LAB #7: Emission Spectra and the Electronic Structure of Atoms
Lab 7.2 Emission Spectrum of Selected Elements
BACKGROUND:
In 1913, Niels Bohr used the idea of a quantum of energy (discrete bundle or photon) introduced
by Max Planck and the model of the nuclear atom introduced by Ernest Rutherford in 1911 to develop
an explanation for the behavior of an electron around a hydrogen atom. He explained that the
characteristic lines of the hydrogen atom spectrum were the result of electron movement from one
energy level to another. The distinct energy levels could be calculated and were called quantum
numbers.
Although Bohr's model could successfully explain the emission spectrum of hydrogen it could
not be used to interpret the spectra lines atoms with more than one electron. The concept of different
quantum levels remained, but additional calculations were necessary.
Our present theory of atomic structure states that electrons in atoms can possess only discrete
energy values. The energy levels are characterized by an integer n called the principal quantum
number. The electrons in atoms in their ground states exist in their lowest energy levels. If energy is
supplied by an electrical discharge or heat, the atoms will absorb energy and electrons will enter higher
energy levels. The energy levels may be compared to rungs on a ladder. The rungs are discrete rather
than continuous; it is not possible to be in between rungs. Atoms having extra energy are called excited
atoms. They can lose some of this extra energy in discrete amounts as the electrons drop back down.
The energy is released as waves with distinct wavelengths, that correspond to the amount of energy
released. The more energy released the shorter the wavelength. These discrete bundles of energy are
called quanta or photons. A quantum of energy is the smallest level of energy that can be transferred to
or from an atom or molecule.
Photons possess energy proportional to the frequency. The
proportionality constant is called Planck's constant. (E = hƒ)
The bright line spectra observed when atoms of gaseous elements are excited by an electric
discharge show sharp lines of colored light with regions of no light. This provides convincing evidence
that the electrons in atoms do not have a continuous range of energies but only certain discrete values of
energy. It was discovered that each atom has distinct characteristic spectrum that could be used to
identify that element.
Purpose: To use a simple spectroscope to observe the bright line spectra of different elements. To
relate the bright line spectra to the energy of the electrons of excited atoms.
Pre-Lab Questions:
1. Describe the photoelectric effect. Explain why this cannot be explained by the wave model.
Procedures:
Observe and record the bright line spectra of the different elements
Label the ends of each spectrum. Which end is high energy? low energy?
Which band of colored light is the result of the fall from the highest level?
lowest level?
Which band of colored light is the lowest frequency? highest frequency?
Which colored band of light has the longest waves? shortest waves?
Analysis:
**2. Given the following data for hydrogen.
color
wavelength (nm)
purple
410.3
violet
434.2
turquoise
486.3
red
656.4
Calculate the frequency and energy of each observed line of light. Tell the story of how this bright
line spectrum was produced. Be detailed and specific about the movement of electrons between
energy levels as energy is absorbed and released. Explain the different energy levels in terms of
their distance from the nucleus and the potential energy and stability of an electron in those levels.
Explain the lines produced in terms of the wavelengths, frequency and energy.
HONORS: Use Bohr’s Equation.
Calculate the energy released as an electron in hydrogen falls from n = 3 to n = 2.
Which observed line in the spectrum is the result of this transition?
Calculate the energy released as an electron in hydrogen falls from n = 5 to n = 2.
Which observed line in the spectrum is the result of this transition?
Calculate the energy released as an electron in hydrogen falls from n = 3 to n = 1.
Explain why this line is not observed?
3. Given the following data for mercury.
Calculate the frequency and energy of each observed line of light.
color
wavelength (nm)
violet
404.7
blue
435.8
green
546.1
yellow
579.0
Describe what happens to the atoms and their electrons within the tube when electric current passes
through the tube. Explain how the observed bright line spectrum is produced. Discuss electron
movement between energy levels when energy is absorbed and then when energy is released.
4. Explain why the emission spectrum cannot be explained by the wave model.
5. What was the significant difference between Rydberg (1885) and Bohr’s (1911) equations?
6. Draw the electron configuration for each atom observed in the emission spectrum.