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NITRIC ACID-MODES OF PREPARATION.

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The nitrate is introduced into the retort, A, through the opening at c, which is closed during the distillation by a stone lid, fitted

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accurately to the aperture; and the oil of vitriol is added by a funnel at e, after the retort is closed. As soon as the acid is introduced, the funnel is withdrawn, and the opening at e is closed with a plug. The nitric acid as it distils over passes through the pipe f, and is condensed in a series of stoneware bottles, the first of which is seen at B. The acid which is collected in the first receiver is always contaminated with sulphuric acid, and that in the last is rather dilute, as water is placed in it to condense the nitrous fumes.

Upon the large scale it is customary to employ a proportion of sulphuric acid smaller than that used when the distillation is performed in glass vessels, for it is quite possible to effect a complete decomposition of the nitrate by heating it with one-half its weight of oil of vitriol. Under these circumstances, however, a higher temperature is needed to expel the last portions of acid, and a considerable quantity of the nitric acid is thereby decomposed and wasted. The residue in the retort, when the smaller quantity of sulphuric acid is used, is much less soluble in water, and consequently is much more difficult of removal: but in the iron cylinder of the manufacturer this is of no moment, because the saline mass can easily be detached by the use of iron tools when the distillation is at an end.

The cause of these differences in the result of the processes adopted on the large and small scale lies in the fact that sulphuric acid by its reaction upon potash gives rise to two different sulphates, one of which contains twice as much potassium as the other; the acid sulphate consisting of KHSO,, while the neutral sulphate contains K,SO.

When nitre and sulphuric acid are mixed in the proportion of

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NITRIC ACID-PREPARATION-PROPERTIES.

equal weights, the acid sulphate of potassium is obtained, and nitric acid distils readily; the change is represented in the following equation :

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Acid sulphate of
potassium.

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but if the nitrate be mixed with sulphuric acid, in the proportion of 2 equivalents of each, the decomposition takes place in two successive stages; in the first of these, half the nitre only is decomposed, and a gentle heat only is needed for the distillation of the nitric acid so produced. The following equation may be employed to represent this part of the change :

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As soon as the first half of the nitric acid has passed over, the temperature begins to rise, and the acid sulphate of potassium reacts on the undecomposed nitre; the second half of the nitric acid is liberated, but at the same time is partially decomposed, particularly towards the end of the operation: the whole of the potassium remains in the retort in the form of the sparingly soluble neutral sulphate. This second stage of the decomposition is exhibited in the subjoined equation :—

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Nitric Acid (monohydrate), HN→ ̧; or HO,NO: Sp. gr. of liquid, 1517, at 69°; boiling-pt. 184°: Comp. in 100 parts, NO, 85'72; H2O, 14'28.-The acid which is obtained by the foregoing process is of a yellowish or red colour, owing to the presence of some of the lower oxides of nitrogen; these may, if necessary, be got rid of by mixing the acid with an equal bulk of oil of vitriol, and submitting the mixture to distillation. If the first portions be collected, and gently warmed while a current of dry air is sent through the acid, sheltered from strong daylight, the acid may be obtained as colourless as water, and quite free from the lower oxides of nitrogen. It is, however, so unstable in this concentrated form that it cannot be redistilled alone without experiencing partial decomposition. When exposed to the sun's light a similar effect is produced; oxygen gas is evolved, and the acid becomes coloured owing to the formation of lower oxides of nitrogen. When pure, nitric acid is a colourless, limpid, fuming,

NITRIC ACID-MODES OF DECOMPOSITION.

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powerfully corrosive liquid, which freezes at about -40°. It begins to boil at 184°, but the temperature rises gradually, owing to the decomposition of the liquid; oxygen and nitrous fumes are evolved; the boiling-point continues to rise slowly till it reaches 250°, at which point the acid in the retort has a composition approaching to (2HNO3, 3 H,→), and distils unchanged.

Owing to the facility with which the acid parts with oxygen, it is continually employed as an oxidizing agent. If it be dropped into hot finely powdered charcoal, the charcoal burns vividly; if it be mixed with a little oil of vitriol, and poured into oil of turpentine, the mixture bursts into flame. Sulphur, phosphorus, and iodine are oxidized by it, the phosphorus almost with explosive violence. Nitric acid is one of the most corrosive substances known. It rapidly destroys all animal textures, and if somewhat diluted stains the skin, wool, feathers, and all albuminous bodies of a bright yellow colour. The acid in a somewhat diluted state is, indeed, often used to impart a permanent yellow dye to woollen and silken goods. Nitric acid acts violently upon tin or iron filings, especially if they be previously moistened with a few drops of water; and indeed it attacks most of the metals except gold, platinum, rhodium, and iridium: but it is most active upon them when diluted to a specific gravity of from 1.35 to 1*25. The action of nitric acid upon the metals varies with its temperature and degree of dilution. pure concentrated acid, HNO,, is in fact without action upon tin, iron, bismuth, and many other metals at ordinary temperatures. The presence of nitrous acid in the nitric acid greatly increases its oxidizing power, for owing to the instability of nitrous acid this compound parts with its oxygen very readily. At a temperature of o° the acid, whether concentrated or dilute, is without action on copper, but it dissolves zinc rapidly.

