Flame test

A flame test is relatively quick test for the presence of some elements in a sample. The technique is archaic and of questionable reliability, but once was a component of qualitative inorganic analysis. The phenomenon is related to pyrotechnics and atomic emission spectroscopy.[1] The color of the flames is understood through the principles of atomic electron transition and photoemission, where varying elements require distinct energy levels (photons) for electron transitions.[2][3]

History

Robert Bunsen invented the now-famous Bunsen burner in 1855, which was useful in flame tests due to its non-luminous flame that did not disrupt the colors emitted by the test materials.[4][1] The Bunsen burner, combined with a prism (filtering the color interference of contaminants), led to the creation of the spectroscope, capable of emitting the spectral emission of various elements.[1] In 1860, the unexpected appearance of sky-blue and dark red was observed in spectral emissions by Robert Bunsen and Gustav Kirchhoff, leading to the discovery of two alkali metals, caesium (sky-blue) and rubidium (dark red).[4][1] Today, this low-cost method is used in secondary education to teach students to detect metals in samples qualitatively.[2]

Process

A flame test involves introducing a sample of the element or compound to a hot, non-luminous flame and observing the color of the flame that results.[4] The compound can be made into a paste with concentrated hydrochloric acid, as metal halides, being volatile, give better results.[5] Different flames can be tried to verify the accuracy of the color. Wooden splints, Nichrome wires, cotton swabs, and melamine foam are suggested for support.[6][7][8] Safety precautions are crucial due to the flammability and toxicity of some substances involved.[9][10][11][6] When using a splint, one must be careful to wave the splint through the flame rather than holding it in the flame for extended periods, to avoid setting the splint itself on fire. The use of a cotton swab or melamine foam (used in “eraser” cleaning sponges) as a support has also been suggested.[7][8][6] Sodium is a common component or contaminant in many samples,[2] and its spectrum tends to dominate many flame tests others.[5] The test flame is often viewed through cobalt blue glass to filter out the yellow of sodium and allow for easier viewing of other metal ions.

The color of the flames also generally depends on temperature and oxygen fed; see flame colors.[5] The procedure uses different solvents and flames to view the test flame through a cobalt blue glass to filter the interfering light of contaminants such as sodium.[12]

Flame tests are subject of a number of limitations. The range of elements positively detectable under standard conditions is small. Some elements emit weakly and others (Na) very strongly. Gold, silver, platinum, palladium, and a number of other elements do not produce a characteristic flame color, although some may produce sparks (as do metallic titanium and iron); salts of beryllium and gold reportedly deposit pure metal on cooling.[12] The test is highly subjective.

Principle

In flame tests, ions are excited thermally. These excited states then relax to the ground state with emission of a photon. The energy of the excited state(s) and associated emitted photon is characteristic of the element. The nature of the excited and ground states depends only on the element. Ordinarily, there are no bonds to be broken, and molecular orbital theory is not applicable. The emission spectrum observed in flame test is also the basis of flame emission spectroscopy, atomic emission spectroscopy, and flame photometry.[4][13]

Common elements

Some common elements and their corresponding colors are:

