Chapter VIII., Note 42.)

It is not only hydrocarbons which are subject to metalepsis. Certain other hydrogen compounds, under the action of chlorine, also give corresponding chlorine derivatives in exactly the same manner; for instance, ammonia, caustic potash, caustic lime, and a whole series of _alkaline_ substances.[27] In fact, just as the hydrogen in marsh gas can be replaced by chlorine and form methyl chloride, so the hydrogen in caustic potash, KHO, ammonia, NH_{3}, and calcium hydroxide, CaH_{2}O_{2} or Ca(OH)_{2}, may be replaced by chlorine and give potassium hypochlorite, KClO, calcium hypochlorite, CaCl_{2}O_{2}, and the so-called chloride of nitrogen, NCl_{3}. For not only is the correlation in composition the same as in the substitution in marsh gas, but the whole mechanism of the reaction is the same. Here also two atoms of chlorine act: one takes the place of the hydrogen whilst the other is evolved as hydrochloric acid, only in the former case the hydrochloric acid evolved remained free, and in the latter, in presence of alkaline substances, it reacts on them. Thus, in the action of chlorine on caustic potash, the hydrochloric acid formed acts on another quantity of caustic potash and gives potassium chloride and water, and therefore not only KHO + Cl_{2} = HCl + KClO, but also KHO + HCl = H_{2}O + KCl, and the result of both simultaneous phases will be 2KHO + Cl_{2} = H_{2}O + KCl + KClO. We will here discuss certain special cases.

[27] By including many instances of the action of chlorine under metalepsis we not only explain the indirect formation of CCl_{4}, NCl_{3}, and Cl_{2}O by one method, but we also arrive at the fact that the reactions of the metalepsis of the hydrocarbons lose that exclusiveness which was often ascribed to them. Also by subjecting the chemical representations to the law of substitution we may foretell metalepsis as a particular case of a general law.

The action of chlorine on ammonia may either result in the entire breaking up of the ammonia, with the evolution of gaseous nitrogen, or in a product of metalepsis (as with CH_{4}). With an excess of chlorine and the aid of heat the ammonia is decomposed, with the disengagement of free nitrogen.[28] This reaction evidently results in the formation of sal-ammoniac, 8NH_{3} + 3Cl_{2} = 6NH_{4}Cl + N_{2}. But if the ammonium salt be in excess, then the reaction takes the direction of the replacement of the hydrogen in the ammonia by chlorine. The principal result is that NH_{3} + 3Cl_{2} forms NCl_{3} + 3HCl.[29] The resulting product of metalepsis, or _chloride of nitrogen_, NCl_{3}, discovered by Dulong, is a liquid having the property of decomposing with excessive ease not only when heated, but even under the action of mechanical influences, as by a blow or by contact with certain solid substances. The explosion which accompanies the decomposition is due to the fact that the liquid chloride of nitrogen gives gaseous products, nitrogen and chlorine.[29 bis]

[28] This may be taken advantage of in the preparation of nitrogen. If a large excess of chlorine water be poured into a beaker, and a small quantity of a solution of ammonia be added, then, after shaking, nitrogen is evolved. If chlorine act on a dilute solution of ammonia, the volume of nitrogen does not correspond with the volume of the chlorine taken, because ammonium hypochlorite is formed. If ammonia gas be passed through a fine orifice into a vessel containing chlorine, the reaction of the formation of nitrogen is accompanied by the emission of light and the appearance of a cloud of sal-ammoniac. In all these instances an excess of chlorine must be present.

[29] The hydrochloric acid formed combines with ammonia, and therefore the final result is 4NH_{3} + 3Cl_{2} = NCl_{3} + 3NH_{4}Cl. For this reason, more ammonia must enter into the reaction, but the metalepsical reaction in reality only takes place with an excess of ammonia or its salt. If bubbles of chlorine be passed through a fine tube into a vessel containing ammonia gas, each bubble gives rise to an explosion. If, however, chlorine be passed into a solution of ammonia, the reaction at first brings about the formation of nitrogen, because chloride of nitrogen acts on ammonia like chlorine. But when sal-ammoniac has begun to form, then the reaction directs itself towards the formation of chloride of nitrogen. The first action of chlorine on a solution of sal-ammoniac always causes the formation of chloride of nitrogen, which then reacts on ammonia thus: NCl_{3} + 4NH_{3} = N_{2} + 3NH_{4}Cl. Therefore, so long as the liquid is alkaline from the presence of ammonia the chief product will be nitrogen. The reaction NH_{4}Cl + 3Cl_{2} = NCl_{3} + 4HCl is reversible; with a dilute solution it proceeds in the above-described direction (perhaps owing to the affinity of the hydrochloric acid for the excess of water), but with a strong solution of hydrochloric acid it takes the opposite direction (probably by virtue of the affinity of hydrochloric acid for ammonia). Therefore there must exist a very interesting case of equilibrium between ammonia, hydrochloric acid, chlorine, water, and chloride of nitrogen which has not yet been investigated. The reaction NCl_{3} + 4HCl = NH_{4}Cl + 3Cl_{2} enabled Deville and Hautefeuille to determine the composition of chloride of nitrogen. When slowly decomposed by water, chloride of nitrogen gives, like a chloranhydride, nitrous acid or its anhydride, 2NCl_{3} + 3H_{2}O = N_{2}O_{3} + 6HCl. From these observations it is evident that chloride of nitrogen presents great chemical interest, which is strengthened by its analogy with trichloride of phosphorus. The researches of F. F. Selivanoff (1891-94) prove that NCl_{3} may be regarded as an ammonium derivative of hypochlorous acid. Chloride of nitrogen is decomposed by dilute sulphuric acid in the following manner: NCl_{3} + 3H_{2}O + H_{2}SO_{4} = NH_{4}HSO_{4} + 3HClO. This reaction is reversible and is only complete when some substance, combining with HClO (for instance, succinimide) or decomposing it, is added to the liquid. This is easily understood from the fact that hypochlorous acid itself, HClO, may, according to the view held in this book, be regarded as the product of the metalepsis of water, and consequently bears the same relation to NCl_{3} as H_{2}O does to NH_{3}, or as RHO to RNH_{2}, R_{2}NH, and R_{3}N--that is to say, NCl_{3} corresponds as an ammonium derivative to ClOH and Cl_{2} in exactly the same manner as NR_{3} corresponds to ROH and R_{2}. The connection of NCl_{3} and other similar explosive chloro-nitrogen compounds (called chloryl compounds by Selivanoff; for example, the C_{2}H_{5}NCl_{2} of Wurtz is chloryl ethylamine), such as NRCl_{2} (as NC_{2}H_{5}Cl_{2}), and NR_{2}Cl (for instance, N(CH_{3}CO)HCl, chlorylacetamide, and N(C_{2}H_{5})_{2}Cl, chloryl diethylamine) with HClO is evident from the fact that under certain circumstances these compounds give hypochlorous acid, with water, for instance, NR_{2}Cl + H_{2}O = NR_{2}H + HClO, and frequently act (like NCl_{3} and HClO, or Cl_{2}) in an oxidising and chloridising manner. We may take chloryl succinimide, C_{2}H_{4}(CO)_{2}NCl for example. It was obtained by Bender by the action of HClO upon succinimide, C_{2}H_{4}(CO)_{2}NH, and is decomposed by water with the re-formation of amide and HClO (the reaction is reversible). Selivanoff obtained, investigated, and classified many of the compounds NR_{2}Cl and NRCl_{2}, where R is a residue of organic acids or alcohols, and showed their distinction from the chloranhydrides, and thus supplemented the history of chloride of nitrogen, which is the simplest of the amides containing chlorine, NR_{3}, where R is fully substituted by chlorine.

