Difference between revisions of "Hydroxylamine"
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'''Hydroxylamine''' is a white crystalline solid, widely used as a reducing agent. It has the chemical formula '''NH<sub>2</sub>OH'''. | '''Hydroxylamine''' is a white crystalline solid, widely used as a reducing agent. It has the chemical formula '''NH<sub>2</sub>OH'''. | ||
| − | As pure hydroxylamine is unstable and explosive, its salts such as hydroxylammonium chloride are more often used. It can also be encountered as aqueous solution. | + | As pure hydroxylamine is unstable and explosive, its salts such as [[hydroxylammonium chloride]] are more often used. It can also be encountered as aqueous solution. |
==Properties== | ==Properties== | ||
| Line 166: | Line 166: | ||
==Projects== | ==Projects== | ||
*Make oximes | *Make oximes | ||
| − | *Make hydroxylammonium nitrate | + | *Make hydroxylammonium salts ([[hydroxylammonium chloride|chloride]], [[hydroxylammonium nitrate|nitrate]], [[hydroxylammonium sulfate|sulfate]], [[hydroxylammonium perchlorate|perchlorate]]) |
*Purify ketones and aldehydes | *Purify ketones and aldehydes | ||
*Make [[nitrous oxide]] | *Make [[nitrous oxide]] | ||
Revision as of 18:04, 8 January 2020
| Names | |
|---|---|
| IUPAC name
Hydroxylamine
| |
| Other names
Azinous acid
Aminol Azanol Hydroxyamine Hydroxyazane Hydroxylazane Nitrinous acid | |
| Properties | |
| NH2OH | |
| Molar mass | 33.03 g/mol |
| Appearance | White hygroscopic crystals |
| Odor | Ammonia-like |
| Density | 1.227 g /cm3 (at 20 °C) |
| Melting point | 33 °C (91 °F; 306 K) |
| Boiling point | 70 °C (158 °F; 343 K) (at 60 mm Hg; decomposes) |
| Soluble | |
| Solubility | Soluble in liq. ammonia, ethanol, methanol Poorly soluble in diethyl ether, carbon disulfide, chloroform, propanol, Insoluble in acetone, benzene, petroleum ether, hydrogen sulfide |
| Vapor pressure | 53 mm Hg at 32 °C |
| Acidity (pKa) | 5.95 |
| Thermochemistry | |
| Std molar
entropy (S |
236.18 J·K−1·mol−1 |
| Std enthalpy of
formation (ΔfH |
−39.9 kJ·mol−1 |
| Hazards | |
| Safety data sheet | None |
| Flash point | Explodes at 129 °C |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
408 mg/kg (oral, mouse) 9–70 mg/kg (intraperitoneal mouse, rat); 29 mg/kg (subcutaneous, rat) |
| Related compounds | |
| Related compounds
|
Ammonia |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Hydroxylamine is a white crystalline solid, widely used as a reducing agent. It has the chemical formula NH2OH.
As pure hydroxylamine is unstable and explosive, its salts such as hydroxylammonium chloride are more often used. It can also be encountered as aqueous solution.
Contents
Properties
Chemical
Hydroxylamine reacts with acids to give hydroxylammonium salts.
Reaction with copper(II) oxide gives nitrous oxide:
- 2 NH2OH + 4 CuO → 2 Cu2O + N2O + 3 H2O
Physical
Hydroxylamine is a white crystalline solid, hygroscopic and unstable when pure. It is very soluble in water and alcohol.
Explosive
Freebase hydroxylamine can explode if heated at high temperatures.
Availability
Pure hydroxylamine is not sold by any entity. However, aqueous solutions are sold by many chemical entities, though they're not readily accessible to the amateur chemist.
Hydroxylammonium salts are also available.
Preparation
Hydroxylamine can be prepared through various means:
The Raschig process is the most common route. It involves the reduction of ammonium nitrite with bisulfite (HSO3−) and sulfur dioxide at 0 °C in water to hydroxylamido-N,N-disulfonate anion, which hydrolyzes to give hydroxylamine sulfate.
- NH4NO2 + 2 SO2 + NH3 + H2O → 2 NH4+ + N(OH)(OSO2)22-
- N(OH)(OSO2)22− + H2O → NH(OH)(OSO2)- + HSO4-
- 2 NH(OH)(OSO2)− + 2 H2O → (NH3OH)2SO4 + SO42-
To obtain pure hydroxylamine, anhydrous ammonia is added:
- (NH3OH)2SO4 + 2 NH3 → 2 NH2OH +(NH4)2SO4
Since free hydroxylamine quickly breaks down, the final product is either dissolved in water or reacted with a strong acid, like hydrochloric acid.
Other nitrites, such as nitrous acid or sodium nitrite can also be reduced with bisulfite to hydroxylamine.
Electrolytic reduction of nitric acid with sulfuric acid at 15-20 °C for 40 minutes at 24 A gives hydroxylamine:
- HNO3 + 3 H2 → NH2OH + 2 H2O
Other acids, such as hydrochloric and phosphoric acids can also be used. The yield of this route is between 50-80%.[1]
Electrolysis of sodium nitrate in the presence of hydrogen chloride will also give hydroxylamine. A lead cathode and a coal anode are used as electrodes. The process takes 3 hours.[2]
Reaction of nitrogen dioxide or nitrous acid with tin(II) chloride will give hydroxylamine. The yield is 90%.[3]
Reduction of nitric oxide with tin(II) chloride in hydrochloric acid at 0 °C will also yield hydroxylamine.[4]
Hydroxylamine can also be made by heating nitromethane with concentrated hydrochloric acid. This gives hydroxylamine hydrochloride:
- CH3NO2 + HCl → NH2OH·HCl + 2 H2O + CO
Projects
- Make oximes
- Make hydroxylammonium salts (chloride, nitrate, sulfate, perchlorate)
- Purify ketones and aldehydes
- Make nitrous oxide
- Caprolactam synthesis
- Make Nylon 6
- Make formaldoxime
Handling
Safety
Hydroxylamine is an irritant to the respiratory tract, skin, eyes, and other mucous membranes. It is harmful if swallowed and may be absorbed through the skin. Hydroxylamine is considered a possible mutagen.
Hydroxylamine explodes upon heating, though the exact mechanism is not well understood. Ferrous and ferric salts accelerate its decomposition in aqueous solution.
Storage
Hydroxylamine is safe to store only as aqueous solution or in salt form. Do not store it in pure free base form.
Disposal
Hydroxylamine can be neutralized by adding a ketone and then gently heating the resulting oxime, which reforms the ketone and releases nitrogen gas and water. It can also be neutralized by diluting it with water and carefully adding iron salts or other compounds that accelerate its decomposition.
Reduction of hydroxylamine with Zn/HCl yields ammonia.
References
- ↑ Schoch, E. P.; Pritchett, R. H.; Journal of the American Chemical Society; vol. 38; (1916); p. 2042
- ↑ Tafel, J.; Z. Anorg. Chem.; vol. 31; (1902); p. 321 - 324
- ↑ Raschig, F.; Z. Anorg. Chem.; vol. 155; (1926); p. 225
- ↑ https://pubs.acs.org/doi/10.1021/j150323a006
