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higher temperature produced by boiling with alcohol may also contribute towards the result.
While sulphate of soda and chloride of calcium dissolved in water are resolved by mutual decomposition into chloride of sodium and precipitated sulphate of lime, an aqueous solution of common salt and sulphate of lime evaporated to dryness, or a pulverized mixture of the two salts moistened with a large quantity of water and then dried, yields chloride of calcium when digested in boiling alcohol, while sulphate of soda may be extracted from the residue by water. (Sch 90, substituting Ca for Mg.) Unless the mixture of the two salts be moistened with water and
then dried, no chloride of calcium will be extracted from it by alcohol (Döbereiner, J. pr. Chem. 1, 112). The reversing of the affinity is perhaps produced by the alcohol: possibly however, as Döbereiner supposes, the affinity of sulphate of lime for sulphate of soda (a combination of the two salts occurs in nature forming the mineral called Glauberite,) may cause the formation of a small quantity of the double salt. By melting the two salts together in equal numbers of atoms, we obtain a hard mass, which becomes soft and afterwards moist by exposure to the air, and consequently contains a small quantity of chloride of calcium. These two facts serve to show how it happens that in analyses of mineral waters, when the residue after evaporation is boiled in alcohol, chloride of magnesium or chloride of calcium and sulphate of soda are frequently obtained, though these substances undoubtedly existed in the water as sulphate of lime or magnesia and chloride of sodium.
If 1 part of carbonate of potash be dissolved in at least 10 parts of water, and the solution shaken up with lime, the carbonic acid is taken up by the lime: with 4 parts of water however no decomposition takes place: on the contrary, a strong solution of caustic potash takes carbonic acid from carbonate of lime (Liebig, Pogg. 24, 365). The affinity of potash for carbonic acid is probably greater than that of lime; but when the quantity of water is increased, the affinity of that liquid for potash perhaps increases more rapidly than its affinity for carbonate of potash, and thus the first mentioned result is brought about.
Aqueous sulphurous acid dissolves iodine, forming a mixture of sulphuric and hydriodic acids; but if the quantity of water in the solution be diminished by evaporation, sulphurous acid is evolved and hydriodic acid containing iodine in solution is left behind. Similarly, concentrated sulphuric acid and hydriodic acid are resolved into sulphurous acid and iodine (Sch. 91). Hence the affinity of water for sulphuric and hydriodic acids gives rise to their formation.
The following experiment of Chevreul also shows the influence of the solvent. Excess of water abstracts half the potash from neutral stearate of potash, and forms bi-stearate of potash. Ether dissolves stearic acid from neutral stearate of potash, and separates a compound of stearic acid with excess of potash. Water has a more especial affinity for potash, and ether for stearic acid.
3. Difference of temperature may give rise to reciprocal affinity in
a. At high temperatures, the affinity of heat for that substance simple or compound which is most disposed to form a gaseous compound with it, often comes into play and determines the result. Heat, in such cases, acts like a fourth body introduced.
Peroxide of ruanganese mixed with hydrochloric acid at ordinary or slightly elevated temperatures, gives up its second atom of oxygen to the
hydrogen of the acid, so that water, chlorine, and chloride of manganese or hydrochlorate of protoxide of manganese are produced. (Sch. 64, 73.) If on the other hand chlorine gas be exposed to light or to a red heat in contact with water, hydrochloric acid is formed and oxygen gas evolved. At one time therefore oxygen (in peroxide of manganese) takes hydrogen from hydrochloric acid, setting chlorine free; at another, chlorine takes hydrogen from water and evolves oxygen gas. We may with probability suppose that the affinity of oxygen for hydrogen is greater than that of chlorine; on this hypothesis the explanation of the first case is evident. On the other hand, heat has a greater affinity for oxygen than for chlorine; for chlorine gas has been liquefied by strong pressure,—which oxygen has
When therefore heat acts with great intensity, its affinity for oxygen + that of chlorine for hydrogen effects the decomposition of water.
