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Spirit of Hartshorn
Colorless gas with strong pungent odor
0.86 kg/m3 (1.013 bar at boiling point)
0.73 kg/m3 (1.013 bar at 15 °C)
681.9 kg/m3 at -33.3°C (liquid)
820 kg/m3 at -80°C (crystal solid)
817 kg/m3 at -80°C (transparent solid)
-77.73 °C (195.42 K)
-33.34 °C (239.81 K)
89.9 g/100 mL at 0 °C
4.75 (reaction with H2O)
Refractive index (nD)
Hazardous gas, caustic, corrosive
Supplementary data page
n, εr, etc.
Solid, liquid, gas
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)
Ammonia is a compound with the
It is normally encountered as a gas with a characteristic pungent
contributes significantly to the nutritional needs of terrestrial organisms by
serving as a precursor to foodstuffs and fertilizers. Ammonia, either directly
or indirectly, is also a building block for the synthesis of many
pharmaceuticals. Although in wide use, ammonia
is both caustic and hazardous. In 2006, worldwide
production was estimated at 146.5 million tonnes.
Ammonia, as used commercially, is often called anhydrous ammonia. This
term emphasizes the absence of water in the material. Because NH3
boils at -33.34 °C, the liquid must be stored under high pressure or at low
temperature. Its heat of vaporization is, however,
sufficiently great that NH3 can be readily handled in ordinary
beakers in a
fume hood. "Household ammonia" or "ammonium hydroxide" is a solution of NH3
in water. The strength of such solutions is measured in units of
with 26 degrees baume (about 30 weight percent ammonia at 15.5 °C) being the
typical high concentration commercial product.
Household ammonia ranges in concentration from 5 to 10 weight percent ammonia.
Structure and basic chemical properties
The ammonia molecule has a trigonal pyramidal shape, as
predicted by VSEPR theory. The
atom in the molecule has a lone electron pair, and ammonia acts as a
base, a proton acceptor. This shape gives the
molecule a dipole
moment and makes it
polar so that ammonia readily dissolves in
water. The degree to which ammonia forms the
ammonium ion increases upon lowering the
pH of the
at "physiological" pH (~7), about 67% of the ammonia molecules are
protonated. Temperature and salinity also affect the proportion of NH4+.
NH4+ has the shape of a regular
The main use of ammonia is for production of
(83% in 2003). Another major application is its conversion to
explosives, because nitric acid is made via oxidation
of ammonia. The entire nitrogen content of all manufactured
organic compounds is derived from ammonia.
Ammonia is found in small quantities in the atmosphere, being produced from the
putrefaction of nitrogenous animal and vegetable matter. Ammonia and
ammonium salts are also found in small quantities in rainwater, whereas
ammonium chloride (sal-ammoniac), and
ammonium sulfate are found in volcanic
districts; crystals of
ammonium bicarbonate have been found in
NH3 to neutralize excess acid.
Ammonium salts also are found distributed through all fertile soil and in
seawater. Substances containing ammonia, or those that are similar to it, are
The Romans called the ammonium chloride deposits they collected from
near the Temple of Jupiter Amun (Greek
Ἄμμων Ammon) in ancient Libya 'sal ammoniacus' (salt of Amun)
because of proximity to the nearby temple.
Salts of ammonia have been known from very early times; thus the term
appears in the writings of
Pliny, although it is not known whether the term is identical with the more
In the form of sal-ammoniac, ammonia was known to the Arabic alchemists as early as the
8th century, first mentioned by Geber (Jabir ibn Hayyan),
and to the European alchemists since the 13th century, being mentioned by
It was also used by dyers
Middle Ages in the form of fermented
to alter the colour of vegetable dyes. In the 15th century,
Basilius Valentinus showed that ammonia
could be obtained by the action of alkalis on sal-ammoniac. At a later period,
when sal-ammoniac was obtained by distilling the hoofs and horns of oxen and
neutralizing the resulting carbonate with
hydrochloric acid, the name "spirit of
hartshorn" was applied to ammonia.
Gaseous ammonia was first isolated by
Joseph Priestley in 1774 and was termed by him
alkaline air; however it was acquired by the alchemist
Eleven years later in 1785,
Claude Louis Berthollet ascertained its
Haber process to produce ammonia from the nitrogen in the air was developed
Haber and Carl Bosch in 1909 and patented in 1910. It was first
used on an industrial scale by the Germans during
following the allied blockade that cut off the supply of nitrates from
ammonia was used to produce explosives to sustain their war effort.
