STUDIES O F SPARINGLY SOLUBLE SALTS, READILY OBTAINED FROM HOT SOLUTIONS O F REACTING SUBSTANCES. I BY K. P. CHATTERJEE AND N. R. DHAR
When solutions of two ionised substances are mixed at the ordinary temperature and if there is the possibility of the formation of a sparingly soluble salt, in general immediate precipitation takes place. It is well known that the precipitate as it first comes out is rather of colloidal nature and is voluminous; this bulky precipitate gradually passes into the less soluble, compact, crystalline form. This behaviour is observable in most reactions of analytical chemistry; however, there are some exceptions. When we mix very dilute solutions of a phosphate and magnesia mixture (magnesium chloride, ammonium chloride and ammonia) a precipitate is not immediately obtained. It is well known that the precipitate of magnesium ammonium phosphate appears on vigorous shaking or rubbing the inside of the containing vessel with a glass rod. If the test-tube, having a very small quantity of the mixture containing magnesium ammonium phosphate, be heated carefully it might appear that more precipitation takes place; but if the tube be heated so that the solution boils vigorously, it will be found that practically the whole of the magnesium ammonium phosphate dissolves at the boiling temperature. Similarly, if solutions of sodium hydrogen tartrate and potassium nitrate are mixed together in dilute solutions, the precipitate of potassium hydrogen tartrate does not appear immediately. Ostwald explains this phenomenon by supersaturation. The potassium hydrogen tartrate that forms in the solution, remains at first in a supersaturated state, but after a time the solution passes through the metastable condition and becomes unstable and crystals of potassium hydrogen tartrate begin to appear. If the mixture of sodium hydrogen tartrate and potassium nitrate be carefully warmed, crystallisation of potassium hydrogen tartrate might take place; but if the heating be continued to boiling, practically the whole of the potassium hydrogen tartrate will dissolve. Similar behaviour is observable in the precipitation of ammonium hydrogen tartrate, potassium chloroplatinate, magnesium ammonium phosphate, magnesium ammonium arsenate etc. Unlike these cases there is a class of substances in which precipitation is hastened by heating the solution to boiling. The well known case of this type is that of calcium citrate. If a citrate and a calcium salt be mixed together and if the solutions are not very concentrated there is no immediate precipitation in the cold. Precipitation takes place immediately on heatjng the mixture, and even if the solution be boiled, no appreciable amount of the precipitate redissolves. It is also known that calcium hydroxide and a few more calcium salts come out as precipitates when their solutions are boiled. We have observed that this phenomenon of increased precipitation on heating the reacting solutions is of very common occurrence. Thus if a nickel salt and an oxalate or an oxalic acid solution be
IO10
K . P. CHATTERJEE AND N. R. DHAR
mixed, there is hardly any immediate precipitation. But if the mixture be heated, even to boiling, a copious precipitate of nickel oxalate is obtained. We have investigated many cases of this type. Thus, oxalates of copper, zinc, manganese, ferrous iron, magnesium, thorium, cobalt, etc, behave exactly in a similar way to the nickel salt. On the other hand, oxalates of calcium, strontium, barium, cadmium, thallium, silver, lead, etc., behave in a quite different way, especially the last two. Thus if dilute solutions of silver nitrate and an oxalate be mixed, a precipitate of silver oxalate is obtained almost immediately. If the tube he heated to boiling practically the whole of silver oxalate is dissolved. At first sight, the difference between these two types of substances may be ascribed to the difference of solubility of these substances a t the different temperatures. It is a possible suggestion that calcium citrate, nickel oxalate, and copper oxalate, etc., are less soluble a t a higher temperature than at a lower one, whilst the solubilities of the oxalates of silver, lead, thallium, cadmium barium, strontium, calcium, etc., in common with most normal salts increase with increase of temperature. We will discuss this point more fully later on. We have observed that the behaviour of strontium citrate i s similar to that of calcium citrate, whilst the behaviour of the third analogue barium citrate is different from the other two This phenomenon is not restricted only to citrates and oxalates. Thus we have observed that tartrates of. nickel, thorium zinc, manganese, cobalt, magnesium, etc., behave in the same way as the oxalates of these metals. In other words, these tartrates are more readily precipitated in the hot state from mixtures containing a soluble tartrate and the respective metallic salt. On the other hand, tartrates of copper, cerium and lanthanum, appear to be more soluble in hot water than in cold. Moreover benzoates of nickel, cobalt, chromium, zinc, etc., come out as precipitates more readily from hot solutions than from cold ones. Also, fluorides of copper, manganese, cobalt, nickel magnesium, etc., appear to be precipitated more quickly from hot solutions. Evidently this phenomenon of the rapid appearance of a precipitate from a hot solution is not restricted to a few calcium salts as has been hitherto supposed, but appears to us to be of common occurrence. The object of the present paper is to investigate this phenomenon thoroughly and if possible to suggest explanations of this peculiar behaviour in all cases. It is apparent that the same explanation will not be applicable to all cases, they being so very diverse. The following might be a possible explanation of this peculiar behaviour. As has already been said, increased precipitation with increase of temperature may take place if the solubility of the salt, in question decreases with increase of temperature. In considering the above explanation of increased precipitation with increase of temperature it should be made certain that the salt is preserving the same composition throughout or whether any change in the amount of hydration is taking place with change of temperature. It can be assumed that the substance precipitated in the cold i s more hydrated and more soluble than the salt precipitated a t the boiling temperature. If these conditions occur, an explanation of the increased precipitation with increase of temperature is obtainable. Calcium citrate is a case in point and
SPARINGLY SOLUBLE SALTS
IO11
furnishes a very good instance of this type. We have observed and definitely proved that calcium citrate can exist as different hydrates containing different amounts of water. It has also been observed that the greater the degree of hydration, the greater the solubility. Thus, if a concentrated solution of sodium citrate and a concentrated solution of calcium chloride be mixed in the cold, a white gelatinous precipitate is immediately obtained. If the substance is rapidly separated from the mother liquor and pressed and dried, it is observed that the substance has the composition Ca3(C6H507)2 16Hz0. This white substance is very easily soluble in water and is hygroscopic in the ordinary air. The solution of this substance decomposes a t the ordinary temperature and readily gives a white precipitate of another hydrate of calcium citrate containing less amount of water of crystallisation. This white substance when purified and dried in air gives on analysis the formula Ca3 (C6H,0i)z6Hz0. If hot solutions of sodium citrate and calcium chloride be mixed, still another hydrate having the composition Ca3(CeHb07)24H,O is formed. The first hydrate containing the maximum amount of water of crystallisation is very easily soluble in water, whilst the hydrate containing 6H20is sparingly soluble and the hydrate containing 4Hz0 is the least soluble of the three and the solubility of the last hydrate decreases with increase of temperature. It appears to us that these facts go a long way in obtaining a complete explanation of the peculiar behaviour of citrates precipitated as calcium citrate in qualitative analysis. We have also