data you record at this station you should be able to identify the unknown element in the discharge tube. You will do this by finding the element that has all or most of the lines that you recorded. No two elements should have all the same lines and relative intensity patterns, so you should be able to identify a unique choice.

Station 6: Observe the spectra coming from the white light in the overhead projector. Now look through the two colored solutions placed on top of the overhead. Observe the spectra with your spectroscope. Record in your notebook the similarities and differences between the white light and the two colored solutions.  Consider the descriptions provided in the table below in assessing these spectra. 

Table summarized from Thermospectronic “Basic UV/Vis Theory, Concepts and Applications.”