[29 bis] In preparing NCl_{3} every precaution must he used to guard against an explosion, and care should he taken that the NCl_{3} remains under a layer of water. Whenever an ammoniacal substance comes into contact with chlorine great care must be taken, because it may be a case of the formation of such products and a very dangerous explosion may ensue. The liquid product of the metalepsis of ammonia may be most safely prepared in the form of small drops by the action of a galvanic current on a slightly warm solution of sal-ammoniac; chlorine is then evolved at the positive pole, and this chlorine acting on the ammonia gradually forms the product of metalepsis which floats on the surface of the liquid (being carried up by the gas), and if a layer of turpentine be poured on to it these small drops, on coming into contact with the turpentine, give feeble explosions, which are in no way dangerous owing to the small mass of the substance formed. Drops of chloride of nitrogen may with great caution be collected for investigation in the following manner. The neck of a funnel is immersed in a basin containing mercury, and first a saturated solution of common salt is poured into the funnel, and above it a solution of sal-ammoniac in 9 parts of water. Chlorine is then slowly passed through the solutions, when drops of chloride of nitrogen fall into the salt water.

Chloride of nitrogen is a yellow oily liquid of sp. gr. 1·65, which boils at 71°, and breaks up into N + Cl_{3} at 97°. The contact of phosphorus, turpentine, india-rubber, &c. causes an explosion, which is sometimes so violent that a small drop will pierce through a thick board. The great ease with which chloride of nitrogen decomposes is dependent upon the fact that it is formed with an absorption of heat, which it evolves when decomposed, to the amount of about 38,000 heat units for NCl_{3}, as Deville and Hautefeuille determined.

Chlorine, when absorbed by a solution of caustic soda (and also of other alkalis) at the ordinary temperature, causes the replacement of the hydrogen in the caustic soda by the chlorine, with the formation of sodium chloride by the hydrochloric acid, so that the reaction may be represented in two phases, as described above. In this manner, sodium hypochlorite, NaClO, and sodium chloride are simultaneously formed: 2NaHO + Cl_{2} = NaCl + NaClO + H_{2}O. The resultant solution contains NaClO and is termed 'eau de Javelle.' An exactly similar reaction takes place when chlorine is passed over dry hydrate of lime at the ordinary temperature: 2Ca(HO)_{2} + 2Cl_{2} = CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. A mixture of the product of metalepsis with calcium chloride is obtained. This mixture is employed in practice on a large scale, and is termed 'bleaching powder,' owing to its acting, especially when mixed with acids, as a bleaching agent on tissues, so that it resembles chlorine in this respect. It is however preferable to chlorine, because the destructive action of the chlorine can be moderated in this case, and because it is much more convenient to deal with a solid substance than with gaseous chlorine. Bleaching powder is also called _chloride of lime_, because it is obtained from chlorine and hydrate of lime, and contains[30] both these substances. It may be prepared in the laboratory by passing a current of chlorine through a cold mixture of water and lime (milk of lime). The mixture must be kept cold, as otherwise 3Ca(ClO)_{2} passes into 2CaCl_{2} + Ca(ClO_{3})_{2}. In the manufacture of bleaching powder in large quantities at chemical works, the purest possible slaked lime is taken and laid in a thin layer in large flat chambers, M (whose walls are made of Yorkshire flags or tarred wood, on which chlorine has no action), and into which chlorine gas is introduced by lead tubes. The distribution of the plant is shown in the annexed drawing (fig. 67).