Prior to the advent of cheap natural gas, hydrogen as a precursor to ammonia
production was produced via the electrolysis of water. The
Vemork 60 MW
hydroelectric plant in Norway constructed in 1911 was used purely for this
purpose and up until the second world war provided the majority of Europe's
Synthesis and production
Because of its many uses, ammonia is one of the most highly produced inorganic
chemicals. Dozens of chemical plants worldwide produce ammonia. The
worldwide ammonia production in 2004 was 109 million
The People's Republic of China produced
28.4% of the worldwide production followed by
India with 8.6%,
8.4%, and the United States with 8.2%. About 80% or more of the
ammonia produced is used for fertilizing agricultural crops.
Before the start of World War I, most ammonia was obtained by the
of nitrogenous vegetable and animal waste products, including
dung, where it
distilled by the reduction of
nitrites with hydrogen; in addition, it was produced by the
distillation of coal,
and also by the decomposition of ammonium salts by alkaline
hydroxides such as
quicklime, the salt most generally used being the chloride (sal-ammoniac)
2 NH4Cl + 2 CaO → CaCl2 + Ca(OH)2 + 2 NH3
(Two molecules of ammonium chloride plus two calcium oxide yields calcium
chloride and calcium hydroxide and two molecules of ammonia)
Today, the typical modern ammonia-producing plant first converts
methane) or liquified petroleum gas (such gases are
hydrogen. The processes used in producing the hydrogen begins with removal
compounds from the natural gas (because sulfur deactivates the
catalysts used in subsequent steps). Catalytic
hydrogenation converts organosulfur compounds into gaseous
H2 + RSH → RH + H2S(g)
- The hydrogen sulfide is then removed by passing the gas through beds of
where it is absorbed and converted to solid
H2S + ZnO → ZnS + H2O
CH4 + H2O → CO + 3 H2
CO + H2O → CO2 + H2
- The final step in producing the hydrogen is to use catalytic methanation to
remove any small residual amounts of carbon monoxide or carbon dioxide from the
CO + 3 H2 → CH4 + H2O
CO2 + 4 H2 → CH4 + 2 H2O
- To produce the desired end-product ammonia, the hydrogen is then catalytically
reacted with nitrogen (derived from process air) to form anhydrous liquid
ammonia. This step is known as the ammonia synthesis loop (also referred to as
the Haber-Bosch process):
3 H2 + N2 → 2 NH3
required for ammonia synthesis could in principle be obtained from other
sources, but these alternatives - apart from the electrolysis of water into
oxygen + hydrogen - are presently impractical. At one time, most of Europe's
ammonia was produced from the Hydro plant at
Vemork, via the
electrolysis route. Various renewable energy electricity sources are also
In certain organisms, ammonia is produced from atmospheric N2 by
nitrogenases. The overall process is called
nitrogen fixation. Although it is unlikely
that biomimetic methods will be developed that are competitive with the
process, intense effort has been directed toward understanding the mechanism
of biological nitrogen fixation. The scientific interest in this problem is
motivated by the unusual structure of the active site of the enzyme, which
consists of an Fe7MoS9 ensemble.
Ammonia is also a metabolic product of
deamination. Ammonia excretion is common in aquatic animals. In humans, it
is quickly converted to
urea, which is much less toxic. This urea is a major component of the dry
weight of urine.
Most reptiles, including birds, as well as insects and snails solely excrete
uric acid as nitrogenous waste.
Ammonia is a colorless
gas with a characteristic pungent smell. It is
lighter than air, its density being 0.589 times
that of air. It is easily liquefied due to the strong
hydrogen bonding between molecules; the
liquid boils at
-33.3 °C, and solidifies at -77.7 °C to white crystals.
possesses strong ionizing
powers reflecting its high ε of 22. Liquid ammonia has a very high
standard enthalpy change of
vaporization (23.35 kJ/mol, cf.
water 40.65 kJ/mol,
phosphine 14.6 kJ/mol) and can therefore be used in laboratories in
non-insulated vessels without additional refrigeration.
It is miscible with water. Ammonia in an aqueous solution can
be expelled by boiling. The aqueous solution of ammonia is
basic. The maximum concentration of ammonia in
water (a saturated solution) has a
0.880 g /cm³ and is often known as '.880 Ammonia'.
Ammonia does not burn readily or sustain
except under narrow fuel-to-air mixtures of 15-25% air. When mixed with
burns with a pale yellowish-green flame. At high temperature and in the presence
of a suitable catalyst, ammonia is decomposed into its constituent elements.
Ignition occurs when chlorine is passed into ammonia, forming nitrogen and
hydrogen chloride; if ammonia is present in
excess, then the highly explosive
nitrogen trichloride (NCl3) is
The ammonia molecule readily undergoes
nitrogen inversion at room temperature; a
useful analogy is an umbrella turning itself inside out in a strong wind.
The energy barrier to this inversion is 24.7 kJ/mol, and the
resonance frequency is 23.79
corresponding to microwave radiation of a
of 1.260 cm. The absorption at this frequency was the first
microwave spectrum to be observed.
