Aqueous solutions Hydration Solubility

Cu+ makes the hydrated ion the more stable in aqueous solution. Thus soluble copper(I) compounds always disproportionate ... 

Properties (L(-)-phenylalanine) Plates and leaflets from concentrated aqueous solutions, hydrated needles from dilute aqueous solutions, decomposes at 283C. Soluble in water slightly soluble in methanol and ethanol. (D(+)-phenylalamne ) Leaflets from water, decomposes 285C. Soluble in water slightly soluble in methanol. (DL-phenylalanine) Leaflets or prisms from water or alcohol sweet tasting. Decomposes 318-320C. Soluble in water. 

Acetaldehyde, CH3CHO, b.p. 21°, is generally used in aqueous solution, which has also a characteristic odour paraldehyde, (CH3CHO)3, is a liquid polymer, b.p. 124°, slightly soluble in water, odour similar to that of acetaldehyde, but less intense. Chloral, CCI3CHO, a liquid, is almost invariably encountered as the stable solid hydrate , CCl3CH(OH)2, m.p. 57 . Both have a characteristic odour the hydrate is readily soluble in water. 

The distinction between pairwise and bulk hydrophobic interactions is often made, although some authors doubt the existence of an intrinsic difference between the two ". Pairwise hydrophobic interactions denote the interactions behveen two isolated nonpolar solutes in aqueous solution. They occur in the regime where no aggregation takes place, hence below the critical aggregation concentration or solubility limit of the particular solute. If any breakdown of the hydrophobic hydration shell occurs, it will be only transient. 

X 10 J/T (5.71 //g) at room temperature. It is air stable at 25°C, but is slowly converted to Fe202 and bromine at 310°C. The light yellow to brown hydroscopic sohd is soluble ia water, alcohol, ether, and acetonitrile. Iron(II) bromide forms adducts with a wide range of donor molecules. Pale green nona-, hexa-, tetra-, and dihydrate species can be crystallized from aqueous solutions at different temperatures. A hydrate of variable water content,... 

The chlorides, bromides, nitrates, bromates, and perchlorate salts ate soluble in water and, when the aqueous solutions evaporate, precipitate as hydrated crystalline salts. The acetates, iodates, and iodides ate somewhat less soluble. The sulfates ate sparingly soluble and ate unique in that they have a negative solubitity trend with increasing temperature. The oxides, sulfides, fluorides, carbonates, oxalates, and phosphates ate insoluble in water. The oxalate, which is important in the recovery of lanthanides from solutions, can be calcined directly to the oxide. This procedure is used both in analytical and industrial apptications. 

The salt is extremely soluble ia water (Table 4), crystallising from aqueous solution as the hydrates LiBr H20 [23303-71-17, LiBr 2H20 [13453-70-8] and LiBr 3H2O [76082-04-7]. The anhydrous salt is obtained by dryiag under vacuum at elevated temperatures. 

Many other metal thiosulfates, eg, magnesium thiosulfate [10124-53-5] and its hexahydrate [13446-30-5] have been prepared on a laboratory scale, but with the exception of the calcium, barium [35112-53-9] and lead compounds, these are of Httle commercial or technical interest. Although thaHous [13453-46-8] silver, lead, and barium thiosulfates are only slightly soluble, other metal thiosulfates are usually soluble in water. The lead and silver salts are anhydrous the others usually form more than one hydrate. Aqueous solutions are stable at low temperatures and in the absence of air. The chemical properties are those of thiosulfates and the respective cation. 

Oxo Ion Salts. Salts of 0x0 ions, eg, nitrate, sulfate, perchlorate, hydroxide, iodate, phosphate, and oxalate, are readily obtained from aqueous solution. Thorium nitrate is readily formed by dissolution of thorium hydroxide in nitric acid from which, depending on the pH of solution, crystalline Th(N02)4 5H20 [33088-17 ] or Th(N02)4 4H20 [33088-16-3] can be obtained (23). Thorium nitrate is very soluble in water and in a host of oxygen-containing organic solvents, including alcohols, ethers, esters, and ketones. Hydrated thorium sulfate, Th(S0 2 H20, where n = 9, 8, 6, or 4, is... 

Zinc chloride melts at 275°C, bods at 720°C, and is stable in the vapor phase up to 900°C. It is very hygroscopic, extremely water-soluble, and soluble in organic Hquids such as alcohols, esters, ketones, ethers, amides, and nitrides. Hydrates with 1, 1.5, 2.5, 3, and 4 molecules of water have been identified and great care must be exercised to avoid hydration of the anhydrous form. Aqueous solutions of zinc chloride are acidic (pH = 1.0 for 6 M) and, when partially neutralized, can form slightly soluble basic chlorides, eg, ZnCl2 4Zn(OH)2 [11073-22-6] and Zn(OH)Cl [14031-59-5]. Many other basic chlorides have been reported (58). 

