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Ammonia
Ammonia is a chemical compound whose
molecule consists of one atom of
nitrogen (N) and three atoms of General
hydrogen (H) with the formula NH3 and Name Ammonia
the structure: Chemical formula NH3
[image:ammonia.png] Appearance Colourless gas
Physical
The molecule is not flat, instead it Formula weight 17.0 amu
has the shape of a tetrahedron with Melting point 195 K (-78 ¡C)
one empty corner. In solution it forms
the positively charged ammonium ion Boiling point 240 K (-33 ¡C)
NH4+ with a hydrogen atom on all four 8.0 ×103
corners of the tetrahedron. Density
kg/m3 (liquid)
At standard temperature and pressure, Solubility 46 g in 100g water
ammonia is a gas with a characteristic Thermochemistry
pungent smell; its main uses are in ΔfH0gas -45.9 kJ/mol
the production of fertilizers,
explosives and polymers. ΔfH0liquid -40.2 kJ/mol
Ammonia is very well suited as a ΔfH0solid ? kJ/mol
refrigerant, since it is readily
liquified under pressure, and was used S0gas, 1 bar 192.77 J/moláK
in virtually all refrigeration units
prior to the advent of freons. Since S0liquid, 1 bar ? J/moláK
the implication of freons as major
greenhouse gases during the 1990s, S0solid ? J/moláK
ammonia is again seeing increasing use
as a refrigerant. Safety
Dangerous. Symptoms
Ammonia is found in small quantities include nausea &
as the carbonate in the atmosphere, Ingestion vomiting; damage to
being produced from the putrefaction lips, mouth and
of nitrogenous animal and vegetable esophagus.
matter; ammonium salts are also found Vapours are
in small quantities in rain-water,
whilst ammonium chloride Inhalation extremely
irritating and
(sal-ammoniac) and ammonium sulfate corrosive.
are found in volcanic districts; and
crystals of ammonium bicarbonate have Concentrated
been found in Patagonian guano. Skin solutions may
Ammonium salts also are found produce severe
distributed through all fertile soil, burns and necrosis.
in sea-water, and in most plant and May cause permanent
animal liquids, and also in urine. Eyes damage, even in
small quantities.
More info Hazardous Chemical
Database
SI units were used where possible.
Unless otherwise stated, standard
conditions were used.
Production
Because of its many uses, ammonia is one of the most highly-produced
inorganic chemicals. Before the start of WWI most ammonia was obtained by
the dry distillation of nitrogenous vegetable and animal products; by the
reduction of nitrous acid and nitrites with nascent hydrogen; and also by
the decomposition of ammonium salts by alkaline hydroxides or by unslaked
lime (quicklime), the salt most generally used being the chloride
(sal-ammoniac) thus
2NH4Cl + 2CaO = CaCl2 + Ca(OH)2 + 2NH3.
It was also obtained by decomposing magnesium nitride (Mg3N2) with water,
Mg3N2 + 6H2O = 3Mg(OH)2 + 2NH3.
Today the Haber process is the most important method for production of
ammonia. The main advantage of the Haber process is that relatively cheap
nitrogen and hydrogen gas are the primary feedstocks. They are reacted over
an iron catalyst at high pressure (3000 psi or 20 MPa) and temperature
(500¡C) to produce the ammonia.
Properties
Ammonia is a colourless gas possessing a characteristic pungent smell and a
strongly alkaline reaction; it is lighter than air, its density being 0.589
times that of air. It is easily liquefied and the liquid boils at -33.7 ¡C,
and solidifies at -75¡C. to a mass of white crystals. Liquid ammonia
possesses strong ionizing powers, and solutions of salts in liquid ammonia
have been much studied.
It is extremely soluble in water, one volume of water at 0¡C and normal
pressure absorbs 1148 volumes of ammonia. All the ammonia contained in an
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aqueous solution of the gas may be expelled by boiling. The aqueous solution
of ammonia is very basic in its reactions, and since it is a weak
electrolyte, one must assume the solution to contain a certain amount of
ammonium hydroxide NH4OH, although it is probably chiefly composed of a
solution of ammonia in water.
It does not support combustion, and it does not burn readily unless mixed
with oxygen, when it burns with a pale yellowish-green flame. However it can
form an explosive mixture with air.