One of the most characteristic properties of ammonia is its basicity. It
combines with acids
to form salts; thus with
hydrochloric acid it forms
ammonium chloride (sal-ammoniac); with
acid, ammonium nitrate, etc. However perfectly dry
ammonia will not combine with perfectly dry
hydrogen chloride: moisture is necessary to
bring about the reaction.
NH3 + HCl →
The salts produced by the action of ammonia on acids are known as the
ammonium salts and all contain the
Anhydrous ammonia is often used for the
production of methamphetamine.
Although ammonia is well-known as a base, it can also act as an extremely weak
acid. It is a
protic substance, and is capable of formation of
contain the NH2− ion), for example when solid
lithium nitride is added to liquid ammonia, forming a
lithium amide solution:
Li3N(s)+ 2 NH3 (l) → 3 Li+(am)
+ 3 NH2−(am)
Brønsted-Lowry acid-base reaction, ammonia serves as an acid.
The combustion of ammonia to nitrogen and
4NH3 + 3O2 → 2N2 + 6H2O (g)
standard enthalpy change of
combustion, ΔHºc, expressed per
mole of ammonia and
with condensation of the water formed, is –382.81 kJ/mol. Dinitrogen is the
thermodynamic product of combustion: all
nitrogen oxides are unstable with respect to nitrogen and
is the principle behind the
catalytic converter. However, nitrogen
oxides can be formed as kinetic products in the presence of appropriate
catalysts, a reaction of great industrial importance in
the production of nitric acid.
4NH3 + 5O2 → 4NO + 6H2O
The combustion of ammonia in air is very difficult in the absence of a catalyst
platinum gauze), as the temperature of the flame is usually lower than the
ignition temperature of the ammonia-air mixture. The flammable range of ammonia
in air is 16–25%.
Formation of other compounds
In organic chemistry, ammonia can act as a
Amines can be
formed by the reaction of ammonia with alkyl
halides, although the resulting –NH2 group is also nucleophilic
and secondary and tertiary amines are often formed as by-products. An excess of
ammonia helps minimise multiple substitution, and neutralises the
hydrogen halide formed. Methylamine is prepared commercially by the reaction
of ammonia with chloromethane, and the reaction of ammonia with
2-bromopropanoic acid has been used to prepare racemic
70% yield. Ethanolamine is prepared by a ring-opening reaction
with ethylene oxide: the reaction is sometimes allowed
to go further to produce diethanolamine and
Amides can be
prepared by the reaction of ammonia with a number of
carboxylic acid derivatives.
chlorides are the most reactive, but the ammonia must be present in at least
a twofold excess to neutralise the
hydrogen chloride formed.
also react with ammonia to form amides. Ammonium salts of carboxylic acids can
dehydrated to amides so long as there are no thermally sensitive groups
present: temperatures of 150–200 °C are required.
in ammonia is capable of replacement by
burns in the gas with the formation of
magnesium nitride Mg3N2,
and when the gas is passed over heated
sodamide, NaNH2, and potassamide, KNH2, are formed. Where
necessary in substitutive nomenclature,
IUPAC recommendations prefer the name azane to ammonia: hence
would be named chloroazane in substitutive nomenclature, not
Ammonia as a ligand
Ball-and-stick model of the diamminesilver(I) cation, [Ag(NH3
Ammonia can act as a ligand in
complexes. It is a pure σ-donor, in the
middle of the spectrochemical series, and shows
hard-soft behaviour. For historical reasons, ammonia is named ammine
in the nomenclature of coordination compounds. Some notable
ammine complexes include:
- Tetraamminediaquacopper(II), [Cu(NH3)4(H2O)2]2+,
a characteristic dark blue complex formed by adding ammonia to solution of
copper(II) salts. Known as
- Diamminesilver(I), [Ag(NH3)2]+, the
active species in Tollens' reagent. Formation of this complex can
also help to distinguish between precipitates of the different silver halides:
AgCl is soluble in dilute (2M) ammonia solution,
AgBr is only soluble in concentrated ammonia solution while
is insoluble in aqueous solution of ammonia.
Ammine complexes of chromium(III) were known in the late 19th century, and
formed the basis of Alfred Werner's theory of coordination compounds.
Werner noted that only two isomers (fac- and mer-) of the complex
[CrCl3(NH3)3] could be formed, and concluded
that the ligands must be arranged around the metal ion at the vertices
octahedron. This has since been confirmed by
An ammine ligand bound to a metal ion is markedly more
acidic than a free
ammonia molecule, although deprotonation in aqueous solution is still rare. One
example is the Calomel reaction, where the resulting
amidomercury(II) compound is highly insoluble.
