Ammonia solubihty

Ammonium lauryl sulfate [2235-54-3] is available as an approximately 28% active, clear, Hquid form. It has greater solubiHty than the sodium salt and is more likely to be used in formulating clear shampoos. Systems using this detergent show their best stabiHty at pH between 6 and 7. Lower pH would tend to hydrolyze the detergent, eventually releasing ammonia. 

The covalent character of mercury compounds and the corresponding abiUty to complex with various organic compounds explains the unusually wide solubihty characteristics. Mercury compounds are soluble in alcohols, ethyl ether, benzene, and other organic solvents. Moreover, small amounts of chemicals such as amines, ammonia (qv), and ammonium acetate can have a profound solubilizing effect (see COORDINATION COMPOUNDS). The solubihty of mercury and a wide variety of mercury salts and complexes in water and aqueous electrolyte solutions has been well outlined (5). 

Ammonia forms a great variety of addition or coordination compounds (qv), also called ammoniates, ia analogy with hydrates. Thus CaCl2 bNH and CuSO TNH are comparable to CaCl2 6H20 and CuSO 4H20, respectively, and, when regarded as coordination compounds, are called ammines and written as complexes, eg, [Cu(NH2)4]S04. The solubiHty ia water of such compounds is often quite different from the solubiHty of the parent salts. For example, silver chloride, AgQ., is almost iasoluble ia water, whereas [Ag(NH2)2]Cl is readily soluble. Thus silver chloride dissolves ia aqueous ammonia. Similar reactions take place with other water iasoluble silver and copper salts. Many ammines can be obtained ia a crystalline form, particularly those of cobalt, chromium, and platinum. 

Sodium is soluble in ethylenediamine (16,17), but solubiHty in other amines such as methyl- or ethylamine may require the presence of ammonia. Sodium solubiHty in ammonia and ethylenediamine solutions has been extensively investigated (18). Sodium is insoluble in most hydrocarbons and is... 

Barium nitrite [13465-94-6] Ba(N02)2, crystallines from aqueous solution as barium nitrite monohydrate [7787-38-4], Ba(N02)2 H2O, which has yellowish hexagonal crystals, sp gr 3.173, solubihty 54.8 g Ba(NO2)2/100 g H2O at 0°C, 319 g at 100°C. The monohydrate loses its water of crystallization at 116°C. Anhydrous barium nitrite, sp gr 3.234, melts at 267°C and decomposes at 270 °C into BaO, NO, and N2. Barium nitrite may be prepared by crystallization from a solution of equivalent quantities of barium chloride and sodium nitrite, by thermal decomposition of barium nitrate in an atmosphere of NO, or by treating barium hydroxide or barium carbonate with the gaseous oxidiation products of ammonia. It has been used in diazotization reactions. 

Potassium and ammonium dichromates are generally made from sodium dichromate by a crystallization process involving equivalent amounts of potassium chloride or ammonium sulfate. In each case the solubiHty relationships are favorable so that the desired dichromate can be separated on cooling, whereas the sodium chloride or sulfate crystallizes out on boiling. For certain uses, ammonium dichromate, which is low in alkaH salts, is required. This special salt may be prepared by the addition of ammonia to an aqueous solution of chromic acid. Ammonium dichromate must be dried with care, because decomposition starts at 185°C and becomes violent and self-sustaining at slightly higher temperatures. 

Hydrolysis of a nitrile to an acid. Reflux 1 g. of the nitrile with 6 ml. of 30-40 per cent, sodium hydroxide solution until ammonia ceases to be evolved (2-3 hours). Dilute with 6 ml. of water and add, with cooling, 7 ml. of 50 per cent, sulphuric acid. Isolate the acid by ether extraction, and examine its solubihty and other properties. 

Ammonia possesses similar physical properties to that of water, which is similarly highly associated. It is a good solvent for many compounds. Owing to the lower dielectric constant (NH3 16.9, H2O 78.3 at 298 K) of ammonia in comparison with water, less polar compounds are more soluble in ammonia and polar compounds, for example, salts, are more soluble in water. Organic compounds tend to have a higher solubihty in ammonia than in water. Armnonium salts, nitrates, nitrites, cyanides, and thiocyanates dissolve readily in ammonia. The solubihty increases from fluorides to chlorides, bromides, and iodides. Salts with higher charged ions dissolve only poorly in ammonia. This results in the reversal of some precipitation reactions in ammonia compared to water. 