investigated the case of strontium citrate and have got results similar to those of calcium citrate. We have already observed that Ca3 (CsH,0,)2 4H20 which is obtained by mixing hot solutions of sodium citrate and calcium chloride is less soluble at the higher temperature than in the ordinary temperature. It is well known that calcium hydroxide and some other calcium salts behave in the same way. It may be argued that some of the substances which show increased precipitation with increase of temperature may behave in a similar way. I n the case of calcium citrate as we have already said, the explanation of the increased precipitation a t the higher temperature rests on the formation of different hydrates with different solubilities a t different temperatures. But it may be assumed that in some cases the same hydrate will be formed when the solutions are mixed in the hot or in the cold. I€ it can be proved that the solubility of one and the same hydrate which does not change its composition when precipitated a t different temperatures is less a t a higher temperature than a t the ordinary temperature then we can have an explanation of this phenomenon of the increased precipitation a t high temperatures. If on the other hand the same hydrate obtainable in the hot or cold does not show a decreased solubility with increase of temperature we have to look to something else for an explanation of increase of precipitation with increase of temperature. As a matter of fact we have observed that the composition of several oxalates, tartrates, etc., remains consistently the same whether the salt is precipitated in the hot or cold. We have also observed that these salts do not show a decreased solubility with increase of temperature, but follow the normal behaviour of increased solubility with increase of temperature. We have
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K. P. CHATTERJEE AND N. R. DHAR
tried to find an explanation for this class of substances in the phenomenon known as “velocit,y of precipitation”. From chemical dynamics, we know that each chemical change has its own velocity or speed, and that rise of temperature markedly accelerates all chemical changes. The influence of temperature on a chemical change depends on the nature of the change itself, but it has been proved that the slower the change the greater is the influence of temperature’. Evidently, temperature should have some effect on the velocity of precipitation of substances. It also follows from our knowledge of chemical dynamics that instantaneous reactions as well as instantaneous precipitations are hardly affected by increase of temperature. I n other words, it will be practically impossible to investigate the influence of temperature on a very rapid precipitation phenomenon such as the formation of silver oxalate or silver chloride. On the other hand temperature should have a marked effect on a slow precipitation process. Consequently, if we assume that the velocity of the appearance of nickel oxalate when dilute solutions of a nickel salt and anoxalate or oxalic acid be mixed, be not extremely great but has a measurable value, we are forced to the conclusion that the temperature will have someeffect on this velocity and will accelerate the precipitation. In other words if we assume that the appearance of nickel oxalate takes a certain amount of time, then temperature will hasten the precipitation as it does accelerate all chemical changes. Evidently, the phenomenon of increased precipitation of substances like nickel oxalate, copper oxalate, etc., a t higher temperatures can be explained on the basis of a measurable velocity of the appearance of the precipitate and the influence of temperature on it. The experimental work in this paper will consist of: (I) determinations of the composition of the substances precipitated from the hot and the cold solutions; (2) determinations of the solubility of these substances a t various temperatures; (3) determinations of the transition temperature of one hydrate into the other; and (4) investigation of the velocity of precipitation and the influence of temperature on it. I n order to find out whether salts prepared at the ordinary temperature differ in composition from those prepared in boiling condition, the reacting substances were mixed in one case a t the ordinary temperature of the laboratory and in another case whilst they were boiling. The salts prepared in this way were freed from the mother liquor and unless otherwise stated, they were dried in air. Then they were analysed by suitable methods. I n order to determine the solubility a t the ordinary temperature and at the higher temperature, a sufficient quantity of the salt is shaken with distilled water in a clean Jena glass flask. For the determination of the solubjlity at the ordinary temperature, the mixture is shaken until no more of the salt dissolves, and the liquid is then filtered, and about 2 0 0 cc of the filtrate is evaporated in a platinum dish. I n most cases the residue was converted into sulphate with a little sulphuric acid and weighed as anhydrous sulphate. From the weight of the salt obtained, the amount of the original salt dissolved ‘Dhar: Ann. chim. ( 9 ) 11, 130 (1919).
SPARINGLY SOLUBLE SALTS
1013
per litre is calculated. For the determination of solubility a t the boiling temperature, the mixture of salt and distilled water was boiled for a sufficient length of time for equilibrium to be attained. The liquid is then quickly filtered in a hot-water funnel and the temperature of the filtrate running down is noted. A measured volume is taken and evaporated and estimated as in the previous case. Calcium Citrate In the following pages, it will be shown that we have got four definite hydrates of calcium citrate. When approximately twice normal solutions of calcium chloride and sodium citrate are mixed, a precipitate similar to starch paste is obtained which is readily soluble in water. Addition of alcohol reprecipitates it from its aqueous solution. It is difficult,to determine the composition of this salt, as even a t ordinary temperature it passes into another hydrate which is described later on and which is sparingly soluble. In presence of an excess of alcohol, the velocity of transformation becomes much less, but as soon as the alcohol is filtered off, the salt readily passes into another hydrate. Thie unstable hydrate is very hygroscopic. Attempts have been made to analyse the unstable hydrate by carrying on the precipitation at o°C and also by rapid centrifuging and quick washing and drying. The following figures will give an idea of the approximate composition of the salt. I n the following analyses as well as in the other cases, a known weight of the substance is taken in a platinum crucible and heated strongly. A few drops of dilute sulphuric acid are added and the excess of acid evaporated and the residue is heated and weighed as sulphate. I n the following, only one result of analysis for each salt is given. 0.4386 gm. of the hydrate gives 0.2262 gm. CaS04 . ' . C a = 15.1% Calc. for Ca3(CsH507)216H20,C a = 15.2% Hence the unstable and easily soluble hygroscopic calcium citrate has most probably the composition Ca3(CoH507) I 6H20. It has been already observed that the aqueous solution of the unstable hydrate gradually gives a precipitate of another hydrate. This hydrate which is much less soluble than the unstable hydrate can be readily obtained by mixing solutions of calcium chloride and sodium citrate a t the ordinary temperature and setting the mixture aside for sometime. This is the hydrate which is generally obtained when a moderately concentrated solution of calcium salt and a citrate are mixed. The salt is filtered, washed, dried in the air at the ordinary temperature, and analysed. 0.3758 gm of the salt gave 0 . 2 5 4 2 gm C a S 0 4 . ' . C a = 19.8% Cdc. for Ca3(C6H507)26Hz0, C a = 1 g . 8 7 ~ . More than ten samples of the salt were made by mixing calcium chloride and sodium citrate of different concentrations at the ordinary temperature and in all cases, the hydrate obtained a t the ordinary temperature has the composition Ca3(C6H507)26Hz0. We could not confirm the existence of the hydrate Ca3(CsH507)2 7H20,described in Beilstein's book.
