World Library  
Flag as Inappropriate
Email this Article

Ammonium chloride

Article Id: WHEBN0000293336
Reproduction Date:

Title: Ammonium chloride  
Author: World Heritage Encyclopedia
Language: English
Subject: Flux (metallurgy), Sodium carbonate, Scandium oxide, Solvay process, Acid
Collection: Alchemical Substances, Ammonium Compounds, Chlorides, Food Additives, Nonmetal Halides
Publisher: World Heritage Encyclopedia

Ammonium chloride

Ammonium chloride
Unit cell of ammonium chloride
Powder of ammonium chloride
IUPAC name
Ammonium chloride
Other names
Sal ammoniac, Salmiac, Nushadir salt, Sal armagnac, Salt armoniack
ATC code B05
ChemSpider  Y
EC number 235-186-4
Jmol-3D images Image
RTECS number BP4550000
Molar mass 53.49 g·mol−1
Appearance White solid, hygroscopic
Odor Odorless
Density 1.5274 g/cm3[1]
Melting point 338 °C (640 °F; 611 K) decomposes, sublimes
Boiling point 520 °C (968 °F; 793 K)
Sublimes at 337.6 °C[2]
ΔsublHo = 176.1 kJ/mol[3]
244 g/L (−15 °C)
294 g/L (0 °C)
283.0 g/L (25 °C)
454.4 g/L (40 °C)
740.8 g/L (100 °C)[4]
30.9 (395 g/L)[5]
Solubility Soluble in liquid ammonia, hydrazine, alcohol, methanol, glycerol
Slightly soluble in acetone
Insoluble in diethyl ether, ethyl acetate[2]
Solubility in methanol 3.2 g/100 g (17 °C)
3.35 g/100 g (19 °C)
3.54 g/100 g (25 °C)[2]
Solubility in ethanol 6 g/L (19 °C)[1]
Solubility in glycerol 97 g/kg[2]
Solubility in sulfur dioxide 0.09 g/kg (0 °C)
0.031 g/kg (25 °C)[2]
Solubility in acetic acid 0.67 g/kg (16.6 °C)[2]
Vapor pressure 133.3 Pa (160.4 °C)[6]
6.5 kPa (250 °C)
33.5 kPa (300 °C)[1]
Acidity (pKa) 9.24
1.642 (20 °C)[2]
84.1 J/mol·K[1]
94.56 J/mol·K[1]
−314.43 kJ/mol[1]
−202.97 kJ/mol[1]
Safety data sheet ICSC 1051
GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[6]
GHS signal word Warning
H302, H319[6]
Harmful Xn Irritant Xi
R-phrases R22, R36
S-phrases (S2), S22
NFPA 704
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
LD50 (Median dose)
1650 mg/kg (rats, oral)
US health exposure limits (NIOSH):
PEL (Permissible)
REL (Recommended)
TWA 10 mg/m3 ST 20 mg/m3 (as fume)[7]
Related compounds
Other anions
Ammonium fluoride
Ammonium bromide
Ammonium iodide
Other cations
Sodium chloride
Potassium chloride
Hydroxylammonium chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
 Y  (: Y/N?)

Ammonium chloride, an salt, highly soluble in water. Solutions of ammonium chloride are mildly acidic. Sal ammoniac is a name of the natural, mineralogical form of ammonium chloride. The mineral is commonly formed on burning coal dumps, due to condensation of coal-derived gases. It is also found around some types of volcanic vents. It is used as a flavouring agent in some types of liquorice. It is the product from the reaction of hydrochloric acid and ammonia.


