The solubility of hydroxides

Water is one of the products formed via mechanochemical interaction between hydroxides. Its presence can substantially effect on the mechanism of transport processes and, thus, on the kinetics of mechanochemical processes, since some reaction products (initial or final) can be dissolved in it. Hence, the reaction could proceed with the participation of the dissolved forms of components. 

The major part of hydroxides are poorly soluble or insoluble in water, except well-soluble hydroxides of alkaline metals and thallium, and much worse soluble hydroxides of alkaline earths. 

The analysis of the solubility of compounds is usually made on the basis of the solubility product (L), which, in fact, is the equilibrium constant of the process 

The equilibrium constant is connected with free energy of dissolution AG by the equation 

The state of ions in an infinitely diluted solution (ion activity product) is selected to be the standard state, and the activity of AxBy(soM) is taken to be unity. 

---

The solubility of hydroxides of the trivalent metals is exceedingly small as a consequence the hydroxides of the trivalent elements are precipitated by ammonium hydroxide in the presence of ammonium salts. 

In the technology of water the solubilities of hydroxides and carbonates are of the greatest importance, these are infuenced by metal ion hydrolysis. With metal cations water forms so-caUed aqua complexes. Usually, four or six molecules of water are coordinated around the central atom (e.g. [Fe(H20)e]"+, [A1(H20)6] +, [Mn(H20)e] +, etc.). 

The quantity of dissolved Mn(II) in natural waters depends on the solubility of hydroxide, carbonate and sulphide. In most natural waters the equilibrium concentration of dissolved manganese is determined by the MnC03 solubility. The production of Mn(OH)2 is considered only in a stronger alkaline medium. In the presence of hydrogen sulphide and its ionic forms the Mn(II) solubility in alkaline media is limited by the MnS solubility and the equilibrium concentrations are the lowest under these conditions. 

The authors discuss systematic trends and specific phenomena on the solubility of hydroxides and oxides of tri-, tetra-, penta- and hexavalent actinides. The reported solubility constants of amorphous Th02(am, hyd), microcrystalline and crystalline Th02(cr), the conclusions on the solubility controlling solid phase and the solubility increasing effect of eigencolloids are taken from [2001NEC/KIM], [2002NEC/MUL],... 

The solubility of a substance can change considerably in response to a number of factors. For example, the solubilities of hydroxide salts, like Mg(OH)2, are dependent upon the pH of the solution. The solubility is also affected by concentrations of other ions in solution, especially common ions. In other words, the numeric value of the solubility of a given solute can and does change as the other species in solution change. In contrast, the solubility-product constant, fQp, has only one value for a given solute at any specific temperature. FIGURE 17.16 summarizes the relationships among various expressions of solubility and K. ... 

The solubility of hydroxides (or acids) of stabilizers elements in water can be an indication, of their effect on the C2S hydraulic activity. Much better solubility of these elements than that of calcimn hydroxide or silicic acid improves the C2S reactivity in respect to water [151]. Ba instead of Ca and A1 instead of Si are the examples. Boron, which is known as a retarder, lowers the C2S hydration rate. 

The solubility of hydroxides can be also expressed in terms of the dependence of the total concentration of the soluble species of metal ion ]M ] on pH (Figure 5) ... 

The pH of the electrolyte does not only have an effect on the passivation potential, but also on the passivation current density, because both the metal dissolution kinetics and the solubility of hydroxides depend on pH. Figure 6.16 shows that the passivation current density of iron becomes smaller at higher pH. This has been explained by a lowering of the solubility of ferrous hydroxide, which precipitates at the surface. Since both the passivation potential and the passivation current density decrease with increasing pH, spontaneous passivation of iron becomes possible in basic, aerated media. This explains why steel reinforcements in concrete (pH >13) resist corrosion well as long as chemical reactions with carbon dioxide from air (carbonation of concrete) do not modify the alkalinity. 