Symbol Name Color[5] Image
Al Aluminium Silver-white, in very high temperatures such as an electric arc, light blue
As Arsenic Blue
B Boron Bright green
Ba Barium Light apple green
Be Beryllium White
Bi Bismuth Azure blue
C Carbon Bright orange
Ca Calcium Brick/orange red; light green as seen through blue glass.
Cd Cadmium Brick red
Ce Cerium Yellow
Co Cobalt Silvery white
Cr Chromium Silvery white
Cs Caesium Blue-violet
Cu(I) Copper(I) Blue-green
Cu(II) Copper(II) (non-halide) Green
Cu(II) Copper(II) (halide) Blue-green
Ge Germanium Pale blue
Fe(II) Iron(II) Gold, when very hot such as an electric arc, bright blue, or green turning to orange-brown
Fe(III) Iron(III) Orange-brown
H Hydrogen Pale blue
Hf Hafnium White
Hg Mercury Red
In Indium Indigo blue
K Potassium Lilac (pink); invisible through cobalt blue glass (purple)
Li Lithium Carmine red; invisible through green glass
Mg Magnesium Colorless due to Magnesium Oxide layer, but burning Mg metal gives an intense white
Mn(II) Manganese(II) Yellowish green
Mo Molybdenum Yellowish green
Na Sodium Bright yellow; invisible through cobalt blue glass. See also Sodium-vapor lamp
Nb Niobium Green or blue
Ni Nickel Colorless to silver-white
P Phosphorus Pale blue-green
Pb Lead Blue-white
Ra Radium Crimson red
Rb Rubidium Violet red
S Sulfur Blue
Sb Antimony Pale green
Sc Scandium Orange
Se Selenium Azure blue
Sn Tin Blue-white
Sr Strontium Crimson to scarlet red; yellowish through green glass and violet through blue cobalt glass
Ta Tantalum Blue
Te Tellurium Pale green
Ti Titanium Silver-white
Tl Thallium Pure green
V Vanadium Yellowish green
W Tungsten Green
Y Yttrium Carmine, crimson, or scarlet red
Zn Zinc Colorless to blue-green
Zr Zirconium Mild/dull red

See also

References
  1. "This Month in Physics History". www.aps.org. Retrieved 2023-11-02.
  2. Moraes, Edgar P.; da Silva, Nilbert S. A.; de Morais, Camilo de L. M.; Neves, Luiz S. das; Lima, Kassio M. G. de (2014-11-11). "Low-Cost Method for Quantifying Sodium in Coconut Water and Seawater for the Undergraduate Analytical Chemistry Laboratory: Flame Test, a Mobile Phone Camera, and Image Processing". Journal of Chemical Education. 91 (11): 1958–1960. doi:10.1021/ed400797k. ISSN 0021-9584.
  3. Wacowich-Sgarbi, Shirley; Langara Chemistry Department (2018). "8.2 Quantization of the Energy of Electrons". Pressbooks BC Campus.
  4. "Robert Bunsen and Gustav Kirchhoff". Science History Institute. Retrieved 2023-10-21.
  5. Helmenstine, Anne (2022-06-15). "Flame Test Colors and Procedure (Chemistry)". Science Notes and Projects. Retrieved 2023-11-01.
  6. Clark, Jim (August 2018). "Flame Tests". chemguide.co.uk. Archived from the original on November 27, 2020. Retrieved January 10, 2021.
  7. Sanger, Michael J.; Phelps, Amy J.; Catherine Banks (2004-07-01). "Simple Flame Test Techniques Using Cotton Swabs". Journal of Chemical Education. 81 (7): 969. doi:10.1021/ed081p969. ISSN 0021-9584.
  8. Landis, Arthur M.; Davies, Malonne I.; Landis, Linda; Nicholas C. Thomas (2009-05-01). ""Magic Eraser" Flame Tests". Journal of Chemical Education. 86 (5): 577. doi:10.1021/ed086p577. ISSN 0021-9584.
  9. "Safety Alert: Do Not Use Methanol-Based Flame Tests on Open Laboratory Desks | NSTA". www.nsta.org. Retrieved 2023-10-24.
  10. Emerson, Jillian Meri. "New and Improved -- Flame Test Demonstration ("Rainbow Demonstration")". American Chemical Society.
  11. Sigmann, Samuella B. (2018-10-09). "Playing with Fire: Chemical Safety Expertise Required". Journal of Chemical Education. 95 (10): 1736–1746. doi:10.1021/acs.jchemed.8b00152. ISSN 0021-9584.
  12. "Flame Test | Explanation, Definition, Information & Summary". Chemistry Dictionary. 2019-10-14. Retrieved 2023-11-02.
  13. "Atomic Absorption Spectroscopy (AAS)|PerkinElmer". www.perkinelmer.com. Retrieved 2023-11-19.
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