Hg2Cl2 + 2 NH3 → Hg + HgCl(NH2) + NH4+
The energy level diagram depicting levels below 350 cm-1 for ammonia
is in the figure to the right. Energies for J=0 to J=5 are included. Note the
decreasing value of the energy with increasing value of K along a single value
of J. This is due to the fact that ammonia is an oblate symmetric top, meaning
A=B<C where A, B, and C are inversely related to the angular momenta along each
molecular axis. The quantum numbers of each level are provided to the right of
the level and the percentage of the molecule in each state assuming LTE at 30K
is given above each level when statistically significant.
A spectral simulation is provided below for ammonia at 30K. Transitions of the
lines are indicated as follows: (J1,K1) - (J2,K2).
The J1=2,3,4 lines are too close together to distinguish
individually. For clarity, the J1=4 lines have been expanded. A
spectral line list of this simulation is below for J1<6. Frequencies are given
in units of cm-1.
A spectrum is also provided of ammonia a T=300K for comparison. The frequencies
of the simulation have been compared to and match Poynter & Kakar (1975)
These spectra and energy level diagrams were produced using the molecular
spectrum simulator package PGopher. The constants used are taken from Table 4 of
Poynter et al. (1975).
Interstellar formation and destruction
The interstellar abundance for ammonia has been measured for a variety of
environments. The [NH3]/[H2] ratio has been estimated to
range from 10-7 in small dark clouds (cf. Ungerechts et al. 1980) up
to 10-5 in the dense core of the Orion Molecular Cloud (Genzel et al.
1982). Although a total of 18 total production routes have been proposed (see
udfa.net ), the principal formation
mechanism for interstellar NH3 is the reaction:
(1) NH4+ + e- --> NH3 + H
The rate constant, k, of this reaction depends on the temperature of the
environment, with a value of 5.2 x 10-6 at 10K (see Vikor,
Al-Khalili, Danared et al., 1999, A&A, 344, 1027). The rate constant was
calculated from the formula k = a(T/300)B. For the primary formation
reaction, a=1.05 x 10-6 and B = -0.47. Assuming an NH4+
abundance of 3 x 10-7 (van Dishoeck & Black 1986) and an electron
abundance of 10-7 typical of molecular clouds, the formation will
proceed at a rate of 1.6 x 10-9 cm-3s-1 in a
molecular cloud of total density 105cm-3.
All other proposed formation reactions have rate constants of between 2 and 13
orders of magnitude smaller, making their contribution to the abundance of
ammonia relatively insignificant. See Astrochemistry.net
for rate constant citations. As an example of the minor contribution other
formation reactions play, equation (4) below (H2 + NH2 -->
NH3 + H) has a rate constant of 2.2 x 10-15. Assuming H2
densities of 105 and NH2/H2 ratio of 10-7,
this reaction proceeds at a rate of 2.2 x 10-12, more than 3 orders
of magnitude slower that the primary reaction above. Some of the other possible
formation reactions are listed below:
(2) H- + NH4+ --> NH3 + H2
(3) PNH3+ + e- --> P + NH3
(4) H2 + NH2 --> NH3 + H
According to the online database UDFA.net, there are 113 total proposed
reactions leading to the destruction of NH3. Of these, 39 were
tabulated by [Prasad & Huntress (1980) who
compiled extensive tables of the chemistry among C, N, and O compounds. A review
of interstellar ammonia by Ho & Townes (1983) cites the following reactions as
the principal dissociation mechanisms:
(5) NH3 + H3+ --> NH4+ +
(6) NH3 + HCO+ --> NH4+ + CO
with rate constants a of 4.39 x 10-9 (Lininger et al. 1975) and 2.2 x
10-9 (Smith & Adams 1977), respectively. For both reactions, B and
gamma are 0, therefore k = a. Equation (5) and (6) run at a rate of 8.8 x 10-9
and 4.4 x 10-13, respectively. These calculations assumed the given
rate constants and abundances of [NH3]/[H2] = 10-5,
[H3+]/[H2] = 2 x 10-5 (Lepp et al.
1986), [HCO+]/[H2] = 2 x 10-9 (Wooten et al.
1980), and total densities of n = 105, typical of cold, dense,
molecular clouds. Clearly, between these two primary reactions, equation (5) is
the dominant destruction reaction, with a rate ~10,000 times faster than
equation (6). This is due to the relatively high abundance of H3+.
Approximately 83% (as of 2003) of ammonia is used as fertilizers either as its
salts or as solutions. Consuming more than 1% of all man-made power, the
production of ammonia is a significant component of the world energy budget.
