Chemical elements
  Nitrogen
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    Chemical Properties
    Ammonia
      Physical Properties of Ammonia
      Chemical Properties of Ammonia
      Liquid Ammonia
      Aqueous Ammonia
      Ammonia in Solutions
      Detection and Estimation
      Ammonia Equilibrium
    Hydroxylamine
    Hydrazine
    Azoimide
    Nitric Acid

Chemical Properties of Ammonia





Chemical Reactions of Ammonia

Ammonia is not oxidised by atmospheric oxygen or mild oxidising agents at ordinary temperatures, and even when ignited in air, the heat of combination is not sufficient to maintain a flame. It will, however, burn vigorously in oxygen with a greenish yellow flame, forming water and nitrogen. A slow combustion is effected by passing sparks through a mixture of ammonia and air. In this connection it is to be noted that ammonia itself is almost completely decomposed by the passage of electric sparks. It is also completely decomposed by ultra-violet light. When the gas, sealed into quartz tubes, is exposed to this light for six hours, the decomposition is so complete that no undecomposed gas can be detected by Nessler's reagent.

Oxidation proceeds at a much lower temperature in the presence of catalysts, and may then take a different course, leading to the formation of oxides of nitrogen, nitrous and nitric acids, and their ammonium salts. Among such catalysts are nickel, iron, copper, also manganates and permanganates. The most effective and useful catalyst is platinum; the finely divided metal accelerates the oxidation to water and nitrogen, while the smooth or compact metal leads to the oxides. This reaction has been studied from the point of view of technical applications. The first product of the reaction is probably nitric oxide, which is converted into the peroxide and then by water into the acids.

The "preferential" combustion of the hydrogen in ammonia when a deficiency of oxygen is used may be represented by the equation:

4NH3+3O2 = 2N2+6H2O,

which represents only the end-products. But the mechanism is considerably more complex, as is shown by the fact that free hydrogen is commonly found when there is a deficiency of oxygen, and oxides of nitrogen when there is an excess.

In the combustion of the theoretical quantities given above, 59 per cent, of nitrogen and 41 per cent, of hydrogen are found in the gaseous products. This is accounted for by the intermediate formation of di-imide or hydrazine. Thus:

2NH3+O2 = N2H2+2H2O;
4NH3+O2 = 2N2H4+2H2O.

The results of the explosion of ammonia with electrolytic gas are somewhat similar. When the ratio of the gas to ammonia is higher than 3, the ammonia is completely decomposed. When the ratio of ammonia to the gas is lower than 1.6, oxides of nitrogen are formed.

Gaseous ammonia is oxidised to water and nitrogen by many oxides, as, for example, by copper oxide. Chlorine monoxide and iodine pent-oxide oxidise it with separation of the halogen, selenium dioxide with separation of selenium. Ammonia is also oxidised by oxides of nitrogen: when the gases are sparked or heated together, the hydrogen is oxidised and nitrogen is separated. The reaction in the case of nitrogen dioxide is violent even at lower temperatures.

In aqueous solution, ammonia is oxidised at the anode during the electrolysis of its salts, also by powerful oxidising agents such as chlorine, bromine, ozone, hydrogen peroxide, chromates, and permanganates. The main product of oxidation is usually nitrogen. Under special conditions, and particularly in the presence of catalysts, nitrites or nitrates may be formed. This is the case, for example, during electrolytic oxidation in the presence of cupric oxide and alkali, or during oxidation by atmospheric oxygen in the presence of cupric oxide, also of metallic copper, iron, or zinc, or of hydrogen peroxide, or of moist ozone:

4NH3+7O3 = NH4NO3+NH4NO2+7O2+2H2O.

The more electro-negative non-metals, chlorine, bromine, and sulphur, readily combine with the hydrogen of ammonia. Thus chlorine oxidises ammonia to free nitrogen and ammonium chloride. In the well-known lecture experiment by which the composition of ammonia is demonstrated, an excess of the latter is added to a measured volume of chlorine, and one-third of this volume of nitrogen is formed. Thus:

2NH3+3Cl2 = N2+6HCl,
followed by
6HCl+6NH3 = 6NH4Cl.

This reaction probably proceeds through the intermediate formation of NCl3. Since the reaction takes place when both gases are anhydrous, the next step is not necessarily the hydrolysis of this compound, but rather consists in the intermediate production of NCl3.HCl3, which then reacts with hydrochloric acid, losing chlorine. Thus:

2NH3+3Cl2 - 6HCl+N2;
6NH3+9Cl2 = 6NCl3+18HCl;
3NCl3.HCl+9HCl = 3NH4Cl+9Cl2.