[30] Quicklime, CaO (or calcium carbonate, CaCO_{3}), does not absorb chlorine when cold, but at a red heat, in a current of chlorine, it forms calcium chloride, with the evolution of oxygen. (This was confirmed in 1893 by Wells, at Oxford.) This reaction corresponds with the decomposing action of chlorine on methane, ammonia, and water. Slaked lime (calcium hydroxide, CaH_{2}O_{2}) also, when dry, does not absorb chlorine at 100°. The absorption proceeds at the ordinary temperature (below 40°). The dry mass thus obtained contains not less than three equivalents of calcium hydroxide to four equivalents of chlorine, so that its composition is [Ca(HO)_{2}]_{5}Cl_{4}. In all probability a simple absorption of chlorine by the lime at first takes place in this case, as may be seen from the fact that even carbonic anhydride, when acting on the dry mass obtained as above, disengages all the chlorine from it, leaving only calcium carbonate. But if the bleaching powder be obtained by a wet method, or if it be dissolved in water (in which it is very soluble), and carbonic anhydride be passed into it, then chlorine is no longer disengaged, but chlorine oxide, Cl_{2}O, and only half of the chlorine is converted into this oxide, while the other half remains in the liquid as calcium chloride. From this it may be inferred that calcium chloride is formed by the action of water on bleaching powder, and this is proved to be the case by the fact that small quantities of water extract a considerable amount of calcium chloride from bleaching powder. If a large quantity of water act on bleaching powder an excess of calcium hydroxide remains, a portion of which is not subjected to change. The action of the water may be expressed by the following formulæ: From the dry mass Ca_{3}(HO)_{6}Cl_{4} there is formed lime, Ca(HO)_{2}, calcium chloride, CaCl_{2}, and a saline substance, Ca(ClO)_{2}. Ca_{3}H_{6}O_{6}Cl_{4} = CaH_{2}O_{2} + CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. The resulting substances are not equally soluble; water first extracts the calcium chloride, which is the most soluble, then the compound Ca(ClO)_{2} and ultimately calcium hydroxide is left. A mixture of calcium chloride and hypochlorite passes into solution. On evaporation there remains Ca_{2}O_{2}Cl_{4}3H_{2}O. The dry bleaching powder does not absorb more chlorine, but the solution is able to absorb it in considerable quantity. If the liquid be boiled, a considerable amount of chlorine monoxide is evolved. After this calcium chloride alone remains in solution, and the decomposition may be expressed as follows: CaCl_{2} + CaCl_{2}O_{2} + 2Cl_{2} = 2CaCl_{2} + 2Cl_{2}O. Chlorine monoxide may be prepared in this manner.

It is sometimes said that bleaching powder contains a substance, Ca(OH)_{2}Cl_{2}, that is calcium peroxide, CaO_{2}, in which one atom of oxygen is replaced by (OH)_{2}, and the other by Cl_{2}; but, judging from what has been said above, this can only be the case in the dry state, and not in solutions.

On being kept for some time, bleaching powder sometimes decomposes, with the evolution of oxygen (because CaCl_{2}O_{2} = CaCl_{2} + O_{2}, _see_ p. 163); the same takes place when it is heated.

The products of the metalepsis of alkaline hydrates, NaClO and Ca(ClO)_{2}, which are present in solutions of 'Javelle salt' and bleaching powder (they are not obtained free from metallic chlorides), must be counted as salts, because their metals are capable of substitution. But the hydrate HClO corresponding with these salts, or _hypochlorous acid_, is not obtained in a free or pure state, for two reasons: in the first place, because this hydrate, as a very feeble acid, splits up (like H_{2}CO_{3} or HNO_{3}) into water and the anhydride, or _chlorine monoxide_, Cl_{2}O = 2HClO-H_{2}O; and, in the second place, because, in a number of instances, it evolves oxygen with great facility, forming hydrochloric acid: HClO = HCl + O. Both hypochlorous acid and chlorine monoxide may be regarded as products of the metalepsis of water, because HOH corresponds with ClOH and ClOCl. Hence in many instances bleaching salts (a mixture of hypochlorites and chlorides) break up, with the evolution of (1) _chlorine_, under the action of an excess of a powerful acid capable of evolving hydrochloric acid from sodium or calcium chlorides, and this takes place most simply under the action of hydrochloric acid itself, because (p. 462) NaCl + NaClO + 3HCl = 2NaCl + HCl + Cl_{2} + H_{2}O; (2) _oxygen_, as we saw in Chapter III.--The bleaching properties and, in general, _oxidising action_ of bleaching salts is based on this evolution of oxygen (or chlorine); oxygen is also disengaged on heating the dry salts--for instance, NaCl + NaClO = 2NaCl + O; (3) and, lastly, _chlorine monoxide_, which contains both chlorine and oxygen. Thus, if a little sulphuric, nitric, or similar acid (not enough to liberate hydrochloric acid from the CaCl_{2}) be added to a solution of a bleaching salt (which has an alkaline reaction, owing either to an excess of alkali or to the feeble acid properties of HClO), then the hypochlorous acid set free gives water and chlorine monoxide. If carbonic anhydride (or boracic or a similar very feeble acid) act on the solution of a bleaching salt, then hydrochloric acid is not evolved from the sodium or calcium chlorides, but the hypochlorous acid is displaced and gives chlorine monoxide,[31] because hypochlorous acid is one of the most feeble acids. Another method for the preparation of chlorine monoxide is based on these feeble acid properties of hypochlorous acid. Zinc oxide and mercury oxide, under the action of chlorine in the presence of water, do not give a salt of hypochlorous acid, but form a chloride and hypochlorous acid, which fact shows the incapacity of this acid to combine with the bases mentioned. Therefore, if such oxides as those of zinc or mercury be shaken up in water, and chlorine be passed through the turbid liquid,[32] a reaction occurs which may be expressed in the following manner: 2HgO + 2Cl_{2} = Hg_{2}OCl_{2} + Cl_{2}O. In this case, a compound of mercury oxide with mercury chloride, or the so-called mercury oxychloride, is obtained: Hg_{2}OCl_{2} = HgO + HgCl_{2}. This is insoluble in water, and is not affected by hypochlorous anhydride, so that the solution will contain hypochlorous acid only, but the greater part of it splits up into the anhydride and water.[32 bis]