1014
K. P. CHATTERJEE AND N . R. DHAR
This hydrate Ca3(C6&07)~ 6 H z 0 is fairly stable a t the ordinary temperature, though it has a very slight tendency to pass into the lower hydrate which will now be described. If the hydrate Ca3(C6HS07)z6 H z 0 be kept in the sun or if it is heated, it loses water and passes into the lower hydrate. When boiling solutions of calcium chloride and sodium citrate are mixed, an immediate and abundant white precipitate is obtained, which is very sparingly soluble. This salt is freed from mother liquor, dried in air and anslysed. 0.3074 gm salt gave 0 . 2 1 9 2 gm CaS04. . ' . Ca=20,97% Calc. for Ca3(C6Hs01)z&O, C a = 2 I . 05% Hence the precipitate immediately obtained by mixing boiling solutions of calcium chloride and sodium citrate has the formula Ca3(CsH507)2 4H&. If any of the above hydrates be heated in an air oven to I ;o"approximately, it gradually loses weight and after a time there is no further change in weight. This substance was analysed with the following results. 0.1156gmsaltgaveo.og3z gmCaSOr. . ' . Ca=23.7% Calc. for Caa(Cs&O.;)z $6 HzO, Ca = 23.6%. This salt in contact with water rehydrates itself and forms Ca3(C6H507)2 4Hz0. It has already been said that an aqueous solution of the unstable highest hydrate gradually gives the precipitate of Ca3(C6H507)2 6M20, in the cold. If however the sohit8ionbe boiled, Ca3(CsHs07)z 4H20 is copiously precipitated. We have also found out the limit of temperature below which only the hydrate Ca3(C6HS07)26Hz0 is precipitated, by effecting the precipitation a t different temperatures which were 32", 40°, 50") 60" 7o0, goo, 90°, 98". Precipitation a t temperatures up to 60°C produced the salt Ca3(C6H507)2 6Hz0, whereas a t 70" and above the salt Ca3(C6H507)24Hz0 was produced. The above results are interesting as they show the limit of temperature beyond which Ca3(C6Ha07)2 6H20 cannot he precipitated. IP other words, in order to precipitate Ca3(C6H5C>7)z 6 H z 0 we should work att a temperature below 7oOC. The results also explain why the hexahydrate decomposes when kept in the sun or slightly heated. This hexahydrate has a very small velocity of decomposition into water and the tetrahydrate when exposed. On keeping for a sufficiently long time in air at a temperature of about 3ooC, for a month or so, the salt is found to be partially decomposed, a8 is seen from the following results. 0 . 0 8 8 0 gm of the exposed salt gave 0.0614 gm CaSOl . ' . Ca=2o. j% Apparently the salt has decomposed into water and the tetrahydrate. We have already said that the hydrate Ca3(C6Hb07)2M H20 readily combines with water and forms the tetrahydrate. Consequently out of the four hydrates which we have obtained (viz, with 16, 6, 4 and % molecules of water) the tetrahydrate is the most stable. We also tried to find out the velocity of precipitation of calcium citrate at ,: different temperatures. With this object in view, solutions of calcium chloride and sodium citrate were mixed at different temperatures and the interval necessary for the appearance of a precipitate was noted. Here are some re-
SPARlNGLY SOLUBLE SALTS
1015
stilts. At the temneratures, 3 2 " 40°, 50' 60°, 70' 80")90" and 98" the intervals respectively were 20 min, I j min, 2 min, I/Z min, and instantaneous in the last four cases. 'I hese results will be discussed later on. Solubility of the citrates of calcium has been investigated. We have already observed that the unstable hydrate is extremely soluble. No quantitative measurement of the so1ubilit)yof this hydrate could be made because of the extremely unstable nature of the substance. 7 he solubility of the hexahydrate ana the tetrahydrate was determined in khe manner indicated previously. The solubility of Can(C6H507)Z I / Z HzO could not he obtained because on mixing with water this hydrate passes into the tetrahydrate. Salt
Temperature
C'an(CsH~0i) 2 6Hz0
30' 85O 30"
Cad C6H 5 0 7 1 2 4H20
9.5"
Grams per litre (of the
salt)
Grams per litre (of the anhydrous salt,)
45 F P
2.01
2.76 gin 2 . 5 1 gm 2 . 1 gm
2.27
2
'
gm gm 2 . 2 gm I . 83 gm
A t a glance it will be seen that the solubility of the hexahydrate increases with temperature, whilst the solubility of t,he tetrahydrate decreases with temperature. When Caa(C&07)2 6H20is boiled with water for a sufficiently long time it passes into Can(C6H507)24H20 as the following results will show. 0.3818 gin of the salt obtained from boiling the hexahydrate in contact with water for a suficientlylong time gave 0 . z 743 gm CaS04. . . Ca = 21. I % Calc. for Ca3(C6HSOi)2 4H20, Ca = 21.05% As has been already said the t.etrahydrate is the most stable of the hydrates. Exposure to sun or moderate heating does not decompose it. The tetrahydrate, when kept in water even for a long time: does not reform the hexahydrate. On the other hand the hexahydrate when kept in the sun for about 1 2 hours, is converted into the tetrahydrate. If the hexahydrate is heated in the water bath at. so', it loses two molecules of water of crystallisation and passes into t)he tetrahydrate in eight hours. Decomposition of the hexahydrate was also investigated a t 65', 70" and 80" and is much quicker than a t 50°C. At 65OC, exposure for an hour and a half is sufficient to decompose 0.3250 gm of the hexahydrate into the tetrahydrate. The solubility curves of the hexahydrate and the tetrahydrate intersect a t 52°C. I n other words, this is the transition temperature of the two hydrates and a t this temperature the two hydrates can be kept for an indefinite period in contact with water. This is remarkable, as this temperature is not far from that beyond which the hexahydrate is not precipitated as we have already shown, and when by mixing calcium chloride and sodium citrate we get thn tetrahydrate. From the foregoing results we get a satisfmtory explanation of the peculiar behaviour observed on mixing a citrate and a calcium salt. In ordinary text boolcs of organic chemistry, in explaining the reactions involved in the test '
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K. P. CHATTERJEE AND N. R. DHAR
