A flame test, invented by Robert Bunsen, is a qualitative analysis technique used in chemistry to detect the presence of certain elements, primarily metal ions, based on each element's characteristic flame emission spectrum (which may be affected by the presence of chloride ions).^{[1][2][3][4][5][6]} 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][4][6][7] The color of the flames also generally depends on temperature and oxygen fed; see flame colors.^[8] 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.^[6][9] Wooden splints, Nichrome wires, cotton swabs, and melamine foam are suggested for support.^{[4][10][11][12]} Safety precautions are crucial due to the flammability and toxicity of some substances involved.^[13][14][15] The test provides qualitative data; therefore, obtaining quantitative data requires subsequent techniques like flame photometry or flame emission spectroscopy.^[1][16]

History

The flame test carried out on a copper halide. The characteristic bluish-green color of the flame is due to the copper.

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.^[1][3] 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.^[3] 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).^[1][3] Today, this low-cost method is used in secondary education to teach students to detect metals in samples qualitatively.^[2]

Electron excitation

Principle

In flame tests, energy is emitted by the flame in the form of heat. When an atom or an ion absorbs the energy, electrons will jump from a lower energy level (the highest occupied molecular orbital, or HOMO at an unexcited state) to a higher energy level (the lowest occupied energy orbital, or LUMO at an excited state), and later fall back to their normal energy level (unexcited state).^[4] As they fall back, they release energy in the form of photons, which are responsible for the light. This is known as the Franck-Condon principle ^[4]

Electron transition of a ${\displaystyle d^{9}}$ transition metal. Electron is promoted from ${\displaystyle t_{2}g}$ set to ${\displaystyle e_{g}}$ set.

For the transition metals, the electrons can be promoted via metal-to-ligand charge transfer (MLCT), ligand-to-metal charge transfer (LMCT) or d-d transitions.^[17] MLCT is when an electron from the ${\displaystyle t_{2}g}$ set or the ${\displaystyle e_{g}}$ set (more metal in character) is transferred to the ${\displaystyle t_{1}}$u set (more ligand in character).^[17] LMCT is when an electron from ${\displaystyle \pi }$ or ${\displaystyle \pi ^{*}}$ orbital (more ligand in character) is promoted to ${\displaystyle \sigma ^{*}}$ orbital (more metal in character).^[17] The d-d transition involves transferring an electron from ${\displaystyle t_{2}g}$ set to ${\displaystyle e_{g}}$ set.^[17]

Different atoms or ions require different amounts of energy to promote one electron to one orbital higher than its normal state.^[4] Electron excitation occurs at specific energy levels, indicating that the energy is quantized and corresponds to specific wavelengths (or frequencies).^[7] Therefore, as the electrons of different atoms fall to their unexcited energy levels, they will emit different amounts of energy, corresponding to lights of different wavelengths and frequencies.^[7] This allows for the detection of different atoms.

In general, a larger transition between energy levels requires higher energy and will emit higher energy photons.^[5] If the photons are in the visible region of the light spectrum (wavelength 380–700 nm), a colored light emitted by the material is observed with the naked eye.^[5][18] If the energy of the emitted photons is higher or lower than 380–700 nm, light will be emitted in the infrared, ultraviolet, or other regions of the electromagnetic spectrum that cannot be observed with the naked eye.^[4][7]

Bulk samples emit light too, but their light is not ideal for analysis. Bulk samples emit light with hydrochloric acid to remove traces of previous analytes.^[19]

Process

A flame test showing the presence of lithium.

The compound is usually made into a paste with concentrated hydrochloric acid, as metal halides, being volatile, give better results. Different flames should be tried to avoid wrong data due to “contaminated” flames or occasionally to verify the accuracy of the color. In high-school chemistry courses, wooden splints are sometimes used, mostly because solutions can be dried onto them and they are inexpensive. Nichrome wire is also sometimes used.^[19] 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.^[11][12]

Sodium is a common component or contaminant in many compounds, and its spectrum tends to dominate over others. 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.^[9]

Results

The flame test is relatively quick and simple to perform and can be carried out with the basic equipment found in most chemistry laboratories^[5]. However, the range of elements positively detectable under these conditions is small, as the test relies on the subjective experience of the experimenter rather than any objective measurements^[4]. The test has difficulty detecting small concentrations of some elements, while too strong a result may be produced for certain others, which tends to cause fainter colors to not appear.

Although the flame test only gives qualitative information, not quantitative data about the proportion of elements in the sample, quantitative data can be obtained by the related techniques of flame photometry or flame emission spectroscopy. Flame atomic absorption spectroscopy instruments, made by, e.g., PerkinElmer or Shimadzu, can be operated in emission mode according to the instrument manuals.^[16]

Common elements

Coloured flames of methanol solutions of different compounds, burning on cotton wool. From left to right: lithium chloride, strontium chloride, calcium chloride, sodium chloride, barium chloride, trimethyl borate, copper chloride, cesium chloride and potassium chloride.

Some common elements and their corresponding colors are:

Symbol	Name	Color[8]	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
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

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.^[6]

Safety hazards

• Methanol is a highly flammable liquid. A flame test or “rainbow demonstration” that uses methanol should be performed under a fume hood to prevent flash fires or deflagrations.^[13]

• Barium chloride is toxic.^[14] It is important to prevent ingestion of the salt or solution.^[14]
• Open flames (Bunsen burner or propane torch) are sources of fire hazards.^[14] It is important to ensure the experiment area is free of flammable materials.
• “Flame jetting”^[15] occurs when an excess flammable solvent is added to a burning or previously ignited setup, causing vapors to flash back into the solvent container, causing a torch.^[15] It is important to not add excess solvent when flames diminish.

See also

• Bead test
• Spark testing
• Colored fire
• Emission spectrum
• Photoelectric flame photometer
• Qualitative inorganic analysis
• Inductively coupled plasma atomic emission spectroscopy
• Robert Bunsen
• Electron excitation

References

External links

Wikimedia Commons has media related to Flame test.

• Flame Test - Coloring Fire - Pictures of Several Flame Tests, Includes Instructions
• WebMineral.com - Flame Coloration by Element