By Ron B. Davis Jr., Georgetown University
Electron clouds are the crucial factor in determining how one atom will interact with others. But there is more to the influence of electrons on the properties of atoms. One of these properties has to do with how atoms and ions interact with light. Exciting the elements with flame seems to cause each element to emit distinctive wavelengths of light.
The Blow-pipe Technique
In the late 1700s, it was realized that purified samples of elements and compounds, when subjected to heating by an intense, clean-burning flame, could emit distinctive colors of light. When this is done, sodium compounds produce yellow, lithium produces red, and potassium produces a pale lavender.
This observation led to the development of the so-called blow-pipe analysis technique, in which a sample is burned in a clean-burning flame. The color of the resulting flame is used to help identify the elements a sample contains.
Throughout the 1800s, this technique was also used to prospect for, and discover, many new elements, since each element turned out to have a distinctive emission spectrum. Certain elements, like thallium, cesium, and rubidium even get the inspiration for their names from the colors that they produce in such a test: Rubidium creates ruby red; Cesium was named for a shade of blue in Latin; and Thallium’s Greek name invokes the color of a green leaf.
Spectroscope and the Line Emission Spectra
In 1859, the use of color for identification got another huge shot in the arm with the development of a new tool—the spectroscope. A spectroscope is really a rather simple device, consisting of a prism or diffraction grating that separates light into its constituent wavelengths.
Of course, Isaac Newton figured this out using sunlight as early as 1672 in his early experiments with prisms. But it took almost two more centuries before Gustav Kirchhoff thought to apply the prism’s light-separating ability to an element’s distinctive light emissions from the flame of a blowpipe. But once he did, far more precise information about the elements became possible.
For example, lithium, rubidium, and strontium emissions may all appear red to the unaided eye. But each actually contains many weaker emission lines of other colors. These so-called ‘line emission spectra’ create something of a fingerprint for each element. This allows known elements to be identified more clearly. And it allows undiscovered elements to be detected, with much greater precision.
What Happens When an Element Is Burned?
Placing an element into a flame promotes some of its outermost electrons into an even higher-energy level or orbital. Then, specific wavelengths of light are produced by excited electrons as they transition back into their most stable position, releasing very specific amounts of energy in the process.
However, the amounts of energy released are very specific—with many wavelengths excluded. Being so specific is only possible if there are defined orbits available to electrons. Line-emission spectra prove that there has to be a significant structure within the electron cloud.
Scientists had been staring spectroscopic evidence in the face for nearly sixty years before the connection to the atomic structure was realized! To be fair, that’s because the electron itself wasn’t discovered until 1898, and the nucleus took another decade to figure out.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
The Relationship between Electron Cloud and Light
The electron cloud also has a special relationship with light, being responsible for the universal tendency of elements to emit a specific—apparently unique—‘fingerprint’ of light wavelengths when stimulated by electricity or heat. The realization that these emissions consist of multiple, discrete lines of color is what first tipped off scientists that the electron cloud was highly structured.
The highly structured orbitals, subshells, and shells are usually—though with a few exceptions—filled according to the Aufbau order as we move through the table. This filling of only one subshell at a time influences the shape of the periodic table. The s-, p-, d-, and f-blocks of the periodic table are the direct result of the clear sequence of subshell filling that helps define the complex structure of each atom’s electron cloud.
And this sorting of the elements by electron structure plays a huge role in allowing the table to do such an amazing job of sorting so many elements into groups and clusters with similar properties.
Common Questions about Wavelengths of Light
When purified samples of elements and compounds are subjected to heating by an intense, clean-burning flame, they emit distinctive colors of light. This observation led to the development of the blow-pipe analysis technique, in which a sample is burned in a clean-burning flame. The color of the resulting flame is used to help identify the elements a sample contains.
Placing an element into a flame promotes some of its outermost electrons into an even higher-energy level or orbital. Then, specific wavelengths of light are produced by excited electrons as they transition back into their most stable position, releasing very specific amounts of energy in the process.
The electron cloud has a special relationship with light, being responsible for the universal tendency of elements to emit a specific—apparently unique—‘fingerprint’ of light wavelengths when stimulated by electricity or heat. The realization that these emissions consist of multiple, discrete lines of color is what first tipped off scientists that the electron cloud was highly structured.