of the citrate, the only suggestion thrown out is that calcium citrate is precipitated in the hot and not in the cold, because calcium citrate is less soluble a t a higher temperature. We presume we have given a complete explanation of this peculiar behaviour. If a calcium salt and a citrate be mixed, there is the possibility of the formation of the readily soluble hydrate Ca3(C6Hs07)2 16H20. But as has been proved, this hydrate is unstable and readily passes into the hexahydrate, which is generally obtained by mixing moderately concentrated solutions of the calcium salt and a neutral citrate in the cold. But if the solutions are dilute the hexahydrate is not usually precipitated in the cold, but on boiling, the sparingly soluble tetrahydrate is readily precipitated, because, like all chemical changes, temperature markedly accelerates the velocity of precipitation. In testing for a citrate by calcium salts, the precipitate which is obtained from a boiling solution is the sparingly soluble tetrahydrate Ca3 (C6Hs07)z 4H20. The above is the explanation of the peculiar behaviour in the test, for a citrate by a calcium salt. I n the following citrates, oxalates and tartrates, only one result of analysis, as a rule, is given. Strontium Citrate The hehaviour of strontium citrate is more or less allied to that of the calcium salt. When concentrated solutions of strontium chloride and sodium citrate are mixed, a precipitate of the nature of the unstable calcium citrate is obtained. This salt is very unstable and is readily soluble in water. This hydrate is too unstable to be partially purified and analysed and very quickly passes into the next hydrate of which the analysis is given below. This second hydrate is to be carefully dried in the shade. Exposure to tropical sunlight makes it lose its water of crystallisation until it attains the composition Sr3 (C6H507)2HzO0.6076 gm of the salt gave 0.4564 gm. SrS04 . * . Sr=35.84y0 Calc. for Sr3(CeH50,)z ~ H z OSr , '35.95% Henre it is evident that by mixing a cold solution of strontium salt and a neutral citrate, we get a white precipitate, which, when air dried at the ordinary temperature, has the formula Sr3(C~H607)2 5HzO. This salt decomposes if it be placed in the sun. A certain specimen of the salt was exposed to the sun for a short time and analysed. 0.1620 gm of the salt gave 0.1234 gm SrS04 . * . Sr =36.34%). It is apparent that the salt has partly decomposed on exposure to tropical sunlight. Another sample was exposed to sunlight for a sufficient length of time until a constant weight was obtained and then it was analysed with the following results. 0.2770 gm of the salt gave 0.2300 gm. SrS04 . * . Sr=39.61% Calc. for Srs(C6H50,)zH20, Sr=39.9Y0 Sr3(C6H607)2 5Hz0 remains unchanged even if it be boiled with water. A certain sample of the salt, prepared in the cold, was boiled with water for some time and was dried in air and analysed. 0.2144 gm of the boiled salt gave 0.1624 gm SrS04 . '. Sr=36.14%
1017
SPARINGLY SOLUBLE SALTS
Calc. for Sr3(C&07)~ 5Hz0, Sr = 35.91% If a solution of the unstable, readily soluble, strontium c i t k t e be boiled,
a copious precipitate is immediately obtained, which is apparently a mixture of Sr3(C,H507)z sH,O and Sr3(C6H&)zH20 8s the following results of analysis will show. 0 , 1 0 7 0 gm of the salt thus precipitated gave 0.0848 gm PrSO., . ' . Sr = 37.8675 Sr3(C7Hb071z HzO can be readily obtained by mixing boiling solutions of a strontiuni salt and a citrate. This salt is very stable. The pentahydrate when heated or exposed to tropical sunlight passes into this hydrate When this monohydrate is kept in contact with water, the pentahydratp is not reformed. 0 . 2 2 2 7 gm of the salt thus kept and then analysed gave 0.1847 gm of SrS04 . ' . Sr=39.6% Calc. for 8r3(CaH0)2 HzO, Sr = 3 9 . g s 1 Many samples of this monohydrate were obtained by mixing boiling solutions of a strontium salt and a citrate and analysed with the followjng results. 0.1294 gm of the salt gave 0.1080gm. SrS04. . ' . Sr=39.8% Calc. for Sr3(C~H507)2H20, Sr=39.g% Evidently it i s clear that by mixing boiling solutions of a strontium salt and a neutral citrate, the white precipitate obtained has the formula Sr3 (C6H607)2 HzO. An old sample of Sr3(C6H507)2 H 2 0 was kept in an air oven a t 110' for a number of days with no loss of weight. Consequently, at I IO', it is not possible to obtain a hydrate lower than the monohydrate. I n this respect strontium citrate differs from calcium citrate. The solubilities of the two hydrates a t different temperatures were determined in the same way as that of the calcium salt and the following results werc obtained. Salt
Sr3(Ci&07)3 6Hz0 Sr3(C6H507)2 HzO
Temperature
30° 95O 3oo
Grams per litre (of the salt)
1.55 1.79
Grams per litre (of the anhydrous salt) I . 26
1.57
3-05 2.97 95O I .83 1.54 It is to be noted that the intersection of the solubility curves of the pentahydrate and the monohydrate is beyond IOOOCas will be shown by producing the curves until they meet. The pentahydrate seems, therefore to be stable a t IOOOCin presence of water.
Nickel Oxalate 2 0 0 cc of seminormal nickel chloride and 2 5 0 cc of seminormal oxalic acid were mixed and kept a t 29OC. An excess of oxalic acid was used to prevent the formation of a basic salt. After enough precipitate separated, it was collected, washed and dried. The analysis of the nickel oxalate was carried out in the following way. A weighed quan'tity of the salt was heated in a platinum crucible (I). for some time and a drop of nitric acid was added and reheated and weighed
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K. P . CHATTERJEE AND N . R. DHAR
as NiO. ( 2 ) . .A weighed quantity of the salt was dissolved in sulphuric acid and the solution was titrated with standard permanganate. (3). By means of combustion analysis, the amounts of water, carbon dioxide and nickelous oxide, left as residue, were directly determined and thus a complete analysis was made, from which the formula was calculated. 0.4950 gm of the salt gave 0.1890 gm NiO. . ' . Ni=30yc (I) Calc. for 3NiC2048Hz0, N i = 3 o . 1% 0.3350 gm of the salt on titration with KMn04 showed 0.2674 gm of (2) anhydrous salt . ' . Ni = 30.14% and H,O = 24.6% Calc. for 3NiCzOl 8Hz0, Ni = 3 0 . 1 % and HzO = 2 4 . 6 % (3) 0 . 1 7 9 2 gni saIt on combustion gave 0.0688 gm NiO, 0.0442 gm HzO and 0.0806 gm COZ, . ' . Ni=30.2%, H z 0 = z 4 . 7 % and COZ=45% Calc. for 3 N C 2 0 48 H ~ 0Ni=30.1YG , H 2 0 = 2 4 . 6 % and COZ=45.I%. Consequently, by mixing solutions of a nickel salt and oxalic acid, a t the ordinary temperature, the formula of the oxalate obtained is 3NiC2048Hz0. By mixing boiling solutions of nickel chloride and .oxalic acid, a copious precipitate was immediately obtained. The salt was dried and analysed by the methods already indicated. Only one analysis is given below. 0.1482 gm of the salt gave on combustion 0 . 0 5 9 2 gm NiO, 0.0330 gm of H z O a n d 0 . 0 6 8 4 g m C O n .. ' . Ni=31.4%, HZO=zz.3Yc,C02=46.170 Calc. for 3NiCz04 7Hz0, N i = 3 1 . 1 7 ~ ,H Z O = z z . 3 % and COz=46.6% Consequently, the formula of the salt obtained by mixing boiling solutions of nickel chloride and oxalic acid is 3Nj C 2 0 47Hz0, whilst the salt obtained a t the ordinary temperature contains one molecule more of the water of crystallisation and has the formula 3NiC2048Hz0. The formula of thc nickel oxalate given in Beilstein is NiC204zH20. We could not verify this result.