Precursor to nitrogenous compounds
Ammonia is directly or indirectly the precursor to most nitrogen-containing
compounds. Practically all synthetic and all inorganic nitrogen compounds are
prepared from ammonia. An important derivative is
acid. This key material is generated via the
Ostwald process by
oxidisation of ammonia with air over a
catalyst at 700 - 850 °C, ~9 atm.
oxide is an intermediate:
NH3 + 2 O2 → HNO3 + H2O
Nitric acid is used for the production of
explosives, and natural organonitrogen other chemical
Minor and emerging uses
Refrigeration - R717
Ammonia's thermodynamic properties made it one of the
refrigerants commonly used prior to the discovery of
Ammonia's toxicity complicates this application. Anhydrous ammonia is widely
used in industrial refrigeration applications because of its high
energy efficiency and low cost. Ammonia is
used less frequently in commercial applications, such as in grocery store
freezer cases and refrigerated displays due to its toxicity.
For remediation of gaseous emissions
Ammonia used to scrub SO2 from the burning of fossil fuels, the
resulting product is converted to ammonium sulfate for use as fertilizer.
Ammonia neutralizes the nitrogen oxides (NOx) pollutants emitted by diesel
engines. This technology, called SCR (selective catalytic reduction), relies on
a vanadia-based catalyst.
As a fuel
Ammonia was used during World War II to power buses in Belgium, and in
engine and solar energy applications prior to 1900. Liquid ammonia was used as
the fuel of the rocket airplane, the X-15. Although not
as powerful as other fuels, it left no soot in the reusable rocket engine and
its density approximately matches that for the oxidizer, liquid oxygen, which
simplified the aircraft's design.
As a vehicle fuel
Ammonia is proposed as a practical and clean alternative to
fuel for internal combustion engines.
The biggest obstacle is the enormous increase in production required since
present production, although the second most produced chemical, is a very small
fraction of world petroleum usage. Ammonia has no more serious issues, as an
alternative vehicle fuel compared to petrol or diesel, including toxicity,
flammability, use in engines, pollution, energy density .It
does require twice the storage volume of petrol/diesel. It can run in existing
engines. It is already widely produced and distributed, and can be manufactured
from renewable energy sources, coal or nuclear power. The main down side is that
overall it is significantly less efficient than batteries. The 60 MW Rjukan dam
in Telemark Norway, was producing ammonia via electrolysis of water for many
years from 1913 producing fertilizer for much of Europe. Ammonia is already
produced, transported and stored on a vast scale. In combination with coal gas
it was used to run 20 buses on 8 routes covering many tens of thousands of miles
with no injuries or engine damage.It can
be used in existing engines with only minor modifications to carburetors /
injectors. If produced from coal, the CO2 can be readily sequestrated.
(the combustion products are nitrogen and water). In 1981 a Canadian company
converted a 1981 Chevrolet Impala to operate using ammonia as fuel.
The use of ammonia as fuel continues to be discussed.
There are prototype solid state processes to use electricity to convert nitrogen
and water directly to ammonia, which are claimed to be cheaper, more efficient
and capable of much smaller scale application ie to otherwise stranded assets
such as remote wind turbines.
The calorific value of ammonia is 22.5 MJ/kg (9690 BTU/lb) which is
about half that of diesel. In a normal engine, in which the water vapor is not
condensed, the calorific value of ammonia will be about 21% less than this
Liquid ammonia is used for treatment of cotton materials, give a properties like
mercerisation using alkalies. And also used for pre-washing of wool.
Solutions of ammonia in water can be applied on the skin to lessen the effects
of acidic animal poisons, especially insect poison and
Ammonia's role in biological systems and human disease
Ammonia is an important source of nitrogen for living systems. Although
atmospheric nitrogen abounds, few living creatures are capable of utilizing this
nitrogen. Nitrogen is required for the synthesis of amino acids, which are the
building blocks of protein. Some plants rely on ammonia and other
nitrogenous wastes incorporated into the soil by decaying matter. Others, such
legumes, benefit from symbiotic relationships with
which create ammonia from atmospheric nitrogen.
Ammonia also plays a role in both normal and abnormal animal
Ammonia is created through normal amino acid metabolism and is toxic in high
ammonia to urea
through a series of reactions known as the
Liver dysfunction, such as that seen in
may lead to elevated amounts of ammonia in the blood (hyperammonemia).
Likewise, defects in the enzymes responsible for the urea cycle, such as
ornithine transcarbamylase, lead to
hyperammonemia. Hyperammonemia contributes to the confusion and
hepatic encephalopathy as well as the
neurologic disease common in people with urea cycle defects and
Ammonia is important for normal animal acid/base balance. After formation of
α-ketoglutarate may be degraded to produce two molecules of
bicarbonate which are then available as buffers for dietary acids. Ammonium
is excreted in the urine resulting in net acid loss. Ammonia may itself diffuse
across the renal tubules, combine with a hydrogen ion, and thus allow for
further acid excretion.
Ammonium ions are a toxic waste
product of the metabolism in
fishes and aquatic invertebrates, it is excreted directly into the water. In
mammals, sharks, and amphibians, it is converted in the
to urea, because it
is less toxic and can be stored more efficiently. In birds, reptiles, and
terrestrial snails, metabolic ammonium is converted into
which is solid, and can therefore be excreted with minimal water loss.