Iodine in aqueous solution gives NH3.NI3. The vapour of phosphorus when passed with ammonia through a red-hot tube gives nitrogen and phosphine. Sulphur vapour gives nitrogen and ammonium sulphide. Carbon at a red heat yields ammonium cyanide. Boron combines with the nitrogen to form the very stable nitride, BN. The metals gold, platinum, silver, and copper decompose ammonia into its elements. Iron forms a nitride in addition, as also do chromium and other metals.


Replacement of Hydrogen

Since ammonia is the hydride of a strongly electro-negative element, it might be supposed that at least one or perhaps two of the hydrogen atoms would exhibit an acidic character and be replaceable by metals. The resulting compounds, some of which are known, are called amides or imides respectively, and may be considered as salts when the metal is strongly electro-positive, although the same groups play a positive part in the amides and imides of non-metals, and of inorganic and organic acid radicals.

The amides of the metals may be formed under the following conditions:
  1. By passing ammonia over the metal heated to a suitable temperature.

    Thus, sodamide is prepared by passing a current of dry ammonia over sodium heated to between 300° and 400° C. in a nickel vessel:

    2NH3+2Na = 2NH2Na+H2.

    It is a white solid which melts at 155° C. and is hydrolysed by water, giving ammonia and sodium hydroxide. The amides of potassium, rubidium, and barium may be made by this method.
  2. By warming a solution of the metal in liquid ammonia. In these reactions an ammine, such as KNH3, Ca(NH3)4, is first formed. On warming, substitution of hydrogen takes place. In this manner the following amides have been prepared: NaNH2, LiNH2 and Ca(NH2)2, KNH2.
  3. By the action of ammonia on the hydrides. The amides of the alkali metals and barium have been prepared in this manner.
  4. Amides may also be prepared by double decomposition in liquid ammonia, notably silver amide, AgNH2. Dry ammonia reacts with halides of many non-metals and their oxides (halanhydrides and halides of oxy-acids) as it does with those of organic acid radicals, giving amides and ammonium chloride. On further heating, these amides may pass into imides and nitrides, or these changes may take place quickly without any possibility of isolation of the intermediate compounds.

Reactions with Acids

Ammonia forms salts in several ways, viz.:
  1. By direct addition of NH3 to a hydracid or oxy-acid. The formation of salts by this method is best exemplified by the halides. Thus:

    NH3+HX = NH4X.

    The solid salts are anhydrous; they can be sublimed with dissociation in the vapour phase and recombination during condensation. Dibasic acids in excess give acid salts, such as NH4HSO4. Hydrogen fluoride behaves in this respect as a dibasic acid, and NH4F.HF is formed by the extensive hydrolysis of the salt, NH4F, on warming. Gaseous ammonia combines with gaseous carbon dioxide to give ammonium carbamate, NH2-CO-ONH4. With hydrogen sulphide, crystals of (NH4)2S are obtained at -18° C., and of NH4HS at 0° C. Compounds of ammonia with weak acids are conveniently prepared by interaction in organic solvents. Thus, NH4HS has been prepared in dry alcohol and ether, and the crystalline alcoholate, (NH4)2S.C2H5OH, in alcohol.
  2. By the neutralisation of aqueous ammonia with acids, and subsequent evaporation and crystallisation of the salts. The majority of ammonium salts may be prepared in this manner in aqueous solution.
  3. By double decomposition, with separation of a less soluble salt from the resulting salt-pair. Ammonium salts crystallise usually in an anhydrous condition on evaporation, and are more or less decomposed with loss of ammonia on long boiling.

The Decomposition of Ammonia

As has already been mentioned, the decomposition of ammonia into its elements, like its formation from these, is greatly accelerated by the presence of solid surfaces. Incandescent wires of various metals are particularly effective in this respect, no doubt on account of the fact that their surfaces adsorb hydrogen in an active, probably an atomic, form. This decomposition presents some points of interest in connection with the theories of heterogeneous, or gas-solid, surface reactions. On a tungsten wire at 856° C. the reaction is unimolecular at very low pressures - that is, it conforms to the requirements of the equation



The active mass of the gas is that which is adsorbed on the surface, and it may be assumed that as long as this is not completely covered the active area is proportional to the pressure, i.e. to the concentration. At higher pressures - 50 mm. or more - the reaction is found to be of zero order; in other words, it is independent of the pressure in the gas phase; under these conditions the surface is supposed to be completely covered.

The decomposition of ammonia on a heated platinum wire at about 1000° C. proceeds with great velocity at first, then falls off rapidly as the hydrogen accumulates, and the reaction proceeds at a rate inversely proportional to its partial pressure. Thus:



This effect is not due to the active mass of the hydrogen in the homogeneous gas phase, but is connected with the fact that hydrogen is strongly adsorbed even at low pressures on the surface, which soon becomes saturated, and thus prevents the ammonia molecules from coming into direct contact with the reactive surface except through a layer of this hydrogen.
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