Cobalt Oxalate The salt was obtained at, 30' and a t the boiling temperature. The following are the analyses, one analysis in each case being given. ( I ) 0.1046 gm of the salt precipitated a t 30°, gave 0.0450 gm Co304, 0.0496 gm COz and 0 . 0 2 1 5 gm HzO. . * . c o = 3 1 . 6 % , coz=47.45 Calc. for 4CoCz04 9HzO, C o = 3 1 . 7 ~ Co2=47%. ~, Consequently the hydrate obtained has the formula 4CoCpO4, 9 H z 0 (2) 0 . I I I 2 gm of the salt precipitated from boiling solutions, gave 0.0490 gm of Coa04, 0.05334 gm COz and 0 . 0 2 2 6 gm HzO . ' . C o = 3 2 . 1 % , C O z = 4 7 . 5 % a n d H z 0 = z o .1%. Calc. for CoCz04zHz0, Co = 3 2 . z%, GOz= 48.1% and H2O = 1 9 . 7 % Hence the salt precipitated at 100' has the formula Co CZO, 2H2O. I n Beilstcin, we find that the oxalate described has the formula CoCz04 2Hz0.We have got this hydrate as well as another new hydrate of the formula 4 Co CZO4gHzO,which is obtainable by mixing oxalic acid and a cobalt salt in the cold and keeping the mixture for some hours.
1019
SPARINGLY SOLUBLE SALTS
Copper Oxalate The following are the analyses, only one analysis in each case being given. (I) 0.1368 gm of the salt, precipitated at 3ooC, gave on combustion, 0 . 0 7 1 2 gm CuO, 0.0792 gm CO? and 0.0006 gin HzO . '.Cu=41.6%, C 0 ~ = 5 8 . 7 %and H 2 0 = 0 . 2 % Calc. for C U C Z O C ~ ,u = q ~ . g O j C , O z = 5 8 . 1 7 ~and HzO nil. Hence the formula of the salt precipitated at 30°C is CuCzO4. 0 . I 104 gm of the salt precipitated from boiling solutions, gave on com(2) bustion 0 . 0 5 7 5 gm CuO, 0.0639 g n ~COz and 0 . 0 0 0 8 gm HzO. . .Cu=41 . 6 % l c02=57.9% and H20= . 7 % Calc. for C)uCzO4,Cu =41,9'31c, COz= 58.1% and HzO nil. Hence, copper oxalate precipitated at IOOOChas the formula CuCzO4. Beilstein describes two hydrates of the formula CuCz04HzOand CuC204 1 / 2 H20. Both in the hot and the cold we got the anhydrous CuCzOd precipitated. '
Manganese Oxalate 0 . 2 9 1 0 gm of the snlt precipitated at 3ooC, gave on titration with (I) K M n 0 4 0.2323 gm of anhydrous oxalate and 0.0587 gm (by difference) of water. . ' . M n = 3 0 . 7 1 ~ ~ a n d H ~ O = 2 0 . 1 ~ ~ Calc. for MnCz04 2Hz01Mn=30.7% and 2 0 . 1 % . (21 0.3683 gm of the salt precipitated from boiling solutions, gave on titration with KMnOd 0.2934 gm of the anhydrous oxalate. . Mn = 30.65% Calc. for MnC204 2Hz0, Mn =30.7(rc The solubility of this oxalate was determined a t 36OC and 93OC. 2 0 0 cc of the clear filtrate obtained after saturation with the oxalate, were evaporated and the residue was strongly heated and weighed as MnsOl. '
Temperature
,
Grams of anhydrous salt per litre
Salt precipitated 36O -375 in the cold 93 O .780 Salt precjpitated 36O ,375 from boiling 93O .789 solutions Evidently the same hydrate is obtained a t the ordinary and also at the boiling temperatures and the solubility of the salt increases with increase of temperature. Beilstein has described two hydrates, namely containing 2.5 and 3 molecules of water of crystallisation, but we got only the hydrate containing zHz0 both in the hot and in the cold.
Magnesium Oxalate ( I ) Concentrated solutions of magnesium sulphate and oxalic acid were mixed and kept overnight a t the ordinary temperature. A white crystalline precipitate was obtained. It was washed, dried, and analysed. Only one analysis is given below. 0.2576 gm of the salt precipitated at 30°C, gave 0.0608 gm MgO . ' .Mg=16.8Y0
K. P. CHATTERJEE A N D N. R. D H A R
I020
Calc. for MgC204 2H20, Mg= 16.2% Hence magnesium oxalate precipitated from concentrated solutions of magnesium sulphatp and oxalic acid at 3ooC has the formula MgC204. zHzO. (2) When concentrated solutions of magnesium sulphate and oxalic acid are mixed in boiling conditions, a copious white precipitate is immediately obtained. The salt was washed, dried, and analysed, one analysis being given underneath. 0 . 2 6 4 2 gm of the salt gave on titration with KMn04, 0.1944 gm of the anhydrous oxalate. . ' . Mg= 16.1% Calc. for MgC204 zHsO, Mg= 1 6 . 2 % Hence the formula of the salt precipitated at roo' is the same as that precipitated at 30'; it is MgCz04 zHz0 For solubility, 2 0 0 cc of the solutions were evaporated and the residues were treated with dmal; quantities of sulphuric acid and heated until fuming ceased. These were weighed as MgS04and calcnlations were made therefrom. Temperature
Salt precipitated in the cold Salt precipitated from boiling solutions.
Grams of anhydrous salt per litre
36'
*
9z0 3 6'
.302
92O
338 .406 .393
Hence the same salt is precipitated both in cold and in hot. This salt also is more soluble at higher temperatures than a t the ordinary temperature. Kohlrauschl prepared supersaturated solutions of magnesium oxalate by neutralization of oxalic acid with magnesium hydrate and determined their molecular conductivities.
Ferrous Oxalate (I) 0.1872 gm of the salt precipitated a t 3ooC gave 0.0824 gm of Fez03. . ' .Fe=30.8% Calc. for FeC204 zHz0, Fe=31.1% When ferrous sulphate and oxalic acid are mixed, there is no immediate precipitation in the cold, but copious precipitation takes place when boiling solutions are mixed. (2) 0.1756 gm of the salt precipitated from boiling solutions gave 0 . 0 7 7 2 gm. Fez03. . ' . Fe=30.8%. Calc for FeC204 2H20, Fe=31.1% Evidently, the hydrate FeC204 2H20 is precipitated both in the hot and in cold. Zinc Oxalate 0 . 7 2 5 8 gm of the salt, precipitated at 3ooC, gave on titration with (I) potassium permanganate, 0.5840 gm of the anhydrous salt. . ' . Zn=34.3Yc Calc. for ZnCzO4, zHzO, Zn=34.6% I
Sitzungsber. Akad. Wiss. Berlin, 1904, 1223.
