Jump to content

User:Difei Shi/Flame test

From Wikipedia, the free encyclopedia

Article Draft

[edit]

Lead

[edit]

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].

Article body

[edit]

History

[edit]

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

[edit]

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[4].

Electron transition of a transition metal. Electron is promoted from set to 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 set or the set (more metal in character) is transferred to the u set (more ligand in character)[17]. LMCT is when an electron from or orbital (more ligand in character) is promoted to orbital (more metal in character)[17]. The d-d transition involves transferring an electron from set to 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].

Safety hazards

[edit]
  • 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.

Common elements

[edit]
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 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 Flame resulting from Barium Chloride combustion in a Bunsen burner
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 Flame test on copper sulfate
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 An iron (III) flame, generated using the thermite reaction
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[6].

References

[edit]
  1. ^ a b c d "Robert Bunsen and Gustav Kirchhoff". Science History Institute. Retrieved 2023-10-21.
  2. ^ a b c 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. ^ a b c d "This Month in Physics History". www.aps.org. Retrieved 2023-11-02.
  4. ^ a b c d e f g "Flame Tests". Chemistry LibreTexts. 2013-10-03. Retrieved 2023-10-24.
  5. ^ a b c "8: Flame Tests of Metal Cations (Experiment)". Chemistry LibreTexts. 2017-10-31. Retrieved 2023-10-24.
  6. ^ a b c d "Flame Test | Explanation, Definition, Information & Summary". Chemistry Dictionary. 2019-10-14. Retrieved 2023-11-02.
  7. ^ a b c d Wacowich-Sgarbi, Shirley; Langara Chemistry Department (2018). "8.2 Quantization of the Energy of Electrons". Pressbooks BC Campus.
  8. ^ a b Helmenstine, Anne (2022-06-15). "Flame Test Colors and Procedure (Chemistry)". Science Notes and Projects. Retrieved 2023-11-01.
  9. ^ "Flame Test". Chemistry LibreTexts. 2019-05-14. Retrieved 2023-11-19.
  10. ^ "flame tests". www.chemguide.co.uk. Retrieved 2023-11-19.
  11. ^ 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.
  12. ^ 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.
  13. ^ a b "Safety Alert: Do Not Use Methanol-Based Flame Tests on Open Laboratory Desks | NSTA". www.nsta.org. Retrieved 2023-10-24.
  14. ^ a b c d Emerson, Jillian Meri. "New and Improved -- Flame Test Demonstration ("Rainbow Demonstration")". American Chemical Society.
  15. ^ a b c 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.
  16. ^ "Atomic Absorption Spectroscopy (AAS)|PerkinElmer". www.perkinelmer.com. Retrieved 2023-11-19.
  17. ^ a b c d "Electronic Spectroscopy - Interpretation". Chemistry LibreTexts. 2013-10-02. Retrieved 2023-11-20.
  18. ^ "Visible Light - NASA Science". science.nasa.gov. Retrieved 2023-10-24.