Theoretical role in alternative biochemistry
Ammonia has been proposed as a possible replacement for water as a bodily
solvent in the theoretical alternative biochemistries of
life-forms that do not use carbon for cellular structure and
water as a
solvent to dissolve bodily solutes and allow essential parts of metabolic
processes to occur. It has been suggested that ammonia would be most favorable
for life-forms that live in temperatures below the freezing point of water .
Liquid ammonia as a solvent
See also: Inorganic nonaqueous solvent
Liquid ammonia is the best-known and most widely studied non-aqueous ionizing
solvent. Its most conspicuous property is its ability to dissolve alkali metals
to form highly coloured, electrically conducting solutions containing
solvated electrons. Apart from these
remarkable solutions, much of the chemistry in liquid ammonia can be classified
by analogy with related reactions in aqueous solutions. Comparison of the
physical properties of NH3 with those of water shows that NH3
has the lower melting point, boiling point, density,
dielectric constant and
electrical conductivity; this is due at
least in part to the weaker H bonding in NH3 and the fact that such
bonding cannot form cross-linked networks since each NH3 molecule has
only 1 lone-pair of electrons compared with 2 for each H2O molecule.
The ionic self-dissociation constant of liquid NH3
at −50 °C is approx. 10-33 mol²·l-2.
Solubility of salts
Liquid ammonia is an ionizing solvent, although less so than water, and
dissolves a range of ionic compounds including many
thiocyanates. Most ammonium salts are soluble, and these salts act as
acids in liquid
ammonia solutions. The solubility of
iodide. A saturated solution of
ammonium nitrate contains 0.83 mol solute per
mole of ammonia, and has a
vapour pressure of less than 1 bar even at 25 °C.
Solutions of metals
See also: Solvated electron, metallic solution
Liquid ammonia will dissolve the
metals and other electropositive metals such as
At low concentrations (<0.06 mol/L), deep blue solutions are formed: these
contain metal cations and
solvated electrons, free electrons which are
surrounded by a cage of ammonia molecules.
These solutions are very useful as strong reducing agents. At higher
concentrations, the solutions are metallic in appearance and in electrical
conductivity. At low temperatures, the two types of solution can coexist as
Redox properties of liquid ammonia
See also: Redox.
E° (V, ammonia)
E° (V, water)
Li+ + e− ⇌ Li
K+ + e− ⇌ K
Na+ + e− ⇌ Na
Zn2+ + 2e− ⇌ Zn
NH4+ + e− ⇌ ½ H2
Cu2+ + 2e− ⇌ Cu
Ag+ + e− ⇌ Ag
The range of thermodynamic stability of liquid ammonia solutions is very narrow,
as the potential for oxidation to
E° (N2 + 6NH4+
+ 6e− ⇌ 8NH3), is only
+0.04 V. In practice, both oxidation to dinitrogen and reduction to
are slow. This is particularly true of reducing solutions: the solutions of the
alkali metals mentioned above are stable for several days, slowly decomposing to
the metal amide
and dihydrogen. Most studies involving liquid ammonia solutions are done in
reducing conditions: although oxidation of liquid ammonia is usually slow, there
is still a risk of explosion, particularly if transition metal ions are present
as possible catalysts.
Detection and determination
Ammonia and ammonium salts can be readily detected, in very minute traces, by
the addition of Nessler's solution, which gives a distinct
yellow coloration in the presence of the least trace of ammonia or ammonium
salts. Sulfur sticks are burnt to detect small leaks in
industrial ammonia refrigeration systems. Larger quantities can be detected by
warming the salts with a caustic alkali or with
quicklime, when the characteristic smell of ammonia will be at once
apparent. The amount of ammonia in ammonium salts can be estimated
quantitatively by distillation of the salts with
potassium hydroxide, the ammonia evolved
being absorbed in a known volume of standard
sulfuric acid and the excess of acid then determined volumetrically; or the ammonia may be
absorbed in hydrochloric acid and the
ammonium chloride so formed precipitated as
ammonium hexachloroplatinate, (NH4)2PtCl6.
Ammonia was first detected in interstellar space in 1968, based on
emissions from the direction of the
This was the first polyatomic
molecule to be so detected. The sensitivity of the molecule to a broad range of
excitations and the ease with which it can be observed in a number of regions
has made ammonia one of the most important molecules for studies of
molecular clouds. The
relative intensity of the ammonia lines can be used to measure the temperature
of the emitting medium.
The following isotopic species of ammonia have been detected:
NH3, 15NH3, NH2D, NHD2,
The detection of triply-deuterated ammonia was considered a surprise as
deuterium is relatively scarce. It is thought that the low-temperature
conditions allow this molecule to survive and accumulate.