SPARINGLY SOLUBLE SALTS
I02I
When solutions of zinc salt and oxalic acid are mixed in the cold there is no immediate precipitate, but the precipitate of zinc oxalate comes out after long time. On the other hand a copious white precipitate is obfiained when boiling solutions of zinc salt and oxalic acid are mixed. The salt obtained a t the boiling tempel'ature was washed, dried and analysed with the following result. (2) 0,9829 gm of the salt showed on titration with potassium permanganate 0.7900 gm. of anhydrous zinc oxalate. . ' . Zn = 34.2% Calc. for Zn C204 2Hz0, Zn=34.6% Consequently the hydrate ZnCzO,, 2H20 is precipitated both in the hot and in cold, when a zinc salt is added to oxalic acid.
Thorium Oxalate Thorium nitrate and oxalic acid were mixed and left for a long time a t A white precipitate was obtained which was washed, dried and analysed. 0.1296 gm of the salt gave 0.0662 gm Tho2 . ' . Th=44.g% Calc. for Th(Cz04)2 6Hz0, Th=45Y0 (2) When boiling solutions of thorium nitrate and oxalic acid are mixed, a copious white precipitate is obtained and the substance has a tendency to pass into colloidal state. The salt was carefully washed and dried in air and analysed. 0 ,I 774 gm of the salt gave o.ogog gm Tho2 . ' . T h "44.8% Calc. for Th(C204)2 6H20, Th=45% Hence the formula of thorium oxalate precipitated both hot and cold, is Th(Cz04)Z 6H&. Brauner' gives the composition of the air dried salt by the formula T h (C204)Z 6H20,obtained by mixing in the cold an oxalate and a thorium salt solution. We have obtained the same composition from boiling solutions as well. I n Beilstein the hydrate described is T h (C204)2zH20 which i s obtained on partial dehydration of the hexahydrate over sulphuric acid. (I) 30'.
Magnesium Tartrate (I) 0.4804 gm of the salt, precipitated a t 30°, gave (on keeping in a hot air oven until a constant weight was obtained) 0.3836 gm anhydrous magnesium tartrate, from which 0.2670 gm anhydrous magnesium sulphate was got. . ' .Mg= I I . 1% and HzO= 2 0 . 7y0 Calc.forMg(C4H406),2j.iH20, M g = 1 1 . 1 % a n d H ~ O = 2 0 . 7 % When boiling solutions of magnesium sulphate and Rochelle salt were (2) mixed, the precipitation was immediate. 0.6966 gm of this salt gave 0.3874 gm MgSOd. . ' . Mg= I I . 2 % Calc. for Mg (C4H406) 2 % HzO, M g = I I . I%,. (3) If any of the above salts be heated to 170°, until a constant weight ip obtained, anhydrous magnesium tartrate, as proved by analysis, is obtained. This anhydrous salt, when moistened with water and dried on a water bath remains unchanged. For solubility, the following results were obtained. 1
J. Chem. SOC.73,984 (1898).
K. P. CHATTERJEE AND N. R. DHAR
I022
Salt
Magnesium tartrate precipitated in the cold Magnesium tartrate precipitated in the hot
Temperature
Grams of the anhydrous salt per litre
30"
7.6
90" 3 oo
14.4 7.5
90"
14.5
It is evident the same salt is precipitated both in hot and cold. Beilstein gives Mg(C4H4Oe) 4 H 2 0 , Mg(C4H406) 3 H 2 0 and anhydrous magnesium tartrate. We could not verify the first two. Manganese Tartrate In the cold, this salt is slowly precipitated, but it comes out in a granular form. The amorphous salt is very quickly formed when boiling solutions are mixed. ( I ) 0.5060 gm of the salt precipitated a t 30°, gave 0,4334 gm anhydrous salt, which yielded 0.3260 gm manganous sulphate . ' . Mn = 23.3%,. Calc. for Mn(C4H4o6)2H20, Mn = 2 3 . 1 7 ~ I . 0904 gm of the salt, obtained from boiling solutions, gave 0,6936 gm (2) manganous sulphate. . ' . Mn = 23. ~7~ Calc. for Mn(C4H406)zHzO Mn = 2 3 . 1 7 ~ The anhydrous manganous tartrate is obtained by heating either of the above salts to 14oOC. On moistening, this does not get hydrated again. The same salt is precipitated both hot and cold. 1000 gm of the solution a t 25°C contains 2 . 4 gm, and a t 90°C 2.45 gm of anhydrous manganese tartrate. Cobalt Tartrate If concentrated solutions of cobalt sulphate and Rochelle salt be mixed, and if the mixture he kept undisturbed, a finely granular purple salt is gradually precipitated. If the mixture be stirred or if the solution be added hot, an amorphous pink salt is obt'ained. I n this, as well as in the previous case, the precipitate is rather retarded if excess of the tartrate be added, owing probably to the formation of complex salt's. ( I ) 0.6956 gm. of the purple salt precipitated in the cold, gave on drying 0.5714 gm of anhydrous salt, which yielded 0.4290 gin of anhydrous cobalt sulphate. . ' . Co = 23.4% and H20 = I 7.8% Cdc. for Co(C4H406) 2 ?&2O, Co = 23.4 % afid HzO= 17.670 Dehydrat,ion of the above salt slowly t>akes place even at 6ooC. On moistening and drying the anhydrous salt, it does not get hydrated again. (2) 0.3570 gm of the pink salt precipitated from hot solutions gave on drying 0.2938 gm of the anhydrous salt which yielded 0 . 2 148 gm anhydrbus cobalt sulphate. . ' . Co = 23.4% and HzO= I 7 . 77T0 Calc. for C O ( C I H ~ O2 ~ ) HzO, Co = 2 3 . 4 % and HzO = I 7 . 6 % The pink salts have composition and properties similar to those of the purple one. The only point of difference is the appearance and crystalline form.