  • Sources 1
  • Reactions 2
  • Applications 3
    • Metalwork 3.1
    • Medicine 3.2
    • Food 3.3
    • In the laboratory 3.4
    • Flotation 3.5
    • Other applications 3.6
  • References 4
  • External links 5


It is a product of the Solvay process used to produce sodium carbonate.[3]

CO2 + 2 NH3 + 2 NaCl + H2O → 2 NH4Cl + Na2CO3

In addition to being the principal method for the manufacture of ammonium chloride, this method is used to minimize ammonia release in some industrial operations. Ammonium chloride is prepared commercially by combining ammonia (NH3) with either hydrogen chloride (gas) or hydrochloric acid (water solution):[3]

NH3 + HCl → NH4Cl

Ammonium chloride occurs naturally in volcanic regions, forming on volcanic rocks near fume-releasing vents (fumaroles). The crystals deposit directly from the gaseous state, and tend to be short-lived, as they dissolve easily in water.[8]


Ammonium chloride appears to sublime upon heating. However, this process is actually decomposition into ammonia and hydrogen chloride gas.[3]

NH4Cl → NH3 + HCl

Ammonium chloride reacts with a strong base, e.g. sodium hydroxide, to release ammonia gas:

NH4Cl + NaOH → NH3 + NaCl + H2O

Similarly, ammonium chloride also reacts with alkali metal carbonates at elevated temperatures, giving ammonia and alkali metal chloride:

2 NH4Cl + Na2CO3 → 2 NaCl + CO2 + H2O + 2 NH3

A 5% by weight solution of ammonium chloride in water has a pH in the range 4.6 to 6.0.[9]

Some of ammonium chloride's reactions with other chemicals are endothermic, for example its reaction with barium hydroxide, and its dissolving in water.


Ammonium chloride crystal(s)

The main application of ammonium chloride is as a nitrogen source in fertilizers (corresponding to 90% of the world production of ammonium chloride), e.g. chloroammonium phosphate. The main crops are rice and wheat in Asia.[10]

Ammonium chloride was used in pyrotechnics in the 18th century but was superseded by safer and less hygroscopic chemicals. Its purpose was to provide a chlorine donor to enhance the green and blue colours from copper ions in the flame.

It had a secondary use to provide white smoke but its ready double decomposition reaction with potassium chlorate producing the highly unstable ammonium chlorate made its use very suspect.[11][12][13]


Ammonium chloride is used as a flux in preparing metals to be tin coated, galvanized or soldered. It works as a flux by cleaning the surface of workpieces by reacting with the metal oxides at the surface to form a volatile metal chloride. For this purpose, it is sold in blocks at hardware stores for use in cleaning the tip of a soldering iron and can also be included in solder as flux.


Ammonium chloride is used as an expectorant in cough medicine. Its expectorant action is caused by irritative action on the bronchial mucosa. This causes the production of excess respiratory tract fluid which presumably is easier to cough up. Ammonium salts are an irritant to the gastric mucosa and may induce nausea and vomiting.

Ammonium chloride is used as a systemic acidifying agent in treatment of severe metabolic alkalosis, in oral acid loading test to diagnose distal renal tubular acidosis, to maintain the urine at an acid pH in the treatment of some urinary-tract disorders.


In several countries, ammonium chloride, known as sal ammoniac, is used as food additive under the E number E510, commonly as a yeast nutrient in breadmaking. It is a feed supplement for cattle and an ingredient in nutritive media for yeasts and many microorganisms.

Ammonium chloride is used to spice up dark sweets called salty liquorice (very popular in Nordic countries, especially in Finland), in baking to give cookies a very crisp texture, and in the vodka Salmiakki Koskenkorva for flavouring. In India and Pakistan, it is called "Noshader" and is used to improve the crispness of snacks such as samosas and jalebi.

In the laboratory

Ammonium chloride is used to produce low temperatures in cooling baths.[14] Ammonium chloride solutions with ammonia are used as buffer solutions.

In paleontology, ammonium chloride vapor is deposited on fossils, where the substance forms a brilliant white, easily removed and fairly harmless and inert layer of tiny crystals. This covers up any coloration the fossil may have, and if lighted at an angle highly enhances contrast in photographic documentation of three-dimensional specimens.[3] The same technique is applied in archaeology to eliminate reflection on glass and similar specimens for photography.[4]


Giant squid and some other large squid species maintain neutral buoyancy in seawater through an ammonium chloride solution which is found throughout their bodies and is lighter than seawater. This differs from the method of flotation used by most fish, which involves a gas-filled swim bladder. The solution tastes somewhat like salmiakki and makes giant squid unattractive for general human consumption.