[31] For this reason it is necessary that in the preparation of bleaching powder the chlorine should be free from hydrochloric acid, and even the lime from calcium chloride. An excess of chlorine, in acting on a solution of bleaching powder, may also give chlorine monoxide, because calcium carbonate also gives chlorine monoxide under the action of chlorine. This reaction may be brought about by treating freshly precipitated calcium carbonate with a stream of chlorine in water: 2Cl_{2} + CaCO_{3} = CO_{2} + CaCl_{2} + Cl_{2}O. From this we may conclude that, although carbonic anhydride displaces hypochlorous anhydride, it may be itself displaced by an excess of the latter.

[32] Dry red mercury oxide acts on chlorine, forming dry hypochlorous anhydride (chlorine monoxide) (Balard); when mixed with water, red mercury oxide acts feebly on chlorine, and when freshly precipitated it evolves oxygen and chlorine. An oxide of mercury which easily and abundantly evolves chlorine monoxide under the action of chlorine in the presence of water may be prepared as follows: the oxide of mercury, precipitated from a mercuric salt by an alkali, is heated to 300° and cooled (Pelouze). If a salt, MClO, be added to a solution of mercuric salt, HgX_{2}, mercuric oxide is liberated, because the hypochlorite is decomposed.

[32 bis] A solution of hypochlorous anhydride is also obtained by the action of chlorine on many salts; for example, in the action of chlorine on a solution of sodium sulphate the following reaction takes place: Na_{2}SO_{4} + H_{2}O + Cl_{2} = NaCl + HClO + NaHSO_{4}. Here the hypochlorous acid is formed, together with HCl, at the expense of chlorine and water, for Cl_{2} + H_{2}O = HCl + HClO. If the crystallo-hydrate of chlorine be mixed with mercury oxide, the hydrochloric acid formed in the reaction gives mercury chloride, and hypochlorous acid remains in solution. A dilute solution of hypochlorous acid or chlorine monoxide may be concentrated by distillation, and if a substance which takes up water (without destroying the acid)--for instance, calcium nitrate--be added to the stronger solution, then the anhydride of hypochlorous acid--_i.e._ chlorine monoxide--is disengaged.

Chlorine monoxide, which corresponds to bleaching and hypochlorous salts, containing as it does the two elements oxygen and chlorine, forms a characteristic example of a compound of elements which, in the majority of cases, act chemically in an analogous manner. Chlorine monoxide, as prepared from an aqueous solution by the abstraction of water or by the action of dry chlorine on cold mercury oxide, is, at the ordinary temperature, a gas or vapour which condenses into a red liquid boiling at +20° and giving a vapour whose density (43 referred to hydrogen) shows that 2 vols. of chlorine and 1 vol. of oxygen give 2 vols. of chlorine monoxide. In an anhydrous form the gas or liquid easily explodes, splitting up into chlorine and oxygen. This explosiveness is determined by the fact that heat is _evolved_ in the decomposition to the amount of about 15,000 heat units for Cl_{2}O.[33] The explosion may even take place spontaneously, and also in the presence of many oxidisable substances (for instance, sulphur, organic compounds, &c.), but the solution, although unstable and showing a strong oxidising tendency, does not explode.[34] It is evident that the presence of hypochlorous acid, HClO, may be assumed in an aqueous solution of Cl_{2}O, since Cl_{2}O + H_{2}O = 2HClO.

[33] All explosive substances are of this kind--ozone, hydrogen peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they cannot be formed directly from the elements or their simplest compounds, but, on the contrary, decompose into them. In a liquid state chlorine monoxide explodes even on contact with powdery substances, or when rapidly agitated--for instance, if a file be rasped over the vessel in which it is contained.

[34] A solution of chlorine monoxide, or hypochlorous acid, does not explode, owing to the presence of the mass of water. In dissolving, chlorine monoxide evolves about 9,000 heat units, so that its store of heat becomes less.

The capacity of hypochlorous acid (studied by Carius and others) for entering into combination with the unsaturated hydrocarbons is very often taken advantage of in organic chemistry. Thus its solution absorbs ethylene, forming the chlorhydrin C_{2}H_{4}ClOH.

The oxidising action of hypochlorous acid and its salts is not only applied to bleaching but also to many reactions of oxidation. Thus it converts the lower oxides of manganese into the peroxide.

Hypochlorous acid, its salts, and chlorine monoxide serve as a transition between hydrochloric acid, chlorides, and chlorine, and a whole series of compounds containing the same elements combined with a still greater quantity of oxygen. The higher oxides of chlorine, as their origin indicates, are closely connected with hypochlorous acid and its salts:

Cl_{2}, NaCl, HCl, hydrochloric acid. Cl_{2}O, NaClO, HClO, hypochlorous acid. Cl_{2}O_{3}, NaClO_{2}, HClO_{2}, chlorous acid.[35] Cl_{2}O_{5}, NaClO_{3}, HClO_{3}, chloric acid. Cl_{2}O_{7}, NaClO_{4}, HClO_{4}, perchloric acid.

When heated, solutions of hypochlorites undergo a remarkable change. Themselves so unstable, they, without any further addition, yield two fresh salts which are both much more stable; one contains more oxygen than MClO, the other contains none at all.