SPARINGLY SOLUBLE SALTS
1023
gm of the solutions of both purple and pink salts contain a t 29°C anhydrous cobalt tartrate, and a t 84'C, 2.45 gm of the same. The molecular volume of HzO in these comes out to be 15. 1000
I . 85 gm
Nickel Tartrate
It is not so easy to prepare this salt as the cobalt compound. In the cold, the precipitation is extremely slow and the salt is fairly soluble in water. (I) 0.2486 gm. of the salt, precipitated a t 3ooC gave 0.1966 gm anhydrous tartrate . . HzO = 2 0 . g y o 0.3240 gm of the original salt gave 0 . 0 9 2 0 gm NiO, . ' . N i = 2 2 . 3 % Calc. for Ni (C4H406) 3Hz0, Ni=22.5% and H 2 0 = z o . 7 % ( 2 ) 0.4608 grri of the salt, precipitated when boiling solutions of nickel chloride and Rochelle salt are mixed, gave 0.1368 gm NiO. . ' . Ni = 23 . 3 % Calc. for Ni(C4H406) 2 % HzO, Ni = 2 3 . 3 % (and HzO= 17.8%) (3) Anhydrous nickel tartratmeis prepared by heating a t 17oOC any of the above salts until the weight, is constant. Moistening the anhydrous salt with water, gives a mixture containing 2 and 3 molecules of water of crystallisation, as the following shows. 0.1653 gm of the substance obtained by moistening the anhydrous salt and drying it in air, lost 0.03 I 2 gm on keeping it at 170°C . ' . Loss of water = 18.8% Calc. for Ni(C4H406)2 % HzO, HzO= 17.8% and for Ni(C~H406)3H20, HzO=20. 7%. '
Calcium Tartrate When concentrated solutions are mixed at the ordinary temperature, a jelly-like precipitate is formed. This passes rapidly to a granular mass with evolution of heat, so much so that ether can be made to boil with this heat of transformation. (I) 0,2402 gm of the salt precipitated in the cold, gave 0.1246 gm CaS04 . ' . Ca=15.3yo Calc. for Ca (C4H406) 4H20, C a = 15.4%. (2) 0.4950 gm of the salt precipitated in the hot, gave 0.2596 gm CaSO{. . ' . Ca'15.4570 Calc. for Ca(C4H406)4Hz0, Ca = 15.4% (3) I . 0650 gm of the salt went on losing weight, when kept a t a temperature of 170'. The loss was slow. After 5 hours heating the loss was found to be about 12%; after 1 5 hours heating it came to about 24%. Finally, it stood constant at 25.6% Calc. for Ca(C4H406) 4Hz0, C n = 15.4% and HzO= 2 7 . 6 % Calc. for Ca(C4H406)3 % HzO, Ca = I 5.9% and HzO= 2 5 . I % Evidently this new salt is 4 Ca(C4H40s)HzO, as later experiments show. (4) This partially anhydrous salt gets hydrated when moistened with water and dried in air. The gain is 25.4%. Calculated for Ca(C4H406) 1/4 H20 passing into Ca(C4H406) 3 H z 0 the gain is 25.1%
1024 0.2522
K . P. CHATTERJEE AND N . R . DHAR
gm of the supposed Ca (C4H406),3 H 2 0 gave 0.1406 gm Cas04
. ' . Ca=I6.4% Calc. for Ca(C4H40d3H20, C a = 16'.5% In Beilstein, we find only the tetrahydrate. Consequently a new hydrate Ca(CcH40e)3H2O is formed on hydration of 4Ca(C4H40s) H20.
Strontium Tartrate These salts are prepared in the same way as the calcium salts (I) 0.6818 gm of the salt precipitated at 3ooC, lost 0.1610 gm on strong heating . ' . H 2 0= 23.6% 0.7412 gm of the salt gaveo.4452 gm Sr804 . ' . Sr=28.7Te Calc. for Sr(C4H406) 4H20, Sr = 28.0% and H20= 23.6% I . 0936 gm of the salt precipitated from boiling solutions lost 0.0402 gm. (2) on heating to I 70" . ' .H 2 0= 3 . 7 % 0.4368 gm of the salt gave 0.3150 gm SrS04:.Sr=34.2% Cnlc.for Sr(C4Hdh) H20, Sr=34.5%andH20=7.og% and Sr(C&$h) 1/2 H20, Sr=35.8% and H2O=3.6% It appears that t)hesalt precipitated from the boiling solutions and dried in air has the formula Sr(C4H40s)H2O and this monohydrate, on heating to 170°C does not lose the whole of its water and most probably becomes converted into the hemihydrate Sr(C4H406)1 / 2 H20. Ferrous Tartrate It is very difficult to prepare this ealt in the cold. Being very soluble, it could not be thoroughly purified. 0.2486 gm of the salt precipitat,ed at the ordinary temperature, lost (I) nothing when heated to 170°C. HzO is nil. The above salt gave 0.0926 gm Fe203 . ' , Fe = 26.1% Calc. for Fe(C4H406),Fe = 27.47@and water is nil. (2) o.713A gm of the salt, precipitated when boiling solutions are mixed, lost nothing when heated to I 70°C and yielded 0.2646 gm Fez03. . ' . Fe = 25.9%
Calc. for Fe(C4H40e),Fe = 2 7 . 4Y0 and water is nil. It has already been observed that the salts could not be purified. It seems the same salt Fe(C4H406)is obtained in the hot and in the cold.
The Velocity of Precipitation and the Influence of the Temperature on it. From the foregoing pages, it will be seen that in most, cases the same hydrate is precipitated in the hot and in the cold. Also it will be seen that in almost all cases the solubility of these hydrates increases with increask of temperature. Consequently it is impossible to find out an explanation of t.he rapid increase of precipitation when a cold mixture of an oxa1at.e solution and a metallic salt solution, say of magnesium, be heated. We have already observed that when a concentrated solution of magnesium sulphate and oxalic acid be mixed a t the ordinary temperature, no precipitate is immediat,ely obtained. But, if the mixt.ure be heated, a copious precipitate of magnesium oxalate is obtained.
SPARINGLY SOLUBLE SALTS
1025
We have also proved that the hydrate obtained in the hot and the cold has the same formula, namely MgCz04 2H20, and that the solubility of this hydrate increases with increase of temperature. I n order to explain the peculiar behaviour of increased precipitation, on heating the reacting sohtions, we have to look for its explanation in the phenomena of the velocity of precipitation of these substances. We have attempted to measure tlhisvelocity of precipitation by adopting different methods. Method ( I ) is, mixing the reacting subst,ances at a definite temperature and taking out from time to time a measured volume of the unacted substance with the help of a pipette, to the nozzle of which was attached a rubber tubing, stuffed with cotton wool. Seminormal manganous sulphate of known strength was mixed with potassium oxalate of equivalent strength and left at a constant t,emperature. From time to time, a measured volume of the unchanged substance was taken out with the help of a pipette and the unchanged oxalate was estimated by titrating against permanganate. The following results were obtained. At 30" Time
cc N/IO K M n 0 4 for
o minutes 2
IO
cc solution
KO(Zero molecular)
72.8 45.7
. I022
4
Jf
21.55
1320
6 8
"
15.65
.I 1 2 0
"
10.9
mean
.1031 .1123
The solutions were diluted and mixed and in most cases the same velocity coefficient was obtained a t t,he same temperature.