Salts
One of the most characteristic properties of ammonia is its power of
combining directly with acids to form salts; thus with hydrochloric acid it
forms ammonium chloride (sal-ammoniac); with nitric acid, ammonium nitrate,
etc. It is to be noted that H. B. Baker (Journal of Chem. Soc., 1894, lxv.
p. 612) has shown that perfectly dry ammonia will not combine with perfectly
dry hydrochloric acid, moisture being necessary to bring about the reaction.
The salts produced by the action of ammonia on acids are known as the
ammonium salts and all contain the compound radical ammonium (NH4). Numerous
attempts have been made to isolate this radical, but so far none have been
successful. By the addition of sodium amalgam to a concentrated solution of
ammonium chloride, the so-called ammonium amalgam is obtained as a spongy
mass which floats on the surface of the liquid; it decomposes readily at
ordinary temperatures into ammonia and hydrogen; it does not reduce silver
and gold salts, a behaviour which distinguishes it from the amalgams of the
alkali metals, and for this reason it is regarded by some chemists as being
merely mercury inflated by gaseous ammonia and hydrogen. M. le Blanc has
shown, however, that the effect of ammonium amalgam on the magnitude of
polarization of a battery is comparable with that of the amalgams of the
alkali metals.
Ammonium bromide, NH4Br, can be prepared by the direct action of bromine on
ammonia. It crystallizes in colourless prisms, possessing a saline taste; it
sublimes on heating and is easily soluble in water. On exposure to air it
gradually assumes a yellow colour and becomes acid in its reaction.
Ammonium chloride, NH4Cl. (See sal-ammoniac.)
Ammonium fluoride, NH4F, may be obtained by neutralizing ammonia with
hydrofluoric acid. It crystallizes in small prisms, having a sharp saline
taste, and is exceedingly soluble in water. It decomposes silicates on being
heated with them.
Ammonium iodide, NH4I, can be prepared by the action of hydriodic acid on
ammonia. It is easily soluble in water, from which it crystallizes in cubes,
and also in alcohol. It gradually turns yellow on standing in moist air,
owing to decomposition with liberation of iodine.
Ammonium chlorate, NH4ClO3, is obtained by neutralizing chloric acid with
either ammonia or ammonium carbonate, or by precipitating barium, strontium
or calcium chlorates with ammonium carbonate. It crystallizes in small
needles, which are readily soluble in water, and on heating, decompose at
about 102¡C, with liberation of nitrogen, chlorine and oxygen. It is soluble
in dilute aqueous alcohol, but insoluble in strong alcohol.
Ammonium carbonates. The commercial salt was formerly known as sal-volatile
or salt of hartshorn and was formerly obtained by the dry distillation of
nitrogenous organic matter such as hair, horn, decomposed urine, etc., but
is now obtained by heating a mixture of sal-ammoniac, or ammonium sulfate
and chalk, to redness in iron retorts, the vapours being condensed in leaden
receivers. The crude product is refined by sublimation, when it is obtained
as a white fibrous mass, which consists of a mixture of hydrogen ammonium
carbonate, NH4.HCO3, and ammonium carbamate, NH2COONH4, in molecular
proportions; on account of its possessing this constitution it is sometimes
called ammonium sesquicarbonate. It possesses a strong ammoniacal smell, and
on digestion with alcohol the carbamate is dissolved and a residue of
ammonium bicarbonate is left; a similar decomposition taking place when the
sesquicarbonate is exposed to air. Ammonia gas passed into a strong aqueous
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solution of the sesquicarbonate converts it into normal ammonium carbonate,
(NH4)2CO3, which can be obtained in the crystalline condition from a
solution prepared at about 30¡C. This compound on exposure to air gives off
ammonia and passes back to ammonium bicarbonate.
Ammonium bicarbonate, NH4HCO3, is formed as shown above and also by passing
carbon dioxide through a solution of the normal compound, when it is
deposited as a white powder, which has no smell and is only slightly soluble
in water. The aqueous solution of this salt liberates carbon dioxide on
exposure to air or on heating, and becomes alkaline in reaction. The aqueous
solutions of all the carbonates when boiled undergo decomposition with
liberation of ammonia and of carbon dioxide:
NH4HCO3 --> NH3 + H2O + CO2
It is therefore occassionally used as baking powder, e.g. for gingerbread.