3MClO = MClO_{3} + 2MCl hypochlorite chlorate chloride

[35] _Chlorous acid_, HClO_{2} (according to the data given by Millon, Brandau, and others) in many respects resembles hypochlorous acid, HClO, whilst they both differ from chloric and perchloric acids in their degree of stability, which is expressed, for instance, in their bleaching properties; the two higher acids do not bleach, but both the lower ones do so (oxidise at the ordinary temperature). On the other hand, chlorous acid is analogous to nitrous acid, HNO_{2}. The anhydride of chlorous acid, Cl_{2}O_{3}, is not known in a pure state, but it probably occurs in admixture with chlorine dioxide, ClO_{2}, which is obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate with such reducing substances as nitric oxide, arsenious oxide, sugar, &c. All that is at present known is that pure chlorine dioxide ClO_{2} (_see_ Notes 39-43) is gradually converted into a mixture of hypochlorous and chlorous acids under the action of water (and alkalis); that is, it acts like nitric peroxide, NO_{2} (giving HNO_{3} and HNO_{2}), or as a mixed anhydride, 2ClO_{2} + H_{2}O = HClO_{3} + HClO_{2}. The silver salt, AgClO_{2}, is sparingly soluble in water. The investigations of Garzarolli-Thurnlackh and others seem to show that the anhydride Cl_{2}O_{3} does not exist in a free state.

Part of the salt--namely, two-thirds of it--parts with its oxygen in order to oxidise the remaining third.[36] From an intermediate substance, RX, two extremes, R and RX_{3} are formed, just as nitrous anhydride splits up into nitric oxide and nitric anhydride (or nitric acid). The resulting salt, MClO_{3}, corresponds with _chloric acid_ and potassium chlorate, KClO_{3}. It is evident that a similar salt may be obtained directly by the action of chlorine on an alkali if its solution be heated, because RClO will be first formed, and then RClO_{3}; for example, 6KHO + 3Cl_{2} = KClO_{3} + 5KCl + 3H_{2}O. Chlorates are so prepared; for instance, _potassium chlorate_, which is easily separated from potassium chloride, being sparingly soluble in cold water.[37]

[36] Hydrochloric acid, which is an example of compounds of this kind, is a saturated substance which does not combine directly with oxygen, but in which, nevertheless, a considerable quantity of oxygen may be inserted between the elements forming it. The same may be observed in a number of other cases. Thus oxygen may be added or inserted between the elements, sometimes in considerable quantities, in the saturated hydrocarbons; for instance, in C_{3}H_{8}, three atoms of oxygen produce an alcohol, glycerin or glycerol, C_{3}H_{5}(OH)_{3}. We shall meet with similar examples hereafter. This is generally explained by regarding oxygen as a bivalent element--that is, as capable of combining with two different elements, such as chlorine, hydrogen, &c. On the basis of this view, it may be inserted between each pair of combined elements; the oxygen will then be combined with one of the elements by one of its affinities and with the other element by its other affinity. This view does not, however, express the entire truth of the matter, even when applied to the compounds of chlorine. Hypochlorous acid, HOCl--that is, hydrochloric acid in which one atom of oxygen is inserted--is, as we have already seen, a substance of small stability; it might therefore be expected that on the addition of a fresh quantity of oxygen, a still less stable substance would be obtained, because, according to the above view, the chlorine and hydrogen, which form such a stable compound together, are then still further removed from each other. But it appears that chloric and perchloric acid, HClO_{3} and HClO_{4}, are much more stable substances. Furthermore, the addition of oxygen has also its limit, it can only be added to a certain extent. If the above representation were true and not merely hypothetical, there would be no limit to the combination of oxygen, and the more it entered into one continuous chain the more unstable would be the resultant compound. But not more than four atoms of oxygen can be added to hydrogen sulphide, nor to hydrochloric acid, nor to hydrogen phosphide. This peculiarity must lie in the properties of oxygen itself; four atoms of oxygen seem to have the power of forming a kind of radicle which retains two or several atoms of various other substances--for example, chlorine and hydrogen, hydrogen and sulphur, sodium and manganese, phosphorus and metals, &c., forming comparatively stable compounds, NaClO_{4}, Na_{2}SO_{4}, NaMnO_{4}, Na_{3}PO_{4}, &c. _See_ Chapter X. Note 1 and Chapter XV.

[37] If chlorine be passed through a _cold_ solution of potash, a bleaching compound, potassium chloride and hypochlorite, KCl + KClO, is formed, but if it be passed through a _hot_ solution potassium chlorate is formed. As this is sparingly soluble in water, it chokes the gas-conducting tube, which should therefore be widened out at the end.

Potassium chlorate is usually obtained on a large scale from calcium chlorate, which is prepared by passing chlorine (as long as it is absorbed) into water containing lime, the mixture being kept warm. A mixture of calcium chlorate and chloride is thus formed in the solution. Potassium chloride is then added to the warm solution, and on cooling a precipitate of potassium chlorate is formed as a substance which is sparingly soluble in cold water, especially in the presence of other salts. The double decomposition taking place is Ca(ClO_{3})_{2} + 2KCl = CaCl_{2} + 2KClO_{3}. On a small scale in the laboratory potassium chlorate is best prepared from a strong solution of bleaching powder by passing chlorine through it and then adding potassium chloride. KClO_{3} is always formed by the action of an electric current on a solution of KCl, especially at 80° (Häussermann and Naschold, 1894), so that this method is now used on a large scale.