It was very soon realised that it was very different to reproduce these results. Very slight shaking affects the results to a great extent. In order to find out the temperature coefficient of this velocity of precipitation, measurements were also made a t 4OC by mixing 50 cc seminormal potassium oxalate and 50 cc seminormal manganous sulphate both kept a t 4OC. The following results were obtained. At 4°C Time in minutes
KO
KMn04 for IOCC solution ( 1 6 8 . 8 ~KMnOa ~ = IOCC oxalate)
4
70.5 56.9
.039 ,042
2
6
47
.042
IO
42
.030
I2
I4
33 e 3 24.6
* 033 .038
16 18
16.04
.045
40
.os I
mean
.04
1026
K. P. CHATTERJEE AND N . R . DHAR
KI A(TI-TP) Applying the Arrhenius formula, log, - = 1 we get A = 1447 Kg TIXTP Kt+ TO and -= 1 . 5 2 Kt Attempts were made to find out the velocity coefficient whilst the reacting mixtures were being continually shaken. But no definite conclusion with regard to the order of the velocity of precipitation could be arrived at. In order to avoid supersaturation, a little manganous oxalate was added at the beginning of the experiment but, this did not help in getting concordant results calculated as zero, uni- or bimolecular formula. It was also obgerved that by mixing seminormal manganous sulphate and an equivalent quantity of potassium oxalate, the appearance of visible precipitation takes place in 40 seconds at 3ooC,whilst a t ooit takes 2 minutes, which gives Arrhenius A = 13 I 5 Kt+Io and - =1.48. Kt A similar experiment made with zinc sulphate and potassium oxalate gives the interval for visible appearance of precipitation as I minute and 5 I seconds Kt+ro at 29°C and 5 minutes 30 seconds a t oc which would give -= I . 5 (nearly).
Kt
Method ( 2 ) . The mixture of zinc sulphate'and potassium oxalate was kept in a beaker over a piece of paper ruled with black lines, and the time was noted when the black lines just disappeared, due to the turbidity of the mixture from precipitation of the solid oxalate. In this case the solution was kept constantly stirred. At 29°C the disappearance of the black lines took place in 136 seconds whilst a t 3" it took 660 seconds. These results give the Arrhenius Kt+Io - I .90 In the former cases, the rea,cting mixconstant A = 2204 and -Nt ture was not stirred at all, while in the latter case it was continually shaken. It seems therefore that the temperature accelerates the velocity of precipitation to a greater extent, when the reacting substances are shaken than when the velocity is determined without shaking the mixture. It is of interest to note that Marc' showed that the velocity of crystallisation of supersaturated solutions of substances like potassium sulphate follows tQe bimolecular formula, whilst we have got agreement according to zero molecular formula2. The value of the temperature coefficient for IO" rise obtained by Marc is I . 6 whilst our value is I . 5 . Moreover Jablczynski3 obtains the value 2 for the temperature coefficient of the velocity of formation of silver bromide and silver chloride. Method (3) We also attempted to obtain the velocity of precipitation from measurements of conductivities. We mixed solutions of a salt (e.g. manganous sulphate or zinc sulphate) and an oxalate and measured the resistance of the mixture. ,4t the very outset we were surprised to find that the resistance changes very slightly though an appreciable amount of the oxalate Z. physik. Chem. 61, 385 (1908). Compare Dhar: Z. anorg. Chem. 121, 156 (1922); J. Chem. SOC.111, 727 (1917). 3 Z. physik. Chem. 82, 115 (1913).
SPARINGLY SOLUBLE SALTS
1027
crystallised out. In order to find out an explanation, we had to make some preliminary experiments. We mixed 50 cc of N/so NaCl and N/so AgN03. We determined the resistance 45 seconds after the mixing. The value of this resistance did not change with time. It appears, therefore, that the precipitation has probably been finished within this 45 seconds. If we mix together, say, potassium oxalate and manganous sulphate we get manganous oxalate and potassium sulphate. As most of the manganePe oxalate goes out as precipitate, the conductivity of the mixture should decrease appreciably. But as a matter of fact, in most rases we got n, slight decrease of conductivity. We thought that the explanation might have something to do with the phenomenon of supersaturation. Incidentally we studied the molecular conductivity of supersaturated solutions of sodium metate, calcium chloride, etc. In the case of sodium acetate and in the case of calcium chloride we found that in concentrated solutions, the curve of molecular conductivity with dilution shows definite breaks. In other words, at certain concentration instead of increasing with dilution, the molecular conductivity slightly decreases with dilution. In a foregoing paper, it has been observed that breaks of that nature which are so common in nonaqueous solutions are expected even in aqueous solutionr of electrolytes. ifheterminntion could be made with sufficiently concentrated solutions.'
Summary (I) Citrates of calcium and stronium, oxalates and tartrates of copper, zinc, magnesium, manganese, cobalt, nickel, iron (ferrous) and thorium, are precipitated more readily from boiling solutions of a citrate or an oxalate or a tartrate and a salt solution of the respective metal than whkn the reacting substances are mixed in the cold. The phenomenon of increased precipitation on heating the solutions of the reacting substances is of common occurrence. The following salts have been prepared and described:(2) (a) Cas(CsHs07)z 16HzO; Cna(CeH607)z 6HzO; Cas(C~H507)z4Hz0 and Cas(CsHsO./)z1/2 HzO. (b) SrdC6H607~~ 5HzO; Sr3(CsH607)z HzO. (c) 3NiC~04SHzO; 3NiC204:HzO (d) 4C0Cz01 9HzO; CoC2042Hz0 (e) CiiC2O4 (f) MnC204 2HzO (g) MgCz04 2Hz0 (h) FeC204 2Hz0 (i) ZnCzOc 2Hz0; (j) T h G 0 4 ) ~ H z O (k) Mg(C4H406) 244 HzO; Mg(CdH4Os) (1) Mn(C4H4Od 2H20; Mn(C4H406) (m) Co(C4H406) 2 ?4 H2O; Co(C4H40a) (n) Ni(C4H406) 3Hz0; I\Ti(C4H406)2 $6 HzO; Ni(C4H400) Compare Dhar: Z. Elektrochem. 20, 60 (1914).
1028
K. P. CHATTERJEE AND N. R. DHAR
(0) Ca(C.&06) 4H-20; 4Ca(C4H406)HzO; Ca(C4H406)3H-20. (p). Sr(C4H406) 4Hz0; Sr(C4H406)HzO; Sr(C4H406)1/2 HzO (9) Fe(CdH406)
(3) Molecular volume of water of crystallisation in some of these sparingly soluble salts, has the value 15 or 16,-practically the same as obtained with the readily soluble salts containing water of crystallisation,
(4) The velocity of precipitation of some oxalates and the influence of temperature on it, have been investigated. It is observed that when the reacting mixture is not agitated, the velocity of the reaction apparently follows the Kt+ro zero-molecular formula (KO= ) and - = I , 5 t Kt ( 5 ) Incidentally it has been observed that the molecular conductivities of very concentrated solutions of calcium chloride, sodium acetate etc. do not go on increasing with dilution, but at certain concentrations slightly fall off. Chemical Laboratow Allahabad, India. " April 16, 1924.