Ammonium nitrate, NH4NO3, is prepared by neutralizing nitric acid with
ammonia, or ammonium carbonate, or by double decomposition between potassium
nitrate and ammonium sulfate. It can be obtained in three different
crystalline forms, the transition points of which are 35¡C, 83¡C and 125¡C.
It is easily soluble in water, a considerable lowering of temperature taking
place during the operation; on this account it is sometimes used in the
preparation of freezing mixtures. On gentle heating, it is decomposed into
water and nitrous oxide. P. E. M. Berthelot in 1883 showed that if ammonium
nitrate be rapidly heated the following reaction takes place with explosive
violence:--2NH4NO3 = 4H2O + 2N2 + O2. In combination with gasoline it is a
widely used explosive.
Ammonium nitrite, NH4NO2, is formed by oxidizing ammonia with ozone or
hydrogen peroxide; by precipitating barium or lead nitrites with ammonium
sulfate, or silver nitrite with ammonium chloride. The precipitate is
filtered off and the solution concentrated. It forms colourless crystals
which are soluble in water and decompose on heating, with the formation of nitrogen.
Ammonium phosphates. The normal phosphate, (NH4)3PO4,is obtained as a
crystalline powder, on mixing concentrated solutions of ammonia and
phosphoric acid, or on the addition of excess of ammonia to the acid
phosphate (NH4)2HPO4. It is soluble in water, and the aqueous solution on
boiling loses ammonia and the acid phosphate NH4H2PO4 is formed. Diammonium
hydrogen phosphate, (NH4)2HPO4, is formed by evaporating a solution of
phosphoric acid with excess of ammonia. It crystallizes in large transparent
prisms, which melt on heating and decompose, leaving a residue of
metaphosphoric acid, (HPO3). Ammonium dihydrogen phosphate, NH4.H2PO4, is
formed when a solution of phosphoric acid is added to ammonia until the
solution is distinctly acid. It crystallizes in quadratic prisms.
Ammonium sodium hydrogen phosphate, NH4.NaHPO4.4H2O. (See microcosmic salt.)
Ammonium sulfate (NH4)2SO4 is prepared commercially from the ammoniacal
liquor of gas-works and is purified by recrystallization. It forms large
rhombic prisms, has a somewhat saline taste and is easily soluble in water.
The aqueous solution on boiling loses some ammonia and forms an acid
sulfate. It is used largely as an artificial manure, and also for the
preparation of other ammonium salts.
Ammonium persulfate (NH4)2S2O8 has been prepared by H. Marshall (Jour. of
Chem. Soc., 1891, lix. p. 777) by the method used for the preparation of the
corresponding potassium salt (see sulfur). It is very soluble in cold water,
a large fall of temperature accompanying solution. It is a mighty strong
oxidizing agent.
Ammonium sulfide, (NH4)2S, is obtained, in the form of micaceous crystals,
by passing sulfuretted hydrogen mixed with a slight excess of ammonia
through a well-cooled vessel; the hydrosulfide NH4.HS is formed at the same
time. It dissolves readily in water, but is probably partially dissociated
in solution. The hydrosulfide NH4.HS can be obtained as a white solid, by
mixing well-cooled ammonia with a slight excess of sulfuretted hydrogen.
According to W. P. Bloxam (Jour. of Chem. Soc., 1895, lxvii. p. 283), if
sulfuretted hydrogen is passed into strong aqueous ammonia at ordinary
temperature, the compound (NH4)2S.2NH4HS is obtained, which, on cooling to
0¡C and passing more sulfuretted hydrogen, forms the compound
(NH4)2S.12NH4HS. An ice-cold solution of this substance kept at 0¡C and
having sulfuretted hydrogen continually passed through it gives the
hydrosulfide. Several complex polysulfides of ammonium have been isolated,
for details of which see Bloxam's paper quoted above. Compounds are known
which may be looked upon as derived from ammonia by the replacement of its
hydrogen by the sulfo-group (HSO3); thus potassium ammon-trisulfonate,
N(SO3K)3.2H2O, is obtained as a crystalline precipitate on the addition of
excess of potassium sulfite to a solution of potassium nitrite, KNO2 +
3K2SO3 + 2H2O = N(SO3K)3 + 4KHO. It can be recrystallized by solution in
alkalies. On boiling with water, it is converted, first into the disulfonate
NH(SO3K)2 thus, N(SO3K)3 + H2O = NH(SO3K)2 + KHSO4, and ultimately into the
monosulfonate NH2.SO3K. The disulfonate is more readily obtained by
moistening the nitrilosulfonate with dilute sulfuric acid and letting it
stand for twenty-four hours, after which it is recrystallized from dilute
ammonia. It forms monosymmetric crystals which by boiling with water yield
amidosulfonic acid. (See also E. Divers, Jour. of Chem. Soc., 1892, lxi. p.
