AP Chemistry
Unit 1 Lab Activity
Name_________________________________________________
Atomic Emission Spectra and the Bohr Model of the Atom
Purpose
:
To measure the wavelengths of the spectra lines for some common atoms and to correlate the results with actual electron transitions in atoms as predicted by the Rydberg equation
Theory :
Some of the most convincing information verifying the current model of the atom comes from the study of absorption and emission spectra of the atom. Looking at the light through a spectroscope as it is emitted from excited atoms, one can see distinct spectra lines. The wavelengths of these lines are related to differences in certain energy levels of the atoms.
Quantum theory predicts that electrons in an atom can only have certain discrete energies. Atoms in the ground state (having a principal quantum number of n = 1) can be excited to a higher energy state , i.e. n = 5 by an external energy source such as an electrical field or flame. When these excited electrons lose energy, they do so by emitting light of discreet wavelengths and revert to lower energy levels, i.e. n = 4, n = 3, n = 2 or the ground state. The energy emitted is given by the formula:
E = h  where  = frequency, and h = Planck's constant = 6.63 x 10 -34 J s
In terms of wavelength:
E = hc where c = speed of light = 3 x 10 8 m/sec, and  = wavelength.

Niels Bohr established two postulates which account for the apparent stability of the hydrogen atom and the resulting line spectrum observed for the hydrogen atom. According to Bohr an electron in an atom can have only specific energy values. Using Planck's work on quantization of energy to explain the motion of an electron in an atom.
According to Bohr's hypothesis
E = R
H where
R
2
H
= Rydberg constant = - 2.18 x 10 -18 Joules

n and n = quantum number
At the same time Bohr showed that an electron can change energy only by changing from one energy level to another energy level. When an electron goes from a higher energy level to a lower level it emits energy as a photon. The energy of the photon can be obtained by subtracting the energy of the lower level from the higher level.
Thus the energy of an observed spectrum line can be given by the value of Einstein's equation

By measuring the wavelength of the light emitted from a hydrogen discharge tube, the transition to which it corresponds can be easily calculated by substituting the values into the modified Rydberg equation shown below

 E = h  hc
 E = (R
H
) (Z 2 ) ((1/n
1
2 ) – (1/n
2
2 ))
Where R
H
= Rydberg's constant = 2.18 x 10 -18 joules, Z = the nuclear charge, and n = quantum number

AP Chemistry
Unit 1 Lab Activity
Name_________________________________________________
Procedure
1.
Set up the apparatus as shown in the diagram below. The transformer should be placed directly behind the screen so that the glow of the discharge tube is clearly visible through the slit. Position the diffraction grating so that it is exactly 100 cm from the slit. (See y on the diagram below.)
Figure 1
2.
One lab partner will view the emission spectrum through the diffraction grating at the slit.
The other partner will stand behind the transformer and move a ruler along the meter stick.
The student viewing the spectrum should indicate when the pencil is at the point where the image of the spectral line appears to be. Measure this distance (x) to the nearest 0.1 cm.
Repeat this process for each line that is visible in the spectrum. Then refer to the diagram
(Figure 1 above) to make the following calculations:
3.
Find the distance z from the diffraction grating to the image of the spectral line using the
Pythagoran theorem
4.
Find the sine of angle ø where sin ø = x/z
5.
Find the wavelength

in cm where

= d sin ø, d = 1.9 x 10 -4 cm
6.
Find the wavelength in nanometers:

(in nanometers) =

cm x 10 7 nm/cm.
7.
Using the equation E =hc/
 one can calculate the Energy for a particular spectral line as follows:: Since h is in joules, convert the wavelength to meters, use 3.00 x 10 8 ms -1 for the speed of light. Finally energy is usually expressed in kJ mole -1 , so multiply by
Avogadro's number 6.022 x 10 23 and divide by 1000 to convert Joules to kJ
8.
Using the Rydberg equation calculate the electron transitions to which the observed wavelengths correspond. Again multiply your answer by Avogadro's number 6.022 x
10 23 and divide by 1000 to convert joules to kJ. How do your results compare? That is, does the “E” in E=hv (calculated from this experiment), match the “E” in

AP Chemistry
Unit 1 Lab Activity
Name_________________________________________________
Data
Element x
E = R
H
Z 2 (1/n
1
2 - (1/n
2
2 ) ? It should! What are some reasons why it wouldn’t?
Hydrogen

E =hc/

E = R
((1/n
1
H
2
Z 2
) - (1/n
2
2 ))
Discussion
Answer the following on a separate sheet of notebook paper and attach it when you turn in this assignment.
1.
How do your results compare? That is, does the “E” in E=hv (calculated from this experiment), match the “E” in
E = R
H
Z 2 (1/n
12
- (1/n
22
) ? It should! What are some reasons why it wouldn’t?
Answer this question for each element you studied.
2.
Is it possible to identify the gases used in a light source, like the lights in this room, based on their emission spectra? Explain.
3.
Lastly, draw a diagram of a hydrogen atom and label the transitions for it’s single electron. For the transition from the first to second level and back again, what wavelength of light would this electron emit? What amount of energy did it absorb and release? Label these on your diagram and use a color pencil if possible.
4.

AP Chemistry
Unit 1 Lab Activity
Name_________________________________________________
Teaching Notes
Equipment for each station
Transformers and electronic discharge tubes
4 ring stands
4 buret clams diffraction grating slit taken from a spectra scope
Discharge tubes
 Hydrogen
 Mercury
 Neon
 Others as available
 Set up 6 tables with equipment as shown in the diagram Set up one lab station per table
 Student groups rotate around the tables and make measurements for each different discharge tube
 Some images are difficult to see. It is best if the room is dark with all outdoor light blocked out
 The students generally do not have sufficient back ground for the calculations. Doing a sample together will help.