The ammonia molecule has also been detected in the atmospheres of the
Jupiter, along with other gases like
and helium. The
interior of Saturn may include frozen crystals of ammonia.
Since its interstellar discovery, NH3 has proved to be an invaluable
spectroscopic tool in the study of the interstellar medium. With a large number
of transitions sensitive to a wide range of excitation conditions, NH3
has been widely astronomically detected - it's detection has been reported in
hundreds of journal articles. Listed below is a sample of journal articles that
highlights the range of detectors that have been used to identify ammonia.
Single Antenna Detections
- Wilson et al. 1979 Radio
observations of NH3 from the 100-m Effelsberg Telescope are reported.
The ammonia line is separated into two components - a background ridge and an
unresolved core. The background corresponds well with the locations
- MacDonald et al. 1981Radio
observations of NH3 from the 25-m Chilbolton telescope in England are
presented. Among the observations are 35 new detections of ammonia in HII
regions, HNH2O MASERS, H-H objects, and other objects associated with
star formation. A comparison of emission line widths indicates that turbulent or
systematic velocities do not increase in the central cores of molecular clouds.
- Morris et al. 1973
Microwave radiation from ammonia was observed in several galactic objects
including W3(OH), Orion A, W43, W51, and five sources in the galactic center.
The high detection rate indicates that this is a common molecule in the
interstellar medium and that high density regions are common in the galaxy.
- Torrelles et al. 1985 VLA
observations of NH3 in seven regions with high-velocity gaseous
outflows is presented. Condensations of less than 0.1 pc were detected in L1551,
S140, and Cepheus A. Three individual condensations were detected in Cepheus A,
one of them with a highly elongated shape. These condensations may play an
important role in creating the bipolar outflow in the region.
- Ho et al. 1990Extragalactic
ammonia is imaged using the VLA in IC 342. The hot gas has temperatures above
70K inferred from ammonia line ratios and appears to be closely associated with
the innermost portions of the nuclear bar seen in CO.
- Cesaroni et al. 1994The
authors present VLA measurements of NH3 towards a sample of four
galactic ultracompact HII regions: G9.62+0.19, G10.47+0.03, G29.96-0.02, and
G31.41+0.31. Based upon temperature and density diagnostics, it is concluded
that in general such clumps are likely to be the sites of massive star formation
in an early evolutionary phase prior to the development of an ultracompact HII
- Knacke et al. 1982 The
authors report the detection of absorption at 2.97 microns due to solid ammonia
on interstellar grains in the Becklin-Neugebauer object and probably in NGC
2264-IR as well. This detection helps explain the physical shape of previously
poorly-understood related ice absorption lines.
- Orton et al. 1982 A
spectrum of the disk of Jupiter was obtained from the Kuiper Airborne
Observatory, covering the 100 to 300 cm^−1 spectral range. Analysis of the
spectrum provides information on global mean properties of ammonia gas and an
ammonia ice haze.
- Benson & Meyers 1989 A
total of 149 dark cloud positions were surveyed for evidence of 'dense cores' by
using the (J,K) = (1,1) rotating inversion line of NH3. The cores are
not generally spherically shaped, with aspect ratios ranging from 1.1 to 4 4. It
is also found that cores with stars have broader lines than cores without stars.
- Mebold et al. 1987NH3
has been detected in the Draco Nebula and in one or possibly two molecular
clouds which are associated with the high latitude galactic infrared cirrus. The
finding is significant because they may represent the birth places for the
Population I metallicity B-type stars in the galactic halo which could have been
borne in the galactic disk.
Astronomical observations and research applications
The study of interstellar ammonia has been important to a number of areas of
research in the last few decades. Some of these are delineated below and
primarily involve using ammonia as an interstellar thermometer.
Observations of nearby dark clouds
By balancing collisions and stimulated emission with spontaneous emission, it is
possible to construct a relation between excitation temperature and density.
Moreover, since the transitional levels of ammonia can be approximated by a
2-level system at low temperatures, this calculation is fairly simple. This
premise can be applied to dark clouds, regions suspected of having extremely low
temperatures and possible sites for future star formation. Detections of ammonia
in dark clouds show very narrow lines – indicative not only of low temperatures,
but also of a low level of inner-cloud turbulence. Line ratio calculations
provide a measurement of cloud temperature that is independent of
previously-done CO observations. The ammonia observations were consistent with
CO measurements of rotation temperatures of ~10 K. With this, densities can be
determined, and have been calculated to range between 104 and 105
cm-3 in dark clouds. Mapping of NH3 gives typical clouds
sizes of 0.1 pc and masses near 1 solar mass. These cold, dense cores are the
sites of future star formation.
