Ammonium bromide
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| ball-and-stick model of an ammonium cation (left) and a bromide anion (right) | |
| Names | |
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| IUPAC name
Ammonium bromide
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| Identifiers | |
12124-97-9  |
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| ChEBI | CHEBI:85364  |
| ChemSpider | 23804 |
| Jmol 3D model | Interactive image |
| RTECS number | BO9155000liugoiugiu |
| UNII | R0JB3224WS |
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| Properties | |
| NH4Br | |
| Molar mass | 97.94 g/mol |
| Appearance | white powder, hygroscopic |
| Density | 2.429 g/cm3 |
| Melting point | 235 °C (455 °F; 508 K) |
| Boiling point | 452 °C (846 °F; 725 K) |
| 60.6 g/100 mL (0 °C) 78.3 g/100 mL (25 °C) 145 g/100 mL (100 °C) |
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Refractive index (nD)
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1.712 |
| Structure | |
| Isometric | |
| Vapor pressure | {{{value}}} |
| Related compounds | |
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Other anions
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Ammonium fluoride Ammonium chloride Ammonium iodide |
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Other cations
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Sodium bromide Potassium bromide |
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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 verify (what is ?) |
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| Infobox references | |
Ammonium bromide, NH4Br, is the ammonium salt of hydrobromic acid. The chemical crystallizes in colorless prisms, possessing a saline taste; it sublimes on heating and is easily soluble in water. On exposure to air it gradually assumes a yellow color because of the oxidation of traces of bromide (Br−) to bromine (Br2).
Contents
Preparation
Ammonium bromide can be prepared by the direct action of hydrogen bromide on ammonia.
- NH3 + HBr → NH4Br
It can also be prepared by the reaction of ammonia with iron(II) bromide or iron(III) bromide, which may be obtained by passing aqueous bromine solution over iron filings.
- 2 NH3 + FeBr2 + 2 H2O → 2 NH4Br + Fe(OH)2
Reactions
Ammonium bromide is a weak acid with a pKa of ~5 in water. It is an acid salt because the ammonium ion hydrolyzes slightly in water.
Ammonium Bromide is strong electrolyte when put in water.:
- NH4Br(s) → NH4+(aq) + Br−(aq)
Ammonium bromide decomposes to ammonia and hydrogen bromide when heated at elevated temperatures:
- NH4Br → NH3 + HBr
Uses
Ammonium bromide is used for photography in films, plates and papers; in fireproofing of wood; in lithography and process engraving; in corrosion inhibitors; and in pharmaceutical preparations.[1]
References
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<references />, or <references group="..." />- ↑ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
