Atomic Theory: Spectroscopy and Flame Tests

Introduction : When sunlight strikes your skin, you feel its heat. This is a sign that you are absorbing some of the sun's energy. Light is only one form of energy known as electromagnetic radiation. The electromagnetic spectrum below shows all of the other forms of radiation: visible light, X-rays, radio waves, infrared (IR), ultraviolet (UV) rays and microwaves.

The Electro-magnetic Spectrum:

Light & Color: Light that can be seen with the naked eye falls within a small range of wavelengths (~400-700 nm) called the “visible” region. Note that these different forms of radiation possess different energies and present different potential health hazards. For example, patients often wear lead shields when getting dental X- rays, and sunscreens are used to block UVA or UVB rays, forms of ultraviolet radiation that can not only tan skin but can also cause severe sunburns. However, people usually do not worry about exposure to the radio waves being broadcast throughout a city, so they can all listen to their favorite station.

These different forms of radiation have electric and magnetic components that travel in the form of a wave. Imagine throwing a pebble into a still pond and watching the circular ripples moving outward. Like those ripples, each energy wave has a series of high points known as crests and a series of low points known as troughs. The figure below shows two different waves : λ | A | crest

trough λ | B |

GCC CHM 151LL: Atomic Theory: Spectroscopy and Flame Tests © GCC, 2008 page 1 of 3 Wavelength and Frequency: The wavelength, symbolized by the Greek letter lambda (λ), is the distance between two wave crests, which is equal to the distance between two troughs. Notice that the wavelength for the top wave, indicated by λA, is greater than the wavelength for the bottom wave, λB. Since energy waves move, we can count the number of crests or peaks that pass a given point in one second. This is called the wave’s frequency, which is symbolized by the Greek letter nu (v) and measured in units of cycles per second called Hertz (Hz = 1/s = s-1). For two waves moving at the same speed, the longer the wavelength, the fewer crests or troughs that pass per second. Thus, the frequency of waves is inversely proportional to their wavelengths. These two properties can be related to one another using a proportionality constant equal to the speed of the wave. Light and other forms of energy move at the fastest speed possible in a vacuum, defined as the speed of light. The speed of light is given the symbol “c” and is exactly equal to 2.99792458×108 m/s. For convenience sake, the speed of light is often rounded to 3.00×108 m/s. Since this value is rounded, we don’t use it to limit significant figures in calculations. Frequency and wavelength are related according to the following equation:

c = λ v Equation 1

Because the heat emitted by the sun and other energy sources is constant, most scientists believed that energy existed as continuous waves. This belief continued for centuries until a German physicist named Max Planck proposed a controversial new theory: Energy was not only a wave but also a particle. Based on his experiments on Blackbody radiation, Planck theorized that energy is emitted in small bundles called “quanta”—or “quantum” for a single bundle, which led to the theory’s name: quantum mechanics. Each quantum of energy has a specific frequency associated with it, and the frequency is directly proportional to its energy:

E = hv h=6.626 x 10-34 J⋅s Equation 2 where h is Planck's constant, which is equal to 6.626 x 10-34 J⋅s. Albert Einstein later applied Planck’s theory to light, so a “particle” of light is now called a photon. Given its wavelength, the energy of a photon or other kind of energy can be determined by combining Equations 1 and 2. In this lab, you will calculate the wavelength, frequency, and/or energy of various forms of electromagnetic radiation.

Spec-20’s: I. USING LIGHT AND COLOR TO ANALYZE SAMPLES A spectrophotometer (often abbreviated as “Spec-20”) is an instrument that measures the intensity of a light beam passing through a solution. Most Spec-20’s operate in the visible and IR regions; for example, the Genesys Spec-20’s used in our labs use wavelengths ranging from 325 nm to 1100 nm. Spec-20’s generally have wavelength control buttons, a display, and a “zero” button. Inside, there is a white light source, a prism to separate the white light into the spectrum of colors (each with a different wavelength), a sample compartment, and a detector.

GCC CHM 151LL: Atomic Theory: Spectroscopy and Flame Tests © GCC, 2008 page 2 of 3 Spec-20’s are generally used to analyze colored solutions. Light of a given wavelength (selected by the control buttons) passes through the sample, hits the detector, and the detector measures the solution’s absorbance (A), the amount of light absorbed by the solution’s molecules and/or ions. In general, darker solutions are more concentrated (i.e., containing more molecules or ions), and thus have a higher absorbance.

In this experiment, you will use a Spec-20 to determine the color of light at various wavelengths. Note that as you change the wavelength you are actually turning the prism so different colors shine through the slit onto the sample.

II. ATOMIC EMISSION SPECTRA AND FLAME TESTS

The sun is 93 million miles away, and other stars are many light years away. (Note that one light year = six trillion miles or 6 x 1012 miles). In spite of these great distances, the elements in stars can be determined by analyzing the light they give off since the atoms of every element selectively absorb and emit light of specific wavelengths, and thus, specific energies. These characteristic wavelengths account for the different colors substances emit when heated.

The unique colors emitted by each element provided experimental evidence for Danish physicist Neils Bohr to propose that electrons in an atom occupy orbitals, and each orbital has a specific energy. Because negatively charged electrons are attracted to the positively charged nucleus, they prefer to be in the lowest energy orbitals close to the nucleus; thus, an atom is in its “ground” (or lowest energy) state when its electrons occupy the lowest energy orbitals. However, when atoms are heated, they absorb the specific energies needed for their electrons to jump up to higher energy orbitals. Now, the atom is in an “excited” state, which is unstable. Since the electrons prefer to be closer to the nucleus, they quickly return to lower energy orbitals, and the excess energy is emitted. When the energy emitted falls within the visible range, specific colors are observed.

Since every element has a different number of protons and electrons, then the energy gap between its orbitals varies. As a result, different elements release light at different wavelengths, and each element emits a characteristic emission spectrum, often called an “atomic fingerprint”. When an element is heated, it may emit a characteristic color, usually corresponding to one of the colors in its emission spectrum. This accounts for the different colors observed with fireworks. In this experiment, you will observe the characteristic colors given off by various elements using flame tests and then use your observations to identify an unknown.

III. ABSORPTION AND EMISSION SPECTRA

The third part of this lab involves an interactive online tutorial to help explain the process electrons go through when emission and absorption spectra are obtained from pure substances. Go to the Website below (a link is provided on the webpage for this week’s experiment):

http://www.wwnorton.com/college/chemistry/gilbert/overview/ch3.htm

Click on “view tutorial” for the second option, Section 3.3 Light Emission and Absorption (p. 106-109).