Potassium chlorate crystallises easily in large colourless tabular crystals. Its solubility in 100 parts of water at 0° = 3 parts, 20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For comparison we will cite the following figures showing the solubility of potassium chloride and perchlorate in 100 parts of water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1 part, 20° about 1-3/4 part, 100° about 18 parts. When heated, potassium chlorate melts (the melting point has been given as from 335°-376°; according to the latest determination by Carnelley, 359°) and decomposes with the evolution of oxygen, potassium perchlorate being at first formed, as will afterwards be described (_see_ Note 47). A mixture of potassium chlorate and nitric and hydrochloric acids effects oxidation and chlorination in solutions. It deflagrates when thrown upon incandescent carbon, and when mixed with sulphur (1/3 by weight) it ignites it on being struck, in which case an explosion takes place. The same occurs with many metallic sulphides and organic substances. Such mixtures are also ignited by a drop of sulphuric acid. All these effects are due to the large amount of oxygen contained in potassium chlorate, and to the ease with which it is evolved. A mixture of two parts of potassium chlorate, one part of sugar, and one part of yellow prussiate of potash acts like gunpowder, but burns too rapidly, and therefore bursts the guns, and it also has a very strong oxidising action on their metal. The sodium salt, NaClO_{3}, is much more soluble than the potassium salt, and it is therefore more difficult to free it from sodium chloride, &c. The barium salt is also more soluble than the potassium salt; O° = 24 parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.

If dilute sulphuric acid be added to a solution of potassium chlorate, _chloric acid_ is liberated, but it cannot be separated by distillation, as it is decomposed in the process. To obtain the free acid, sulphuric acid must be added to a solution of barium chlorate.[38] The sulphuric acid gives a precipitate of barium sulphate, and free chloric acid remains in solution. The solution may be evaporated under the receiver of an air-pump. This solution is colourless, has no smell, and acts as a powerful acid (it neutralises sodium hydroxide, decomposes sodium carbonate, gives hydrogen with zinc, &c.); when heated above 40°, however, it decomposes, forming chlorine, oxygen, and perchloric acid: 4HClO_{3} = 2HClO_{4} + H_{2}O + Cl_{2} + O_{3}. In a concentrated condition the acid acts as an exceedingly energetic oxidiser, so that organic substances brought into contact with it burst into flame. Iodine, sulphurous acid, and similar oxidisable substances form higher oxidation products and reduce the chloric acid to hydrochloric acid. Hydrochloric acid gas gives chlorine with chloric acid (and consequently with KClO_{3} also) acting in the same manner as it acts on the lower acids: HClO_{3} + 5HCl = 3H_{2}O + 3Cl_{2}.

[38] Barium chlorate, Ba(ClO_{3})_{2},H_{2}O, is prepared in the following way: impure chloric acid is first prepared and saturated with baryta, and the barium salt purified by crystallisation. The impure free chloric acid is obtained by converting the potassium in potassium chlorate into an insoluble salt. This is done by adding tartaric or hydrofluosilicic acid to a solution of potassium chlorate, because potassium tartrate and potassium silicofluoride are very sparingly soluble in water. Chloric acid is easily soluble in water.

By cautiously acting on potassium chlorate with sulphuric acid, the _dioxide_ (_chloric peroxide_), ClO_{2},[39] is obtained (Davy, Millon). This gas is easily liquefied in a freezing mixture, and boils at +10°. The vapour density (about 35 if H = 1) shows that the molecule of this substance is ClO_{2}.[40] In a gaseous or liquid state it very easily explodes (for instance, at 60°, or by contact with organic compounds or finely divided substances, &c.), forming Cl and O_{2}, and in many instances[41] therefore it acts as an oxidising agent, although (like nitric peroxide) it may itself be further oxidised.[42] In dissolving in water or alkalis chloric peroxide gives chlorous and hypochlorous acids--2ClO_{2} + 2KHO = KClO_{3} + KClO_{2} + H_{2}O--and therefore, like nitric peroxide, the dioxide may be regarded as an intermediate oxide between the (unknown) anhydrides of chlorous and chloric acids: 4ClO_{2} = Cl_{2}O_{3} + Cl_{2}O_{3}.[43]

[39] To prepare ClO_{2} 100 grams of sulphuric acid are cooled in a mixture of ice and salt, and 15 grams of powdered potassium chlorate are gradually added to the acid, which is then carefully distilled at 20° to 40°, the vapour given off being condensed in a freezing mixture. Potassium perchlorate is then formed: 3KClO_{3} + 2H_{2}SO_{4} = 2KHSO_{4} + KClO_{4} + 2ClO_{2} + H_{2}O. The reaction may result in an explosion. Calvert and Davies obtained chloric peroxide without the least danger by heating a mixture of oxalic acid and potassium chlorate in a test tube in a water-bath. In this case 2KClO_{3} + 3C_{2}H_{2}O_{4},2H_{2}O = 2C_{2}HKO_{4} + 2CO_{2} + 2ClO_{2} + 8H_{2}O. The reaction is still further facilitated by the addition of a small quantity of sulphuric acid. If a solution of HCl acts upon KClO_{3} at the ordinary temperature, a mixture of Cl_{2} and ClO_{2} is formed, but if the temperature be raised to 80° the greater part of the ClO_{2} decomposes, and when passed through a hot solution of MnCl_{2} it oxidises it. Gooch and Kreider proposed (1894) to employ this method for preparing small quantities of chlorine in the laboratory.

[40] By analogy with nitric peroxide it might be expected that at low temperatures a doubling of the molecule into Cl_{2}O_{4} would take place, as the reactions of ClO_{2} point to its being a mixed anhydride of HClO_{2} and HClO_{3}.

[41] Owing to the formation of this chlorine dioxide, a mixture of potassium chlorate and sugar is ignited by a drop of sulphuric acid. This property was formerly made use of for making matches, and is now sometimes employed for setting fire to explosive charges by means of an arrangement in which the acid is caused to fall on the mixture at the moment required. An interesting experiment on the combustion of phosphorus under water may be conducted with chlorine dioxide. Pieces of phosphorus and of potassium chlorate are placed under water, and sulphuric acid is poured on to them (through a long funnel); the phosphorus then burns at the expense of the chlorine dioxide.

