Ammonium dichromate

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Ammonium dichromate
(NH4)2Cr2O7.JPG
Ammonium-dichromate-2D.png
Ammonium-dichromate-xtal-2007-CM-3D-balls.png
Identifiers
CAS number 7789-09-5 YesY
ChemSpider 23002 YesY
UN number 1439
RTECS number HX7650000
Jmol-3D images Image 1
Properties
Molecular formula (NH4)2Cr2O7
Molar mass 252.07 g/mol
Appearance Orange-red crystals
Density 2.115 g/cm3
Melting point 180 °C (356 °F; 453 K) decomposes
Solubility in water 18.2 g/100ml (0 °C)
35.6 g/100ml (20 °C)
40 g/100ml (25 °C)
156 g/100ml (100 °C)
Solubility insoluble in acetone
soluble in alcohol
Hazards
MSDS ICSC 1368
GHS pictograms GHS-pictogram-rondflam.svgGHS-pictogram-acid.svgGHS-pictogram-skull.svgGHS-pictogram-silhouete.svgGHS-pictogram-pollu.svg[1]
GHS hazard statements H272, H301, H312, H314, H317, H330, H334, H340, H350, H360, H372, H410[1]
GHS precautionary statements P201, P220, P260, P273, P280, P284[1]
EU Index 024-003-00-1
EU classification Explosive EOxidizing Agent OVery Toxic T+Corrosive CDangerous for the Environment (Nature) NHarmful Xn
Carc. Cat. 2
Muta. Cat. 2
Repr. Cat. 2
R-phrases R45, R46, R60, R61, R2, R8, R21, R25, R26, R34, R42/43, R48/23, R50/53
S-phrases S53, S45, S60, S61
NFPA 704
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Autoignition temperature 190 °C
Related compounds
Other cations Potassium dichromate
Sodium dichromate
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Ammonium dichromate is the inorganic compound with the formula (NH4)2Cr2O7. In this compound, as in all chromates and dichromates, chromium is in a +6 oxidation state, commonly known as hexavalent chromium. It is a salt consisting of ammonium ions and dichromate ions.

Ammonium dichromate is sometimes known as Vesuvian Fire, because of its use in demonstrations of tabletop "volcanoes".[2] It has been used in pyrotechnics and in the early days of photography.

Properties[edit]

At room temperature and pressure, the compound exists as orange, acidic crystals soluble in water and alcohol. It is formed by the action of chromic acid on ammonium hydroxide with subsequent crystallisation.[3]

The (NH4)2Cr2O7 crystal (C2/c, z=4) contains a single type of ammonium ion, at sites of symmetry C1(2,3). Each NH4+ centre is surrounded irregularly by eight oxygen atoms at N—O distances ranging from ca. 2.83 to ca. 3.17 Å, typical of hydrogen bonds.[4]

Uses[edit]

It has been used in pyrotechnics and in the early days of photography as well as in lithography, as a source of pure nitrogen in the laboratory, and as a catalyst.[5] It is also used as a mordant for dyeing pigments, in the manufacturing of alizarin, chrome alum, leather tanning and oil purification.[3]

Photosensitive films containing PVA, ammonium dichromate, and a phosphor are spin-coated as aqueous slurries in the production of the phosphor raster of television screens and other devices. The ammonium dichromate acts as the photoactive site.[6]

Reactions[edit]

Tabletop volcanoes and thermal decomposition[edit]

The volcano demonstration involves igniting a pile of the salt, which initiates the following exothermic conversion:[7]

(NH
4
)
2
Cr
2
O
7
(s)Cr
2
O
3
(s) + N
2
(g) + 4 H
2
O
(g) (ΔH=−429.1 ± 3 kcal/mol)
Ammonium dichromate decomposition

Like the well-known explosive ammonium nitrate, it is thermodynamically unstable.[8][9] Its decomposition reaction proceeds to completion once initiated, producing voluminous dark green powdered chromium(III) oxide. Not all of the ammonium dichromate decomposes in this reaction. When the green powder is brought into water a yellow/orange solution is obtained from left over ammonium dichromate.