943.) Amidosulfonic acid crystallizes in prisms, slightly soluble in water,
and is a stable compound.
Other compounds
Ammonia finds a wide application in organic chemistry as a synthetic
reagent; it reacts with alkyl iodides to form amines, with esters to form
acid amides, with halogen fatty acids to form amino acids; while it also
combines with isocyanic esters to form alkyl ureas and with the mustard oils
to form alkyl thioureas. Aldehydes also combine directly with ammonia.
Ammonia gas has the power of combining with many substances, particularly
with metallic halides; thus with calcium chloride it forms the compound
CaCl2.8NH3, and consequently calcium chloride cannot be used for drying the
gas. With silver chloride it forms two compounds -- one, AgCl.3NH3 at
temperatures below 15¡C; the other, 2AgCl.3NH3 at temperatures above 20¡C.
On heating these substances, ammonia is liberated and the metallic chloride
remains. It was by the use of silver chloride ammonia compounds that in 1823
Michael Faraday was first able to liquefy ammonia. It can be shown by
Isambert's results that the compound AgCl.3NH3 cannot be formed above 20¡C,
by the action of ammonia on silver chloride at atmospheric pressure; whilst
2AgCl.3NH3, under similar conditions, cannot be formed above about 68¡C.
Liquid ammonia is used for the artificial preparation of ice. It readily
dissolves sodium and potassium, giving in each case a dark blue solution. At
a red heat ammonia is easily decomposed into its constituent elements, a
similar decomposition being brought about by the passage of electric sparks
through the gas. Chlorine takes fire when passed into ammonia, nitrogen and
hydrochloric acid being formed, and unless the ammonia be present in excess,
the highly explosive nitrogen trichloride NCl3 is also produced.
With iodine it reacts to form nitrogen iodide. This compound was discovered
in 1812 by Bernard Courtois, and was originally supposed to contain nitrogen
and iodine only, but in 1840 R. F. Marchand showed that it contained
hydrogen, whilst R. Bunsen showed that no oxygen was present. As regards its
constitution, it has been given at different times the formulae NI3, NHI2,
NH2I, N2H3I3, &c., these varying results being due to the impurities in the
substance, owing to the different investigators working under unsuitable
conditions, and also to the decomposing action of light. F. D. Chattaway
determined its composition as N2H3I3, by the addition of excess of standard
sodium sulfite solution, in the dark, and subsequent titration of the excess
of the sulfite with standard iodine. The constitution has been definitely
determined by O.Silberrad (Jour. of Chem. Soc., 1905, lxxxvii. p. 55) by the
interaction of nitrogen iodide with zinc ethyl, the products of the reaction
being triethylamine and ammonia; the ammonia liberated was absorbed in
hydrochloric acid, and 95% of the theoretical amount of the ammonium
chloride was obtained. On these grounds O. Silberrad assigns the formula
NH3.NI3 to the compound, and explains the decomposition as taking place,
2NH3.NI3 + 6Zn(C2H5)2 = 6ZnC2H5.I + 2NH3 + 2N(C2H5)3.
The hydrogen in ammonia is capable of replacement by metals, thus magnesium
burns in the gas with the formation of magnesium nitride Mg3N2, and when the
gas is passed over heated sodium or potassium, sodamide, NaNH2, and
potassamide, KNH2, are formed.
Detection
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.
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 sodium or 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 chlorplatinate, (NH4)2PtCl6.
History
Salts of ammonia have been known from very early times; thus the term
Hammoniacus sal appears in the writings of Pliny, although it is not known
whether the term is identical with the more modern sal-ammoniac.
In the form of sal-ammoniac, ammonia was known, however, to the alchemists
as early as the 13th century, being mentioned by Albertus Magnus, while in
the 15th century Basil Valentine showed that ammonia could be obtained by
the action of alkalies 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 spirits of hartshorn
was applied to ammonia.