Potassium at a red heat decomposes black oxide of iron forming potash and metallic iron: at a white heat on the contrary potash is decomposed by metallic iron, the products being black oxide of iron and rapour of potassium. In this case it must be supposed that the affinity of potassium for oxygen is greater than that of iron; nevertheless, at a white heat the affinity of heat for potassium, with which it combines and forms a vapour, comes into play and determines the result.
At a red heat potassium decomposes carbonic oxide, forming potash and charcoal; at a low white heat charcoal decomposes potash, producing carbonic oxide gas and potassium vapour. (Sch. 6, substituting K for Zn.) In the latter case, the weaker affinity of carbon for oxygen is assisted by that of heat for carbonic oxide and potassium.
When potash (or soda) is in combination with phosphoric acid, boracic acid, or silica, sulphuric acid will separate these substances at ordinary temperatures and combine with the potash by virtue of its greater afti. nity. If on the contrary sulphate of potash be ignited in contact with phosphoric acid, boracic acid, or silica, these acids will take hold of the potash and separate the sulphuric acid in the state of vapour. In this case it is the affinity of heat for the volatile sulphuric acid, with which it forms a vapour, that enables the much weaker affinity of the above-mentioned acids for potash to gain the mastery.
When carbonate of ammonia is added to an aqueous solution of nitrate of lime, nitrate of ammonia and precipitated carbonate of lime are formed. But when a mixture of nitrate of ammonia and carbonate of lime is heated above 100° C. carbonate of ammonia volatilizes and nitrate of lime remains. (Sch. 92.) Here the result is determined by the affinity of heat for the volatile carbonate of ammonia. Similar relations are exhibited between carbonate of ammonia and chloride of calcium in the cold, and between sal-ammoniac and carbonate of lime at a higher temperature. For a similar reason, borate of ammonia and common salt act upon each other only at a high temperature, evolving sal-ammoniac in the state of vapour,
To the same category may also be assigned the following facts, so far as they may prove to be correct. When sulphuric acid acts upon zinc under the ordinary pressure of the atmosphere, sulphate of zinc and hydrogen gas are evolved (Sch. 17); but according to Babinet (Ann. Chim. Phys. 37, 183; also Pogg. 12, 523) this decomposition ceases when the process is conducted in a strong copper vessel closed by a stop cock, as soon as the disengaged hydrogen gas has attained a certain pressure; at 10° C, the decomposition and evolution of gas cease when the gas presses with a force of 13 atmospheres; at 25°, when the pressure amounts to 33
atmospheres. This seems to show that the affinity of zinc for oxygen + that of sulphuric acid for oxide of zinc is less than the affinity of hydrogen for oxygen – that of sulphuric acid for water, and therefore the decomposition does not take place under strong pressure;—but at lower pressures, when the affinity of heat for hydrogen, with which it forms a gas, likewise comes into play, the action goes on. On the contrary, Faraday has observed (Qu. J. of Sc. 3, 474) that at this increased pressure, the decomposition is not arrested but merely retarded, because the effervescence ceases and with it the motion of the liquid, by which the chemical action is materially assisted. I have also obtained the following results at a summer heat of from 20° to 30° C. (68° to 86° Fab.), using very thick and narrow glass tubes 5 inches in length. On filling the tube full of moderately strong hydrochloric acid, introducing a piece of zinc just above the acid, sealing the tube, and laying it in a horizontal position, it burst after four hours with a violent explosion. Now since a tube of equal strength is capable of holding liquid carbonic acid at 25° C. without bursting, and the elasticity of carbonic acid at that temperature amounts to 50 atmospheres, that of the hydrogen gas in the experiment just described must have exceeded 50 atmospheres. When a similar experiment is made with a mixture of 1 part of oil of vitriol and 8 of water, the tube does not burst even when left in the horizontal position for several weeks and placed upright every day. On cutting off the end, the gas escapes with a slight detonation without bursting the tube, and the acid is found to be nearly saturated with zinc. This seems to show that the decomposition is not arrested by strong pressure but only retarded.