Observations obtained using relatively high magnification microscopy during a kinetic study of the thermal decomposition of ammonium dichromate provided evidence that salt breakdown proceeds with the intervention of an intermediate liquid phase rather than a solid phase. The characteristic darkening of (NH
4
)
2
Cr
2
O
7
crystals as a consequence of the onset of decomposition can be ascribed to the dissociative loss of ammonia accompanied by progressive anion condensation to Cr
3
O2−
10
, Cr
4
O2−
13
, etc., ultimately yielding CrO
3
. The CrO
3
has been identified as a possible molten intermediate participating in (NH
4
)
2
Cr
2
O
7
decomposition.[10]

Oxidation reactions[edit]

Ammonium dichromate is a strong oxidising agent and reacts, often violently, with any reducing agent. The stronger the reducing agent, the more violent the reaction.[8] It has also been used to promote the oxidation of alcohols and thiols. Ammonium dichromate, in the presence of Mg(HSO4)2 and wet SiO2 can act as a very efficient reagent for the oxidative coupling of thiols under solvent free conditions. The reactions produces reasonably good yields under relatively mild conditions.[11] The compound is also used in the oxidation of aliphatic alcohols to their corresponding aldehydes and ketones in ZrCl4/wet SiO2 in solvent free conditions, again with relatively high yields.[12][13]

Safety[edit]

Ammonium dichromate, like all chromium(VI) compounds, is highly toxic and a proven carcinogen.[14] It is also a strong irritant.

Incidents[edit]

In sealed containers, ammonium dichromate is likely to explode if heated.[8] On January 19, 1986, The New York Times reported that two workers had been killed and 14 others injured at Diamond Shamrock Chemicals, Ashtabula, Ohio, when 2,000 lbs of ammonium dichromate exploded as it was being dried in a heater.[15]

References[edit]

  1. ^ a b c Sigma-Aldrich Co., Ammonium dichromate. Retrieved on 2013-07-20.
  2. ^ "Ammonium Dichromate Volcano". Chemistry Comes Alive!. J. Chem. Ed. 
  3. ^ a b Richard J. Lewis Hawley's Condensed Chemical Dictionary. Wiley & Sons, Inc: New York, 2007 ISBN 978-0-471-76865-4
  4. ^ Keresztury, G. and Knop, O. (1982). "Infrared spectra of the ammonium ion in crystals. Part XII. Low-temperature transitions in ammonium dichromate, (NH4)2Cr2O7". Can. J. Chem.: 1972–1976. 
  5. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  6. ^ Havard, J. M.; Shim, S. Y.; and Fréchet, J. M. (1999). "Design of Photoresists with Reduced Environmental Impact. 1. Water-Soluble Resists Based on Photo-Cross-Linking of Poly(vinyl alcohol)". Chem. Mater. 11 (3): 719–725. doi:10.1021/cm980603y. 
  7. ^ Neugebauer, C. A. and Margrave, J. L. (1957). "The Heat Formation of Ammonium Dichromate". J. Phys. Chem. 61 (10): 1429–1430. doi:10.1021/j150556a040. 
  8. ^ a b c Young, A.J. (2005). "CLIP, Chemical Laboratory Information Profile: Ammonium Dichromate". J. Chem. Educ. 82 (11): 1617. doi:10.1021/ed082p1617. 
  9. ^ G. A. P. Dalgaard, A. C. Hazell and R. G. Hazell (1974). "The Crystal Structure of Ammonium Dichromate, (NH4)2Cr2O7". Acta Chemica Scandinavica A28: 541–545. doi:10.3891/acta.chem.scand.28a-0541. 
  10. ^ Galwey, Andrew K.; Pöppl, Làszlò; Rajam, Sundara (1983). "A Melt Mechanism for the Thermal Decomposition of Ammonium Dichromate". J. Chem. Soc., Faraday Trans. 1 79 (9): 2143–2151. doi:10.1039/f19837902143. 
  11. ^ Shirini, F., et al. (2003). "Solvent free oxidation of thiols by (NH4)2Cr2O7 in the presence of Mg(HSO4)2 and wet SiO2". J. Chem. Research (S) 2003: 28–29. doi:10.3184/030823403103172823. 
  12. ^ Shirini, F., et al. (2001). "ZrCl4/wet SiO2 promoted oxidation of alcohols by (NH4)2Cr2O7 in solution and solvent free condition". J. Chem. Research (S) 2001 (11): 467–477. doi:10.3184/030823401103168541. 
  13. ^ F. Shirini, M. A. Zolfigol,FOO† and M. Khaleghi (2003). "Oxidation of Alcohols Using (NH4)2Cr2O7 in the Presence of Silica Chloride/Wet SiO2 in Solution and under Solvent Free Conditions". Bull. Korean Chem. Soc. 24 (7): 1021–1022. doi:10.5012/bkcs.2003.24.7.1021. 
  14. ^ Volkovich, V. A. and Griffiths, T. R. (2000). "Catalytic Oxidation of Ammonia: A Sparkling Experiment". J. Chem. Ed. 77 (2): 177. doi:10.1021/ed077p177. 
  15. ^ Diamond, S. The New York Times, 1986, p. 22.

External links[edit]