UC HII regions
Ultra-compact HII regions are among the best tracers of high-mass star
formation. The dense material surrounding UCHII regions is likely primarily
molecular. Since a complete study of massive star formation necessarily involves
the cloud from which the star formed, ammonia is an invaluable tool in
understanding this surrounding molecular material. Since this molecular material
can be spatially resolved, it is possible to constrain the heating/ionizing
sources, temperatures, masses, and sizes of the regions. Doppler-shifted
velocity components allow for the separation of distinct regions of molecular
gas which can trace outflows and hot cores originating from forming stars.
NH3 has been detected in external galaxies, and by simultaneously
measuring several lines, it is possible to directly measure the gas temperature
in these galaxies. Line ratios imply that gas temperatures are warm (~50 K),
originating from dense clouds with sizes of tens of pc. This picture is
consistent with the picture within our Milky Way galaxy – hot dense molecular
cores form around newly-forming stars embedded in larger clouds of molecular
material on the scale of several hundred pc (giant molecular clouds; GMCs).
Toxicity and storage information
Hydrochloric acid sample releasing HCl fumes which are reacting with ammonia
fumes to produce a white smoke of ammonium chloride.
The toxicity of ammonia solutions does not usually cause problems for humans and
other mammals, as a specific mechanism exists to prevent its build-up in the
bloodstream. Ammonia is converted to
carbamoyl phosphate by the enzyme
carbamoyl phosphate synthetase,
and then enters the urea cycle to be either incorporated into
or excreted in the urine. However fish and amphibians lack this mechanism, as they can usually
eliminate ammonia from their bodies by direct excretion. Ammonia even at dilute
concentrations is highly toxic to aquatic animals, and for this reason it is
classified as dangerous for the
environment. Ammonium compounds should never be allowed to come in contact
with bases (unless in an intended and contained reaction), as dangerous
quantities of ammonia gas could be released.
Solutions of ammonia (5–10% by weight) are used as household cleaners,
particularly for glass. These solutions are irritating to the eyes and
mucous membranes (respiratory and digestive tracts), and to a lesser extent
the skin. Caution should be used that the chemical is never mixed into any
liquid containing bleach, or a poisonous gas may result. Mixing with
products or strong oxidants, for example household
bleach can lead
to hazardous compounds such as chloramines.
Laboratory use of ammonia solutions
The hazards of ammonia solutions depend on the concentration: "dilute" ammonia
solutions are usually 5–10% by weight (<5.62 mol/L); "concentrated" solutions
are usually prepared at >25% by weight. A 25% (by weight) solution has a density
of 0.907 g/cm³, and a solution which has a lower density will be more
concentrated. The European Union classification of ammonia
solutions is given in the table.
The ammonia vapour from concentrated ammonia solutions is severely irritating to
the eyes and the respiratory tract, and these solutions should only be handled
in a fume hood. Saturated ("0.880") solutions can develop a significant pressure
inside a closed bottle in warm weather, and the bottle should be opened with
care: this is not usually a problem for 25% ("0.900") solutions.
Ammonia solutions should not be mixed with
toxic and/or explosive products are formed. Prolonged contact of ammonia
silver, mercury or
can also lead to explosive products: such mixtures are often formed in
qualitative chemical analysis, and should
be lightly acidified but not concentrated (<6%w/v) before disposal once the test
Laboratory use of anhydrous ammonia (gas or liquid)
Anhydrous ammonia is classified as toxic (T) and dangerous for
the environment (N). The gas is flammable (autoignition temperature: 651 °C) and
can form explosive mixtures with air (16–25%). The
permissible exposure limit (PEL) in
the United States is 50 ppm
(35 mg/m³), while the
IDLH concentration is estimated at 300 ppm. Repeated exposure to ammonia
lowers the sensitivity to the smell of the gas: normally the odour is detectable
at concentrations of less than 50 ppm, but desensitized individuals may not
detect it even at concentrations of 100 ppm. Anhydrous ammonia corrodes
alloys, and so
brass fittings should not be used for handling the gas. Liquid ammonia can
also attack rubber and certain plastics.
Ammonia reacts violently with the halogens.
Nitrogen triiodide, a
high explosive, is formed when ammonia comes in contact with
causes the explosive polymerization of
ethylene oxide. It also forms explosive
fulminating compounds with compounds of
and with stibine.
Violent reactions have also been reported with
acetaldehyde, hypochlorite solutions,
potassium ferricyanide and
The U. S.
and Health Administration (OSHA) has set a 15-minute exposure limit for
gaseous ammonia of 35 ppm by volume in the environmental air and an 8-hour
exposure limit of 25 ppm by volume.
Exposure to very high concentrations of gaseous ammonia can result in lung
damage and death.
Although ammonia is regulated in the United States as a non-flammable gas, it
still meets the definition of a material that is toxic by inhalation and
requires a hazardous safety permit when transported in quantities greater than
13,248 L (3,500 gallons).
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