Gaseous ammonia was first isolated by J. Priestley in 1774 and was termed by
him "alkaline air." In 1777 Karl Wilhelm Scheele showed that it contained
nitrogen, and C. L. Berthollet, in about 1785, ascertained its composition.
The Haber process to produce ammonia from the nitrogen contained in the air
was developed by Fritz Haber and Carl Bosch in 1909 and patented in 1910. It
was first used on an industrial scale by the Germans during WWI. The ammonia
was used to produce explosives to sustain their war effort.
Etymology of "Ammonia"
In classical times, sal ammoniac was discovered by accident through burning
the dung of camels in the temple of Jupiter Ammon at Siwa oasis in Libya.
"Ammonia" is a genus name in the Foraminifera (marine planktonic protozoa
with a calcium carbonate shell, whose remains have contributed to limestone
and chalk deposits), and "ammonites" are an extinct group of cephalopod
whose fossil shells are abundant from the Paleozoic. In both cases, the
shell is formed of a series of chambers, arranged in a spiral, and the name
is given for the "Horn of Ammon", the ram's horns that the god by whose
temple the ammoniacal camel dung was to be found (see above) was supposed to
have had
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http://en.wikipedia.org/wiki/Ammonia
^ easier to read here.
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(Rated as: UTF PM Function)
(Hive Bee)
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Ammonium sulfide and hydrogen sulfides as crystals? In my experience, passing H2S through NH3aq, forms the sulfide as a thin-ish bright amber yellow oil with a foul and nauseating smell, slightly reminiscent of rotten eggs, yet different, and utterly foul, it seems to have a wierd effect too, I've smelled it, when I used to make stinkbombs at home, and smelling or skin contact with traces of it makes my feel all jittery and nervous, for quite some time after.
Any idea what causes this? possible residual H2S causing very slight toxicity? direct toxicity of the sulfide itself? I know that its toxic by skin contact to atleast some degree, but only trace amounts have ever come into skin contact.
Over unity and rivers of blood for all!
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(Rated as: insignificant)
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No 536317
Hydrigen Sulfide (purity 99.25% minimum)
Typical impurities: propylene, propane, carbon dioxide, nitrogen, methane, ethane, water, all less than 1%. May contain carbonyl sulfide, carbon disulfide, and/or mercaptans all less than 1%.
Exposure Limits: OSHA* ACC: 20 ppm (30 mg/m3) with Acceptable Maximum Peak above ACC of 50 ppm (75 mg/m3) 10 minute/day if no other measurable exposure occurs. ACGIH TLV-TWA 10 ppm; TLV-STEL 15 ppm (15 minutes, not more than 4 times per 8 hr day. NIOSH recommendation: C=10 ppm, [10 minutes]. Manufacturer's Recommendation to avoid eye irritation 10 ppm, avg. NIOSH IDLH [30 minutes]=100 ppm, {NOTE: IDLH is not working level; it is deemed a concentration from which escape may be made in 30 minutes without injury or irreversible health effects and without deleterious / severe impediment to escape--e.g. irritation.}.
Since combustion of hydrogen sulfide in air produces water vapor and sulfur dioxide (SO2) mixed with large volumes of nitrogen the exposure limits for SO2 are shown here for information purposes only: Sulfur Dioxide -OSHA*=5 ppm TWA (13 mg/m3). ACGIH TLV=2 ppm; STEL=5 ppm; NIOSH TWA=2 ppm. NIOSH IDLH=100 ppm. See above for explanation of IDLH.
Effects of overexposure: Deadening of sense of smell, headache, eye and breathing passage irritation or pain, dizziness, coughing, loss of consciousness, respiratory paralysis, coma, death. Target Organs are the Eyes and Airway (respiratory) Membranes (irritation) and the Nervous System (olfaction, dizziness, unconsciousness, interference with respiratory centers resulting in stopped breathing, gastrointestinal effects, etc.). Other effects have appeared in the literature associated with this gas (generally involving exposures to mixtures of hydrogen sulfide and other toxics). These effects include fatigue, irritability, incoordination, nausea, vomiting, chemical pneumonia, edema, hyperpnea, low blood pressure, convulsions, hematuria, loss of appetite, sleep disturbance, conjunctivitis, excitement, sensation of dryness or pain in eye, nose, throat, chest, etc.