[42] Potassium permanganate oxidises chlorine dioxide into chloric acid (Fürst).

[43] The euchlorine obtained by Davy by gently heating potassium chlorate with hydrochloric acid is (Pebal) a mixture of chlorine dioxide and free chlorine. The liquid and gaseous chlorine oxide (Note 35), which Millon considered to be Cl_{2}O_{3}, probably contains a mixture of ClO_{2} (vapour density 35), Cl_{2}O_{3} (whose vapour density should be 59), and chlorine (vapour density 35·5), since its vapour density was determined to be about 40.

As the salts of chloric acid, HClO_{3}, are produced by the splitting up of the salts of hypochlorous acid, so in the same way the salts of perchloric acid, HClO_{4}, are produced from the salts of chloric acid, HClO_{3}. But this is the highest form of the oxidation of HCl. _Perchloric acid_, HClO_{4}, is the most stable of all the acids of chlorine. When fused potassium chlorate begins to swell up and solidify, after having parted with one-third of its oxygen, potassium chloride and potassium perchlorate have been formed according to the equation 2KClO_{3} = KClO_{4} + KCl + O_{2}.

The formation of this salt is easily observed in the preparation of oxygen from potassium chlorate, owing to the fact that the potassium perchlorate fuses with greater difficulty than the chlorate, and therefore appears in the molten salt as solid grains (_see_ Chapter III. Note 12). Under the action of certain acids--for instance, sulphuric and nitric--potassium chlorate also gives potassium perchlorate. This latter may be easily purified, because it is but sparingly soluble in water, although all the other salts of perchloric acid are very soluble and even deliquesce in the air. The perchlorates, although they contain more oxygen than the chlorates, are decomposed with greater difficulty, and even when thrown on ignited charcoal give a much feebler deflagration than the chlorates. Sulphuric acid (at a temperature not below 100°) evolves volatile and to a certain extent stable perchloric acid from potassium perchlorate. Neither sulphuric nor any other acid will further decompose perchloric acid as it decomposes chloric acid. Of all the acids of chlorine, perchloric acid alone can be distilled.[44] The pure hydrate HClO_{4}[45] is a colourless and exceedingly caustic substance which fumes in the air and has a specific gravity 1·78 at 15° (sometimes, after being kept for some time, it decomposes with a violent explosion). It explodes violently when brought into contact with charcoal, paper, wood, and other organic substances. If a small quantity of water be added to this hydrate, and it be cooled, a crystallo-hydrate, ClHO_{4},H_{2}O, separates out. This is much more stable, but the liquid hydrate HClO_{4},2H_{2}O is still more so. The acid dissolves in water in all proportions, and its solutions are distinguished for their stability.[46] When ignited both the acid and its salts are decomposed, with the evolution of oxygen.[47]

[44] If a solution of chloric acid, HClO_{3}, be first concentrated over sulphuric acid under the receiver of an air-pump and afterwards distilled, chlorine and oxygen are evolved and perchloric acid is formed: 4HClO_{3} = 2HClO_{4} + Cl_{2} + 3O + H_{2}O. Roscoe accordingly decomposed directly a solution of potassium chlorate by hydrofluosilicic acid, decanted it from the precipitate of potassium silicofluoride, K_{2}SiF_{6}, concentrated the solution of chloric acid, and then distilled it, perchloric acid being then obtained (_see_ following footnote). That chloric acid is capable of passing into perchloric acid is also seen from the fact that potassium permanganate is decolorised, although slowly, by the action of a solution of chloric acid. On decomposing a solution of potassium chlorate by the action of an electric current, potassium perchlorate is obtained at the positive electrode (where the oxygen is evolved). Perchloric acid is also formed by the action of an electric current on solutions of chlorine and chlorine monoxide. Perchloric acid was obtained by Count Stadion and afterwards by Serullas, and was studied by Roscoe and others.

[45] Perchloric acid, which is obtained in a free state by the action of sulphuric acid on its salts, may be separated from a solution very easily by distillation, being volatile, although it is partially decomposed by distillation. The solution obtained after distillation may be concentrated by evaporation in open vessels. In the distillation the solution reaches a temperature of 200°, and then a very constant liquid hydrate of the composition HClO_{4},2H_{2}O is obtained in the distillate. If this hydrate be mixed with sulphuric acid, it begins to decompose at 100°, but nevertheless a portion of the acid passes over into the receiver without decomposing, forming a crystalline hydrate HClO_{4},H_{2}O which melts at 50°. On carefully heating this hydrate it breaks up into perchloric acid, which distills over below 100°, and into the liquid hydrate HClO_{4},2H_{2}O. The acid HClO_{4} may also be obtained by adding one-fourth part of strong sulphuric acid to potassium chlorate, carefully distilling and subjecting the crystals of the hydrate HClO_{4},H_{2}O obtained in the distillate to a fresh distillation. Perchloric acid, HClO_{4}, itself does not distil, and is decomposed on distillation until the more stable hydrate HClO_{4}, H_{2}O is formed; this decomposes into HClO_{4} and HClO_{4},2H_{2}O, which latter hydrate distils without decomposition. This forms an excellent example of the influence of water on stability, and of the property of chlorine of giving compounds of the type ClX_{7}, of which all the above hydrates, ClO_{3}(OH), ClO_{2}(OH)_{3}, and ClO(OH)_{5}, are members. Probably further research will lead to the discovery of a hydrate Cl(OH)_{7}.

[46] According to Roscoe the specific gravity of perchloric acid = 1·782 and of the hydrate HClO_{4},H_{2}O in a liquid state (50°) 1·811; hence a considerable contraction takes place in the combination of HClO_{4} with H_{2}O.