The following experiment (recorded in Berzelius's Lehrbuch, 5, 9) is also connected with this matter. When pieces of carbonate of lime are placed in a strong glass vessel, a somewhat dilute acid poured upon them, and the vessel closed air-tight, the solution ceases after a time, and the lime is no longer attacked for whatever length of time it may be left in the acid; but on opening the vessel the lime is in a few minutes completely dissolved. From this it might be inferred that the affinity of carbonic acid for lime is greater than that of sulphuric, nitric, or hydrochloric acid, that these acids decompose carbonate of lime only under comparatively small pressure, when the action is assisted by the affinity of heat for carbonic acid with which it forms a gas,—that at strong pressures, on the contrary, carbonic acid would expel these acids from their combinations with lime. (Sch. 12.) But hydrochloric acid of moderate strength exhibits a different relation from this, at least according to my experiments. A quantity of acid like that above mentioned having been sealed up in a tube together with an excess of calcspar, the tube laid horizontally but turned upright every day in order to renew the points of contact, the liquid was found after 14 days to be covered with a highly moveable film of liquid carbonic acid, 2 lines in thickness. On cutting off the point, the upper half of the tube burst with a loud report, and the remaining liquid was neutral to litmus-paper. This experiment shows that hydrochloric acid decomposes carbonate of lime even under a pressure at which carbonic acid becomes liquid, and therefore that the affinity of hydrochloric acid for lime is greater than that of carbonic acid.
Lastly, with reference to this subject, we may mention an experiment of Petzhold (J. pr. Chem. 17, 464), according to which pulverized quartz heated to whiteness in an open vessel with an equal weight of carbonate
of lime expels the carbonic acid, but produces no such effect when the mixture is heated in a strong closed iron vessel. In this case, it may with great probability be supposed that the affinity of silica for lime is weaker than that of carbonic acid, and that the formation of silicate of lime takes place only when the action is assisted by the affinity of heat for carbonic acid,
6. In other cases, difference of temperature appears to modify the result in consequence of the cohesion (or affinity?) of bodies increasing and diminishing in different degrees at lower and higher temperatures, -and here in particular Berthollet's law regarding the decomposition of salts by double affinity finds its application.
A solution of common salt and sulphate of magnesia evaporated at ordinary temperatures, or a little above, allows both salts to crystallize out unaltered (page 127); but at 0° C, or at lower temperatures, as was long ago observed by Scheele, hydrated sulphate of soda crystallizes out and the solution retains chloride of magnesium; on gently warming the whole, common salt and sulphate of magnesia are again obtained. But above 50° C. (122° Fah.) the solution again deposits sulphate of soda, though in the anhydrous state. (H. Rose, Pogg. 35, 180.) These results may be explained by the different solubility of sulphate of soda at different temperatures: at 0°C. one part of sulphate of soda requires 8.2 parts of water to dissolve it, at 330 the smallest quantity, viz. 0:33, and at 50:4° again 0:38 parts. Below 0° and above 50° the solubility must be considerably less." Since now, according to Berthollet's law, the least soluble salt is always produced, sulphate of soda separates both below 0° and above 50°, because at these extremes of temperature its solubility is less than that of common salt or sulphate of magnesia; at medium temperatures, on the contrary, at which sulphate of soda is more able than common salt or sulphate of magnesia, these salts remain unaltered.
In a similar manner, sulphate of soda and chloride of potassium dissolved together in water resolve themselves, at ordinary temperatures, into sulphate of potash and chloride of sodium, whereas, according to Hahnemann and Richter (Stöch. 2, 224) a solution of the last-named salts at —20° C (-4° Fah.) deposits sulphate of soda. At ordinary temperatures sulphate of potash, at low temperatures sulphate of soda is the less soluble salt. According to Constantini, alum and common salt yield crystals of Glauber's salt at freezing temperatures, and according to Hahnemann, Glauber's salt crystallizes at very low temperatures, even from a mixture of saturated solutions of gypsum and common salt.
The explanation of the following cases—by supposing a disproportionate alteration of cohesion to be produced by change of temperature is not quite so satisfactory.
