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Electronic Structure of Atoms

Amina Khalifa El-Ashmawy, Ph.D.

Collin College

Department of Chemistry

Introduction: It is the fourth of July and you are enjoying the brilliant colors of the fireworks display. For a

fleeting instant you wonder what gives each burst its distinctive color. On your way home from the

fireworks show you drive by many establishments vying for your attention with their glowing signs,

many of which are bright red. You probably don’t think much of the signs, but, in fact, the

phenomenon going on in the bulbs used in those signs is identical to the fireworks.

The phenomenon we witness everyday is actually quite amazing and has many uses. It was first

discovered in the mid 1600s when Sir Isaac Newton passed light through a prism, producing a

rainbow. Later experiments indicated that light is a form of energy. The colors of a rainbow make up

a small portion of the electromagnetic spectrum, visible light. In the 1850s the spectroscope, an

instrument that bends light at different angles based on the constituent wavelengths, was developed

by the German physicist Gustav Robert Georg Kirchhoff and chemist Robert Wilhelm Bunsen. It

consists of a lens or slit that allows light of parallel rays into the chamber and a diffraction grating.

When sunlight is passed through a prism or spectroscope, a continuous spectrum is produced. That

is, the spectrum of energy contains all the possible wavelengths in that range of energies.

Conversely, if you use a spectroscope to view fireworks you will see a line spectrum that is

comprised of distinct, separate lines that are characteristic for the particular metal that was used in

that shell. Fireworks of different colors are produced by incorporating different metal salts.

Scientists use line spectra to identify elements present in an unknown sample. One of the uses of

most interesting historic significance was the analysis of Napoleon’s hair. Samples of his hair were

analyzed some 150 years after his death only to discover that he was poisoned with arsenic on a

fairly regular basis over a period of time. Line spectra are also used in astronomy to determine the

composition, temperature, motion and density of stars and nebulae.

Electromagnetic Spectrum

Electromagnetic radiation is a combination of orthogonal electric and magnetic waves that travel in

phase in the direction of energy propagation (Figure 1).

Figure 1. Electric and magnetic waves of electromagnetic radiation (Electromagnetic Spectrum)

The distance from crest to crest or any two identical points of the oscillating wave is the wavelength,

λ, expressed in meters. The number of cycles of the wave in a given amount of time is the frequency,

ν. Frequency can be expressed in the following units, all of which are equivalent: cps (cycles per

second), 1/s, s -1

, or Hz (Hertz). Wavelength and frequency are inversely proportional. Additionally,

the energy of the wave, E, is directly proportional to frequency.

c = λ ν, where c is the speed of light, 2.998x10 8 m/s

E = h ν, where h is Plank’s constant, 6.626x10 -34

J•s.

Different regions of the electromagnetic spectrum have different wavelengths and frequencies. At

the low frequency end of the spectrum are radio waves and microwaves. At the high frequency end

are gamma rays and x-rays. The sun’s rays span a wide range of energies; some are visible but most

are not.

Bohr’s Work

Danish physicist Niels Henrik David Bohr calculated, based on the wavelengths of the hydrogen line

spectrum, the allowable distances where hydrogen’s electron can be with respect to the nucleus.

These distances, or orbits (n), were determined based on the idea that when excited or heated the

hydrogen atom absorbs energy and its electron jumps to a distance farther from the nucleus or higher

energy shell. The electron falls back down closer to the nucleus and emits a photon of specific

wavelength. According to the photon wavelength, one can calculate the shells associated with the

electronic transition that gave rise to that observed line in the line spectrum. This is done using the

Rhydberg equation shown below.

where R = 109737 cm -1

nf = final shell of transition

ni = initial shell of transition.

So, we have a basic understanding of electromagnetic radiation and the spectroscope, but how are

these related to fireworks, neon signs, and Napoleon’s hair? When a substance is heated until it

glows, it emits a certain amount of radiant energy that corresponds to particular energies in the

electromagnetic spectrum. Transitions to n = 1 as the final shell fall in the ultraviolet region of the

electromagnetic spectrum and constitute what is known as the Lyman series. Those to n = 2, in the

visible region, constitute the Balmer series. Those to n = 3, in the infrared region, constitute the

Paschen series.

The Problem:

You will heat different metal salts in a flame to witness their distinctive colors. You will then be

given an unknown salt to identify based on your flame test observations. Additionally, you will

observe emission tubes of different elements (H, He, Ne, Ar, Kr) and record their line spectra in your

laboratory notebook. Focusing only on the hydrogen emission spectrum, calculate the transitions that

are causing the different lines observed.

Conducting the Flame Tests: Soak some cotton swabs in deionized water for at least 5 minutes. Pour a very small amount of the

metal salt you wish to test onto a watch glass. Roll one of the soaked cotton swabs in the metal salt

then place in the inner blue cone of a Bunsen burner flame. Record you observations. Place the used

cotton swab in the appropriate waste receptacle. Repeat the previous procedures for each of the

metal salts to be tested using a clean watch glass and a new cotton swab each time.

Thought Questions Before Performing the Lab:

Following are some questions to consider before doing this lab.

What equipment do you need to conduct the flame tests?

What is happening at the atomic level to give rise to the observed energy?

Why does each salt have its own distinctive flame test color?

How are the line spectra of the different elements related to the observed color with the unaided eye?

How many lines should be in the emission spectrum of each element listed above?

Will everyone see all the lines in all the emission spectra of these elements? Explain.

Are the lines seen representative of all the transitions that are taking place in the atom? Explain.

Why are no two emission spectra for different elements ever the same?

What are the wavelengths of the lines in the hydrogen emission spectrum?

Using the Rhydberg equation, calculate the transitions in the H-atom that are taking place to give rise to each line?

Critical Data/Discussion to Include in Your Lab Report:

Succinct and complete list of procedures

All data collected

Answer all questions posed above

Calculations of the transitions taking place in the H-atom that give rise to each of the lines in its emission spectrum

Works cited in the format specified by instructor

Work Cited

The Electromagnetic Spectrum. Retrieved May 22, 2008, from Google Images Web site:

http://images.google.com/imgres?imgurl=http://www.geo.mtu.edu/rs/back/spectrum/e_mag.gif&

imgrefurl=http://www.geo.mtu.edu/rs/back/spectrum/&h=268&w=506&sz=5&hl=en&start=8&u

m=1&tbnid=BLYdvKWr6t5XfM:&tbnh=69&tbnw=131&prev=/images%3Fq%3Delectromagne

tic%2Bwaves%26um%3D1%26hl%3Den%26newwindow%3D1%26client%3Dfirefox-

a%26rls%3Dorg.mozilla:en-US:official