In general, hydrated borates of heavy metals ate prepared by mixing aqueous solutions or suspensions of the metal oxides, sulfates, or halides and boric acid or alkali metal borates such as borax. The precipitates formed from basic solutions are often sparingly-soluble amorphous soHds having variable compositions. Crystalline products are generally obtained from slightly acidic solutions. 

The dichromate(VI) salts may be obtained by the addition of acid to the chromate(VI) salts. However, they are better prepared by adding one-half the acid equivalent of a metal hydrate, oxide, or carbonate to an aqueous solution of CrO, then removing the water and/or CO2. Most dichromates(VI) are water-soluble, and the salts contain water(s) of hydration. However, the normal salts of K, Cs, and Rb are anhydrous. Dichromate(VI) compounds of the colorless cations are generally orange-red. The geometry of Ci2 is described as two tetrahedral CrO linked by the shared odd oxygen (72). 

Ethyleneamines are soluble in water. However, in concentrated aqueous solutions, amine hydrates may form the reaction is mildly exothermic. The hydrates of linear TETA and PIP melt around 50°C (65,66). 

The acid is precipitated from aqueous solution as the mono-hydrate, which is soluble in cold dilute hydrochloric acid to the extent of about 6 g. per 1. 

Gaseous SO2 is readily soluble in water (3927 cm SO2 in lOOg H2O at 20°). Numerous species are present in this aqueous. solution of sulfurous acid" (p. 717). At 0° a cubic clathrate hydrate also forms with a composition S02.6H20 it.s dissociation pressure reaches I atm at 7.1°. The ideal composition would be SO2.55H2O (p. 627). 

The data given in Tables 1.9 and 1.10 have been based on the assumption that metal cations are the sole species formed, but at higher pH values oxides, hydrated oxides or hydroxides may be formed, and the relevant half reactions will be of the form shown in equations 2(a) and 2(b) (Table 1.7). In these circumstances the a + will be governed by the solubility product of the solid compound and the pH of the solution. At higher pH values the solid compound may become unstable with respect to metal anions (equations 3(a) and 3(b), Table 1.7), and metals like aluminium, zinc, tin and lead, which form amphoteric oxides, corrode in alkaline solutions. It is evident, therefore, that the equilibrium between a metal and an aqueous solution is far more complex than that illustrated in Tables 1.9 and 1.10. Nevertheless, as will be discussed subsequently, a similar thermodynamic approach is possible. 

Similarly, concepts of solvation must be employed in the measurement of equilibrium quantities to explain some anomalies, primarily the salting-out effect. Addition of an electrolyte to an aqueous solution of a non-electrolyte results in transfer of part of the water to the hydration sheath of the ion, decreasing the amount of free solvent, and the solubility of the nonelectrolyte decreases. This effect depends, however, on the electrolyte selected. In addition, the activity coefficient values (obtained, for example, by measuring the freezing point) can indicate the magnitude of hydration numbers. Exchange of the open structure of pure water for the more compact structure of the hydration sheath is the cause of lower compressibility of the electrolyte solution compared to pure water and of lower apparent volumes of the ions in solution in comparison with their effective volumes in the crystals. Again, this method yields the overall hydration number. 

CO3 species was formed and the X-ray structure solved. It is thought that the carbonate species forms on reaction with water, which was problematic in the selected strategy, as water was produced in the formation of the dialkyl carbonates. Other problems included compound solubility and the stability of the monoalkyl carbonate complex. Van Eldik and co-workers also carried out a detailed kinetic study of the hydration of carbon dioxide and the dehydration of bicarbonate both in the presence and absence of the zinc complex of 1,5,9-triazacyclododecane (12[ane]N3). The zinc hydroxo form is shown to catalyze the hydration reaction and only the aquo complex catalyzes the dehydration of bicarbonate. Kinetic data including second order rate constants were discussed in reference to other model systems and the enzyme carbonic anhy-drase.459 The zinc complex of the tetraamine 1,4,7,10-tetraazacyclododecane (cyclen) was also studied as a catalyst for these reactions in aqueous solution and comparison of activity suggests formation of a bidentate bicarbonate intermediate inhibits the catalytic activity. Van Eldik concludes that a unidentate bicarbonate intermediate is most likely to the active species in the enzyme carbonic anhydrase.460... 

Weissenbom PK, Pugh RJ (1996) Surface tension of aqueous solutions of electrolytes relationship with hydration, oxygen solubility, and bubble coalescence. J Colloid Interface Sci 184 550-553... 

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