User:Difei Shi/Flame test

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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). 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. The color of the flames also generally depends on temperature and oxygen fed; see flame colors. 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. Wooden splints, Nichrome wires, cotton swabs, and melamine foam are suggested for support. Safety precautions are crucial due to the flammability and toxicity of some substances involved. The test provides qualitative data; therefore, obtaining quantitative data requires subsequent techniques like flame photometry or flame emission spectroscopy.

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. 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. 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). Today, this low-cost method is used in secondary education to teach students to detect metals in samples qualitatively.

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). As they fall back, they release energy in the form of photons, which are responsible for the light.

Electron transition of a $d^{9}$ transition metal. Electron is promoted from $t_{2}g$ set to $e_{g}$ 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. MLCT is when an electron from the $t_{2}g$ set or the $e_{g}$ set (more metal in character) is transferred to the $t_{1}$u set (more ligand in character). LMCT is when an electron from $\pi$ or $\pi ^{*}$ orbital (more ligand in character) is promoted to $\sigma ^{*}$ orbital (more metal in character). The d-d transition involves transferring an electron from $t_{2}g$ set to $e_{g}$ set.

Different atoms or ions require different amounts of energy to promote one electron to one orbital higher than its normal state. Electron excitation occurs at specific energy levels, indicating that the energy is quantized and corresponds to specific wavelengths (or frequencies). 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. 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. 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. 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.

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.

• Barium chloride is toxic. It is important to prevent ingestion of the salt or solution.
• Open flames (Bunsen burner or propane torch) are sources of fire hazards. It is important to ensure the experiment area is free of flammable materials.
• “Flame jetting” 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. It is important to not add excess solvent when flames diminish.

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	Image
Al	Aluminium	Silvery 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 the 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.