[47] The decomposition of salts analogous to potassium chlorate has been more fully studied in recent years by Potilitzin and P. Frankland. Professor Potilitzin, by decomposing, for example, lithium chlorate LiClO_{3}, found (from the quantity of lithium chloride and oxygen) that at first the decomposition of the fused salt (368°) takes place according to the equation, 3LiClO_{3} = 2LiCl + LiClO_{4} + 5O, and that towards the end the remaining salt is decomposed thus: 5LiClO_{3} = 4LiCl + LiClO_{4} + 10O. The phenomena observed by Potilitzin obliged him to admit that lithium perchlorate is capable of decomposing simultaneously with lithium chlorate, with the formation of the latter salt and oxygen; and this was confirmed by direct experiment, which showed that lithium chlorate is always formed in the decomposition of the perchlorate. Potilitzin drew particular attention to the fact that the decomposition of potassium chlorate and of salts analogous to it, although exothermal (Chapter III., Note 12), not only does not proceed spontaneously, but requires time and a rise of temperature in order to attain completion, which again shows that chemical equilibria are not determined by the heat effects of reactions only.

P. Frankland and J. Dingwall (1887) showed that at 448° (in the vapour of sulphur) a mixture of potassium chlorate and powdered glass is decomposed almost in accordance with the equation 2KClO_{3} = KClO_{4} + KCl + O_{2}, whilst the salt by itself evolves about half as much oxygen, in accordance with the equation, 8KClO_{3} = 5KClO_{4} + 3KCl + 2O_{2}. The decomposition of potassium perchlorate in admixture with manganese peroxide proceeds to completion, KClO_{4} = KCl + 2O_{2}. But in decomposing by itself the salt at first gives potassium chlorate, approximately according to the equation 7KClO_{4} = 2KClO_{3} + 5KCl + 11O_{2}. Thus there is now no doubt that when potassium chlorate is heated, the perchlorate is formed, and that this salt, in decomposing with evolution of oxygen, again gives the former salt.

In the decomposition of barium hypochlorite, 50 per cent. of the whole amount passes into chlorate, in the decomposition of strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of calcium hypochlorite about 2·5 per cent. Besides which Potilitzin showed that the decomposition of the hypochlorites and also of the chlorates is always accompanied by the formation of a certain quantity of the oxides and by the evolution of chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Prost (1889) represent the evolution of oxygen from KClO_{3} as due to the salt first splitting up into base and anhydride, thus (1) 2MClO_{3} = M_{2}O + Cl_{2}O_{5}; (2) Cl_{2}O_{5} = Cl_{2} + O_{3}; and (3) M_{2}O + Cl = 2MCl + O.

I may further remark that the decomposition of potassium chlorate as a reaction evolving heat easily lends itself for this very reason to the contact action of manganese peroxide and other similar admixtures; for such very feeble influences as those of contact may become evident either in those cases (for instance, detonating gas, hydrogen peroxide, &c.), when the reaction is accompanied by the evolution of heat, or when (for instance, H_{2} + I_{2}, &c.) little heat is absorbed or evolved. In these cases it is evident that the existing equilibrium is not very stable, and that a small alteration in the conditions at the surfaces of contact may suffice to upset it. In order to conceive the _modus operandi_ of contact phenomena, it is enough to imagine, for instance, that at the surface of contact the movement of the atoms in the molecules changes from a circular to an elliptical path. Momentary and transitory compounds may he formed, but their formation cannot affect the explanation of the phenomena.

On comparing chlorine as an element not only with nitrogen and carbon but with all the other non-metallic elements (chlorine has so little analogy with the metals that a comparison with them would be superfluous), we find in it the following fundamental properties of _the halogens_ or salt-producers. With metals chlorine gives salts (such as sodium chloride, &c.); with hydrogen a very energetic and monobasic acid HCl, and the same quantity of chlorine is able by metalepsis to replace the hydrogen; with oxygen it forms unstable oxides of an acid character. These properties of chlorine are possessed by three other elements, bromine, iodine, and fluorine. They are members of one natural family. Each representative has its peculiarities, its individual properties and points of distinction, in combination and in the free state--otherwise they would not be independent elements; but the repetition in all of them of the same chief characteristics of the family enables one more quickly to grasp all their various properties and to classify the elements themselves.

In order to have a guiding thread in forming comparisons between the elements, attention must however be turned not only to their points of resemblance but also to those of their properties and characters in which they differ most from each other. And the atomic weights of the elements must be considered as their most elementary property, since this is a quantity which is most firmly established, and must be taken account of in all the reactions of the element. The halogens have the following atomic weights--

F = 19, Cl = 35·5, Br = 80, I = 127.

All the properties, physical and chemical, of the elements and their corresponding compounds must evidently be in a certain dependence on this fundamental point, if the grouping in one family be natural.[47 bis] And we find in reality that, for instance, the properties of bromine, whose atomic weight is almost the mean between those of iodine and chlorine, occupy a mean position between those of these two elements. The second measurable property of the elements is their equivalence or their capacity for forming _compounds of definite forms_. Thus carbon or nitrogen in this respect differs widely from the halogens. Although the form ClO_{2} corresponds with NO_{2} and CO_{2}, yet the last is the highest oxide of carbon, whilst that of nitrogen is N_{2}O_{5}, and for chlorine, if there were an anhydride of perchloric acid, its composition would be Cl_{2}O_{7}, which is quite different from that of carbon. In respect to the forms of their compounds the halogens, like all elements of one family or group, are perfectly analogous to each other, as is seen from their hydrogen compounds:

HF, HCl, HBr, HI.

[47 bis] See, for example the melting point of NaCl, NaBr, NaI in