An aqueous solution of sulphate of lime gives with chloride of barium a precipitate of sulphate of baryta, while chloride of calcium remains in solution. (Sch. 52.) If, on the other hand, chloride of calcium be fused with sulphate of baryta, a mixture of sulphate of lime and chloride of barium is formed, the latter of which may be removed by rapidly boiling the powdered mass in water and filtering; but by longer standing under water the whole would again be converted into sulphate of baryta and chloride of calcium. Does the affinity of water for chloride of calcium contribute to this result?
Sulphate of baryta is decomposed both by fusion with carbonate of soda and by boiling with the aqueous solution of that salt (though but imperfectly), yielding carbonate of baryta and sulphate of soda: on the contrary, as Kölreuter has shown, carbonate of baryta is decomposed by digestion with sulphate of soda at ordinary temperatures, the products being sulphate of baryta and carbonate of soda.
A dilute solution of nitrate of lime remains clear when mixed with a solution of sulphate of soda, but deposits sulphate of lime when warmed. (Persoz.)
A solution of alum does not become turbid when mixed with very small quantities of carbonate of lime or soda, but on evaporation at a gentle heat, yields crystals of cubic alum. But when more strongly heated, the solution becomes turbid and deposits basic sulphate of alu. mina which redissolves as the solution cools.
A solution of pure acetate of alumina does not become turbid on heating, but undergoes that change if it contains sulphate of ammonia, potash, soda, or magnesia ; a fainter turbidity is produced by the addition of nitrate of potash, none by the addition of nitrate or acetate of baryta, chloride of calcium, or acetate of lead. The precipitate which consists of hydrate of alumina disappears when the solution is cooled, and appears again on heating. (Gay-Lussac, Ann. Chim. 74, 193, also Schw. 5, 49; further, Ann. Chim. Phys. 6, 201, also Schw. 21, 96.)
Persulphate of iron mixed with acetate of potash deposits hydrated peroxide of iron on boiling.
An aqueous mixture of borate of soda and sulphate of magnesia yields, on application of heat, a precipitate of borate of magnesia, which however disappears each time on cooling.
Metallic silver takes oxygen from persulphate of iron dissolved in water, when boiled in the liquid, so that a solution is formed containing sulphate of silver and protosulphate of iron; but on cooling, all the silver is reprecipitated in the metallic state, and the solution once more contains sulphate of peroxide of iron. (Sch. 94.)
In many other cases the occurrence of reciprocal affinity is only apparent.
When ammonia is added to neutral sulphate (nitrate or hydrochlorate) of magnesia, it is taken up and magnesia precipitated; on the other hand, magnesia expels ammonia from the neutral sulphate (nitrate or hydrochlorate) of ammonia and is itself dissolved. In both cases however the decomposition is only half complete, in whatever excess the ammonia or magnesia may be added. In the first case, half of the sulphate of magnesia remains undecomposed and unites with the sulphate of ammonia in the form of a double salt containing 2 atoms of sulphuric acid, 1 atom of ammonia, and 1 atom of magnesia; in the second, half of the ammoniacal salt remains undecomposed, and forms the same double salt with the sulphate of magnesia produced. (Sch. 95 and 96.)
Nitric acid added to chloride of potassium forms nitrate of potash, and sets hydrochloric acid free; on the other hand, nitrate of potash is converted by excess of hydrochloric acid into chloride of potassium. The affinity of potassium for oxygen + that of chlorine for hydrogen + that of nitric acid for potash is undoubtedly greater than the affinity of potassium for chlorine + that of hydrogen for oxygen; and thus the first case explains itself. (Sch. 97.) If, on the contrary, nitrate of potash is to be converted by hydrochloric acid into chloride of potassium, a great excess of the acid must be used and heat be applied: moreover the excess of hydrochloric acid does not expel the nitric acid in its unaltered state, but the two together are resolved into hyponitric acid (N O“), chlorine, and water. (Sch. 98.) Thus it is not nitric acid but the much weaker hypo