Other applications

Ammonium chloride is used in a ~5% aqueous solution to work on oil wells with clay swelling problems. It is also used as electrolyte in zinc–carbon batteries. Other uses include in hair shampoo, in the glue that bonds plywood, and in cleaning products. In hair shampoo, it is used as a thickening agent in ammonium-based surfactant systems, such as ammonium lauryl sulfate. Ammonium chloride is used in the textile and leather industry in dyeing, tanning, textile printing and to luster cotton.


  1. ^ a b c d e f g Pradyot, Patnaik (2003). Handbook of Inorganic Chemicals. The McGraw-Hill Companies, Inc.  
  2. ^ a b c d e f g
  3. ^ a b c d Wiberg, Egon; Wiberg, Nils (2001). Inorganic Chemistry (illustrated ed.). Academic Press. p. 614.  
  4. ^ Seidell, Atherton; Linke, William F. (1919). Solubilities of Inorganic and Organic Compounds (2nd ed.). D. Van Nostrand Company. 
    Results here are multiplied by water's density at temperature of solution for unit conversion.
  5. ^ "Solubility Products of Selected Compounds". Salt Lake Metals. Retrieved 2014-06-11. 
  6. ^ a b c d Sigma-Aldrich Co., Ammonium chloride. Retrieved on 2014-06-11.
  7. ^ a b c "NIOSH Pocket Guide to Chemical Hazards #0029".  
  8. ^ Rowley, Steven P. (2011). General Chemistry I Laboratory Manual (Second ed.).  
  9. ^ Dr. K. G. Bothara (7 October 2008). Inorganic Pharmaceutical Chemistry. Pragati Books Pvt. Ltd. pp. 13–.  
  10. ^ Karl-Heinz Zapp "Ammonium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2012, Wiley-VCH, Weinheim. doi:10.1002/14356007.a02_243
  11. ^ John A. Conkling; Christopher J. Mocella (2010). Chemistry of Pyrotechnics (2nd ed.). CRC Press.  
  12. ^ Tenney L Davis (2012). Chemistry of Powder and Explosives. Angriff Press.  
  13. ^ K. L. Kosanke; B. J. Kosanke; Barry T. Sturman; Robert M. Winokur (2012). Encyclopedic Dictionary of Pyrotechnics (and Related Subjects). Journal of Pyrotechnics.  
  14. ^ "A New Frigorifick Experiment Shewing, How a Considerable Degree of Cold May be Suddenly Produced without the Help of Snow, Ice, Haile, Wind, or Niter, and That at Any Time of the Year" (PDF). Philosophical Transactions 1: 255–261. 18 July 1666. 

External links

  • Calculators: surface tensions, and densities, molarities and molalities of aqueous ammonium chloride
  • CDC - NIOSH Pocket Guide to Chemical Hazards
This article was sourced from Creative Commons Attribution-ShareAlike License; additional terms may apply. World Heritage Encyclopedia content is assembled from numerous content providers, Open Access Publishing, and in compliance with The Fair Access to Science and Technology Research Act (FASTR), Wikimedia Foundation, Inc., Public Library of Science, The Encyclopedia of Life, Open Book Publishers (OBP), PubMed, U.S. National Library of Medicine, National Center for Biotechnology Information, U.S. National Library of Medicine, National Institutes of Health (NIH), U.S. Department of Health & Human Services, and, which sources content from all federal, state, local, tribal, and territorial government publication portals (.gov, .mil, .edu). Funding for and content contributors is made possible from the U.S. Congress, E-Government Act of 2002.
Crowd sourced content that is contributed to World Heritage Encyclopedia is peer reviewed and edited by our editorial staff to ensure quality scholarly research articles.
By using this site, you agree to the Terms of Use and Privacy Policy. World Heritage Encyclopedia™ is a registered trademark of the World Public Library Association, a non-profit organization.

Copyright © World Library Foundation. All rights reserved. eBooks from Hawaii eBook Library are sponsored by the World Library Foundation,
a 501c(4) Member's Support Non-Profit Organization, and is NOT affiliated with any governmental agency or department.