Silver Chloride. Silver chloride, AgCl, is a white precipitate that forms when chloride ion is added to a silver nitrate solution. The order of solubihty of the three silver halides is Cl > Br > I . Because of the formation of complexes, silver chloride is soluble in solutions containing excess chloride and in solutions of cyanide, thiosulfate, and ammonia. Silver chloride is insoluble in nitric and dilute sulfuric acid. Treatment with concentrated sulfuric acid gives silver sulfate. 

To summarize the chemistry which can actually be expected for element 113, one may say that it is in general expected to fall between that of T1+ and Ag+. 113+ is expected to bind anions more readily than T1+ so that (113)C1 will be rather soluble in excess HCl whereas the solubihty of TlCl is essentially imchanged. Similarly, (113)C1 is expected to be soluble in ammonia water in contrast to the behavior of TlCl. The behavior of the 113+ ion should tend toward Ag+ in these respects. Also, although Tl(OH) is soluble and a strong base, the 113+ ion should form a sUghtly soluble oxide that is soluble in aqueous ammonia. 

The chemistry of the 115+ ion can be summarized as follows the complexing ability of 115+ can be expected to be low with such anions as the hahdes, cyanide and ammonia. Hydrolysis should occur readily for 115 in the oxidation state of 1, and the hydroxide, carbonate, oxalate and fluoride should be soluble. The sulfide should be insoluble and the chloride, bromide, iodine, and thiocyanide only slightly soluble. For example, excess HCl will not appreciably affect the solubihty of (115)C1. 

Inorganic heavy metals are usually removed from aqueous waste streams by chemical precipitation in various forms (carbonates, hydroxides, sulfide) at different pH values. The solubiHty curves for various metal hydroxides, when they are present alone, are shown in Figure 7. The presence of other metals and complexing agents (ammonia, citric acid, EDTA, etc) strongly affects these solubiHty curves and requires careful evaluation to determine the residual concentration values after treatment (see Table 9) (38,39). 

Most gases obey Henry s law, but there are some important exceptions. For example, if the dissolved gas reacts with water, higher solubilities can result. The solubihty of ammonia is much higher than expected because of the reaction... 

The aqueous solubihty of ammonia decreases rapidly as temperature increases. The existence of undissociated ammonium hydroxide [1336-21 -6], NH OH,... 

There are a considerable number of stable crystalline salts of the ammonium ion [14798-03-9] NH. Several are of commercial importance because of large scale consumption in fertilizer and industrial markets. The ammonium ion is about the same size as the potassium and mbidium ions, so these salts are often isomorphous and have similar solubiHty in water. Compounds in which the ammonium ion is combined with a laige, uninegative anion are usually the most stable. Ammonium salts containing a small, highly charged anion generally dissociate easily into ammonia (qv) and the free acid (1). At about 300°C most simple ammonium salts volatilize with dissociation, for example... 

The solubility of nitrogen in B4C was measured by Oscroft and Thompson (1991) [256]. Fast-flowing ammonia (12 h) was used in the temperature range from 1073 to 1373 K. X-ray diffraction, chemical analysis (SEM/EDX), combustion analysis were used and a solubihty of 3.8 mass% was found. However, it can be assumed that the use of ammonia with high nitrogen activities may have resulted in a metastable solid solution B CyNj. 

When sodium chloride dissolves in water, the ions are stabilized in solution by hydration, which involves ion-dipole interaction. In general, we predict that ionic compounds should be much more soluble in polar solvents, such as water, liquid ammonia, and liquid hydrogen fluoride, than in nonpolar solvents, such as benzene and carbon tetrachloride. Because the molecules of nonpolar solvents lack a dipole moment, they cannot eflfeclively solvate the Na and Q ions. (Solvation is the process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner. The process is called hydration when the solvent is water.) The predominant intermolecular interaction between ions and nonpolar compounds is ion-induced dipole interaction, which is much weaker than ion-dipole interaction. Consequently, ionic compounds usually have extremely low solubihty in nonpolar solvents. 

The above calculation assumes that the liquid makes only a single pass through the tower. In practice, a more effective way of treating the gas would be to recirculate the hquid with a small feed of concentrated acid to the recirculation loop and a small bleed of the recirculation loop contents. This would permit the NH4 concentration to build up in the recirculation loop. However, even if this concentration built up to the solubihty limit of ammonium nitrate (NH4 NO3), about 30 (g mol)/L, the partial pressure exerted by the dissolved ammonia at the top of the tower would still be nearly two orders of magnitude below the partial pressure of ammonia in the gas at the top of the tower. Therefore, for all practical conditions, the reaction is effectively irreversible, given the acid concentrations calculated in step 4. 

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