Experiment 4.-Remove the stopper downwards from a bottle of ammonia, and afterwards immerse the mouth of the bottle in water. The water will rise rapidly, and if the bottle was quite full of gas, will fill it. If a piece of red litmus-paper is dipped in the water of the bottle its blue colour will return. Ammonia is soluble in water, and its solution is a base.

One volume of water at freezing point will dissolve 1149, and at ordinary temperatures about 700 volumes of the gas, so that it is even more soluble than hydrochloric acid. A concentrated solution of it may be prepared like that of hydrochloric acid (page 113). It is sold under the name of solution of ammonia, or "liquor ammoniæ," and has a specific gravity of 0.88. Weaker solutions are sometimes called "hartshorn," or "spirits of hartshorn," because they were formerly obtained by the distillation of horn.

Experiment 5.-Ammonia gas, like hydrochloric acid, can easily be prepared by gently heating its concentrated solution. By prolonged boiling all trace of ammonia may be removed from water.

Experiment 6.-The direct combination of ammonia and hydrochloric acid gases has already been noticed (page 13). The experiment may be repeated with equal volumes of the pure gases. Sal-ammoniac or ammonium chloride is formed: NH3+ HCl = NH,CI.

Experiment 7.-Carefully neutralize a dilute solution of ammonia with nitric acid, so that it has no effect on either blue or red litmus. The solution will contain a salt called ammonium nitrate, which may be obtained in the dry state by evaporation :

NH, + HNO3 = NH1NO3.

Ammonium compounds.-The compounds of ammonia are so similar in many of their properties to those of potassium and sodium, that they are generally supposed to be analogous in structure, and to contain a sort of compound metal called ammonium, NH,, which has recently been isolated. When the gas dissolves in water it is assumed that the hydrate of this quasi metal is formed: NH, H2O = NH,HO. The theory cannot be proved, but it is probable and useful. The

following table shows the analogy which on this view exists between potassium and ammonium salts:

It will be seen that to represent the composition of any ammonium salt we need only take the formula of the corresponding potassium or sodium compound and write NH instead of K or Na. It must be remarked that although solution of ammonia may contain the hydrate NH HỌ, no proof that such is the case has been obtained.

Experiment 8.-Put about an ounce of mercury into a mortar and press into it with the pestle a few fragments of clean dry sodium. The two metals will combine with a flash of light, and a semi-fluid mass called sodium amalgam will be obtained. When cold, throw this into a tumbler about one-third full of a cold saturated solution of sal-ammoniac. The amalgam will immediately begin to swell, and will soon assume such an enormous volume as to float in the liquid as a metallic mass, which feels like butter. From the moment of its formation, however, this curious substance begins to change into ammonia, hydrogen and mercury. It is called ammonium amalgam, and it is believed to be formed by the temporary combination of the ammonium, which has lost its chlorine, with mercury. The following formula, in which n denotes an unknown number of atoms, explains its formation:

NH,C1+NaHg = NHHg + NaCl.

By washing with water the mercury is afterwards recovered without loss.

The specific gravity of ammonia gas is 8.5, a little more than half that of air (14.47). By a pressure seven times that of the atmosphere the gas can be condensed into a liquid, and if cooled to -75° C. (-103° F.) it freezes to a transparent, ice-like solid.

It must be borne in mind that the name ammonia is often given to the solution, which should be called ammonium hydrate.

Nitrous anhydride. HNO, Nitrous acid.

N2O4 or NO2 Nitric peroxide.

N2O

Nitric anhydride.

HNO, Nitric acid.

Of these compounds nitric acid is undoubtedly the most important, and as, moreover, it is the source from whence the other oxides are obtained, it claims our first attention.

The oxide N2O5, nitric anhydride, corresponding to nitric acid, is prepared with difficulty. It is a white, crystalline solid, which combines eagerly with water:

N2O + H2O = 2HNO3.

Nitric Acid or Aquafortis, HNO3.

Experiment 1.-Introduce into a small retort half an ounce of powdered saltpetre and half an ounce of common sulphuric acid, and let the retort stand erect for some time, in order

Fig. 59.

that as much as possible of the sulphuric acid remaining in the neck may flow down into the retort. Then surround the latter with sand contained in an iron basin, adapt to the beak of it a receiver, wrap round the joint some strips of moistened blottingpaper, and heat gently.