The

(360) Action of Acids on Metals.-The chemical action of nitric acid upon the metals is a process of considerable importance, but in order to study with advantage the effects to which it gives rise it will be useful to consider the action of acids upon the metals from a general point of view. It has already been stated that the metals unite directly with many of the non-metallic elements, such as chlorine, oxygen, and sulphur. Antimony, for example, will take fire spontaneously if allowed to fall in fine powder into chlorine. Iron will burn in oxygen if first heated to the point of ignition; and copper turnings, if mixed with powdered sulphur, will, on the application of heat, combine with the sulphur, and will emit during the action a vivid light. The metallic oxides, when presented to the acids, become quickly dissolved; oxide of copper is brought

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NITRIC ACID-ACTION OF ACIDS UPON METALS.

into solution by diluted sulphuric acid, oxide of zinc quickly disappears when agitated with hydrochloric acid, and oxide of lead is rapidly dissolved by acetic acid.

But a metal will not unite directly with an anhydride. Union between a metallic oxide and an anhydride may, however, occur, though, even then, the action is much favoured by the presence of water. Sulphuric anhydride, for instance, does not act upon iron, but the anhydride is immediately absorbed by caustic potash, ᏦᎻᎾ + ᎦᎾ, becoming ᏦᎻᎦᎾ, ; and in like manner, carbonic anhydride is rapidly absorbed by slaked lime.

When the metals are presented to the acids other phenomena are observed; a brisk action frequently takes place, accompanied by the evolution of a gas, and it is very often stated that the metal first becomes oxidized, and is then dissolved by the acid.

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It is not, however, necessary to assume that the metal always undergoes oxidation as a preliminary step, for it may be supposed that the metal simply displaces the hydrogen of the acid. When, for example, zinc is placed in diluted sulphuric acid, the metal is dissolved with rapidity, whilst hydrogen escapes in the gaseous form; H2SO+ Zn yielding ZnSO4 + H2. A similar result is obtained when iron or tin is dissolved in hydrochloric acid, ferrous or stannous chloride being produced, whilst hydrogen is given off, the reaction in the case of iron being; 2 HCl + Fe = FeCl2, + H2. When an oxide is employed instead of a metal, the hydrogen, instead of escaping as gas, is eliminated in the form of water; for instance, in the action of oxide of zinc upon sulphuric acid, the change may be represented as Zn✪ + H2SO4 ZnᎦᎾ + H2; and again with oxide of copper and hydrochloric acid, Єu+2 HCl = EuCl2+ H2O.

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The ordinary action of metals upon sulphuric acid, in which the components of the acid are united by powerful chemical ties, is, as we have just seen, comparatively simple; but where the elements of the acid are but feebly held together, as is the case with nitric acid, the reactions are often more complicated. When, for example, silver or copper is dissolved in nitric acid, hydrogen may as before be displaced from the acid by the metal which becomes dissolved; but owing to the facility with which nitric acid parts with its oxygen no hydrogen is set free-for at the moment of its liberation it becomes oxidized at the expense of the elements of the nitric acid-one of the lower oxides of the nitrogen is formed, and occasions the disengagement of ruddy fumes instead of the colourless and inflammable hydrogen. In many instances it is probable that the radicle of the acid itself is deoxidized by the direct

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action of the metal, the oxide of the metal then forming a salt with the undecomposed portion of acid by double decomposition, as already explained where an oxide acts upon an acid-for example, when metallic silver acts upon heated nitric acid, a portion of the acid furnishes oxygen, disengaging nitric oxide, whilst the freshly formed oxide of silver reacts upon another portion of the acid, Ag2→ + 2 HNO, yielding H2O + 2 AgN→ ̧. The exact nature of the decomposition, however, varies in different cases; silver when allowed to become dissolved slowly in the cold in an excess of diluted nitric acid produces nitrous acid (HNO), which remains in solution; 2 Ag + 3 HNO, giving 2 AgNO3 + H ̧→ + HNO„, and the metal is dissolved without evolution of gas; a similar effect is also produced by palladium. With metals which attack the acid more vigorously, such as copper or mercury, nitric acid of moderate concentration (sp. gr. 125 to 13) disengages nitric oxide in large quantity: for example, 3 Єu + 8 HN→ yields 2 NO + 3 (Єu 2 NO3) + 4 H ̧Ð; but if the acid be more highly concentrated (sp. gr. 1'42), peroxide of nitrogen is disengaged abundantly; Eu + 4 HNO, yielding Єu 2 NO3 + 2 NO2 + 2H2O. And when the decomposition occurs at a high temperature, free nitrogen is usually disengaged in considerable quantity, the acid undergoing complete deoxidation. If the metal has a still more energetic action, as zinc, for example, the acid when dilute yields nitrous oxide amongst its gaseous products; 4 Zn + 10 HN→ N2O + 4 (Zn 2 NO3) + 5 H‚Ð. When zinc or tin is used with a stronger acid, ammonia is amongst the products; for instance, 4 Zn + 9 HNO2 = 4 (Zn 2 NO3) + 3H2+ H2N, the ammonia combining with the excess of acid employed.

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(361) Hydrates of Nitric Acid.-It is doubtful whether any definite hydrates of the acid HNO, really exist. When concentrated nitric acid is exposed to the air it absorbs moisture, and if 70 parts of the concentrated acid be mixed with 30 of water it emits a sensible amount of heat. Many chemists believe that a hydrate of nitric acid of considerable stability is formed under these circumstances, and that it has a composition represented by the formula (2 HNO3, 3 H2O.) This hydrate would contain 60 per cent. of the anhydride N→, and 40 of water: such an acid has a sp. gr. of 1424; it boils at 250°, and may be distilled under ordinary pressures apparently unaltered. A weaker acid when heated parts with its water till it arrives at this density, and a stronger acid, when distilled, also loses acid until reduced to this point, the liquid in the retort eventually, in both cases, acquiring a density of 1424. This apparent stability does not

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