In a short time a yellowish fuming fluid passes over into the receiver, which is placed in a vessel filled with water, and must frequently be sprinkled with cold water; this fluid is heavier than water, and is called nitric acid.

Saltpetre is potassium nitrate, KNO,. When it is acted upon by sulphuric acid the metal and part of the hydrogen change places; we get nitric acid and a salt, called hydrogen potassium sulphate, HKSO, in which half the hydrogen of the acid is replaced by potassium :

KNO, H2SO, HNO, + HKSO..

=

At a higher temperature the hydrogen potassium sulphate will decompose another molecule of potassium nitrate:

KNO+HKSO, HNO3 + K2SO1;

=

but on the small scale it is better not to push the action so far or the retort will be apt to break. K2SO, is called neutral, or di-potassium sulphate.

By using perfectly dry potassium nitrate and concentrated sulphuric acid, true nitric acid, HNO,, can be obtained as a yellowish liquid, which fumes strongly when exposed to the air. It is hence called fuming nitric acid, and is the strongest that can be prepared. It is one and a half times heavier than water, its specific gravity being 1.52. A weaker kind is commonly met with in commerce as ordinary nitric acid, or aquafortis. It consists of 60 parts of true nitric acid and 40 of water, and has a specific gravity of 1-42. This acid is colourless when pure, but usually possesses a yellowish tint. When very strong or very weak nitric acid is heated, acid of this strength is obtained. It can be distilled unchanged.

Experiment 2.-A drop of nitric acid is sufficient to acidify several spoonfuls of water, and even at a greater dilution it will redden blue litmus-paper; nitric acid is accordingly distinctly characterised as an acid.

Experiment 3.—If lead be heated for a long time in the air it abstracts oxygen from it, and becomes converted into a reddish-yellow powder, called lead oxide, or litharge. Take up a small portion of this litharge on the point of a knife, put it into a test-tube, and add some dilute nitric acid. The greater part will be dissolved by gentle heating. Filter the solution while warm, and put it in a cold place; a salt will be deposited from it in white brilliant crystals; this is lead nitrate. This shows that lead oxide is a basic oxide, as it combines with acids forming salts.

Nitric acid dissolves most of the metallic oxides, and forms with them salts, all of which are soluble in water. For this reason, nitric acid is often used for cleaning metals, for instance, copper and brass instruments, which, during the process of annealing, soldering, &c., have become covered with a coating of oxide.

Experiment 4.-Pour over some shot, common nitric acid,

slightly diluted with water; a solution is also effected in this instance, but it is accompanied by the evolution of a yellowish-red vapour of a suffocating smell. This vapour is nitric peroxide. Part of the nitric acid is decomposed, while another part of it combines with the lead, and forms the same salt, as in the former experiment. This likewise crystallises from its solution, if it is evaporated until a film forms on its surface.

In this case the lead is apparently dissolved, but it is obvious that this is quite a different kind of solution from that of common salt or sugar in water. The salt and sugar are unchanged in the solution, while the lead is not contained in the liquid as a metal, but as a salt, a nitrate. The same

thing occurs with all other metals which are soluble in nitric acid; as, for example, with silver, mercury, copper, iron, &c. Gold is not dissolved by it; hence gold may be separated from silver by means of nitric acid.

Nitric acid readily parts with a portion of its oxygen. It is therefore a powerful oxidizing agent.

Experiment 5.-Place a few fragments of tin in a wineglass, and pour over them a little nitric acid. A violent action is set up, and red suffocating fumes are copiously evolved. The tin is not dissolved and converted into nitrate as the lead was, but becomes a white insoluble powder, an oxide of tin called meta-stannic acid.

Experiment 6.-Some of the non-metallic elements, as well as of the metals, are oxidized by nitric acid; charcoal, on being boiled in it, becomes carbonic anhydride; sulphur, sulphuric acid; phosphorus, phosphoric acid, &c. In all these cases yellowish-red fumes are evolved.

Experiment 7.-Organic substances also, as wool, feathers, wood, indigo, &c., are oxidized and decomposed by heating them with nitric acid. This sort of decomposition may be regarded as combustion in the moist way. If substances of animal origin are allowed to remain for a short time only in contact with this acid, they will assume a yellow colour. In this manner wood may be stained, and silk may be dyed yellow; the hands and clothes are also stained yellow by nitric acid. Cotton undergoes a most remarkable change if soaked for a short time in the strongest nitric acid; it will then detonate and explode, like gunpowder, only far more