Aqueous ammonia can also behave as a weak base giving hydroxide ions in solution. However, addition of aqueous ammonia to a solution of a cation which normally forms an insoluble hydroxide may not always precipitate the latter, because (a) the ammonia may form a complex ammine with the cation and (b) because the concentration of hydroxide ions available in aqueous ammonia may be insufficient to exceed the solubility product of the cation hydroxide. Effects (a) and (b) may operate simultaneously. The hydroxyl ion concentration of aqueous ammonia can be further reduced by the addition of ammonium chloride hence this mixture can be used to precipitate the hydroxides of, for example, aluminium and chrom-ium(III) but not nickel(II) or cobalt(II). 

Originally, general methods of separation were based on small differences in the solubilities of their salts, for examples the nitrates, and a laborious series of fractional crystallisations had to be carried out to obtain the pure salts. In a few cases, individual lanthanides could be separated because they yielded oxidation states other than three. Thus the commonest lanthanide, cerium, exhibits oxidation states of h-3 and -t-4 hence oxidation of a mixture of lanthanide salts in alkaline solution with chlorine yields the soluble chlorates(I) of all the -1-3 lanthanides (which are not oxidised) but gives a precipitate of cerium(IV) hydroxide, Ce(OH)4, since this is too weak a base to form a chlorate(I). In some cases also, preferential reduction to the metal by sodium amalgam could be used to separate out individual lanthanides. 

Divide the saturated solution of n-butyl alcohol in water into three approximately equal parts. Treat these respectively with about 2-5 g. of sodium chloride, potassium carbonate and sodium hydroxide, and shake each until the soli have dissolved. Observe the effect of these compounds upon the solubility of n-butanol in water. These results illustrate the phenomenon of salting out of organic compounds, t.e., the decrease of solubility of organic compounds in water when the solution is saturated with an inorganic compound. The alcohol layer which separates is actually a saturated solution of water in n-butyl alcohol. 

Study of the solubility behaviour of the compound. A semi-quantitative study of the solubility of the substance in a hmited number of solvents (water, ether, dilute sodium hydroxide solution, dilute hydrochloric acid, sodium bicarbonate solution, concentrated sulphuric and phosphoric acid) will, if intelligently apphed, provide valuable information as to the presence or absence of certain classes of organic compounds. 

Most metals will precipitate as the hydroxide in the presence of concentrated NaOH. Metals forming amphoteric hydroxides, however, remain soluble in concentrated NaOH due to the formation of higher-order hydroxo-complexes. For example, Zn and AP will not precipitate in concentrated NaOH due to the formation of Zn(OH)3 and Al(OH)4. The solubility of AP in concentrated NaOH is used to isolate aluminum from impure bauxite, an ore of AI2O3. The ore is powdered and placed in a solution of concentrated NaOH where the AI2O3 dissolves to form A1(0H)4T Other oxides that may be present in the ore, such as Fe203 and Si02, remain insoluble. After filtering, the filtrate is acidified to recover the aluminum as a precipitate of Al(OH)3. 

Another important parameter that may affect a precipitate s solubility is the pH of the solution in which the precipitate forms. For example, hydroxide precipitates, such as Fe(OH)3, are more soluble at lower pH levels at which the concentration of OH is small. The effect of pH on solubility is not limited to hydroxide precipitates, but also affects precipitates containing basic or acidic ions. The solubility of Ca3(P04)2 is pH-dependent because phosphate is a weak base. The following four reactions, therefore, govern the solubility of Ca3(P04)2. 

The reaction time depends on the quality of the potassium hydroxide employed. An induction period is often observed when older potassium hydroxide samples are used, possibly because surface formation of carbonates reduces the solubility of the salt in acetonitrile. An attempt was made to monitor the cinnamonitrile reaction by GLC, following loss of starting... 

The amine (Imol) is added to a solution of anhydrous zinc chloride (Imol) in concentrated hydrochloric acid (42mL) in ethanol (200mL, or less depending on the solubility of the double salt). The solution is stirred for Ih and the precipitated salt is filtered off and recrystallised from ethanol. The free base is recovered by adding excess of 5-ION NaOH (to dissolve the zinc hydroxide that separates) and is steam distilled. Mercuric chloride in hot water can be used instead of zinc chloride and the salt is crystallised from 1% hydrochloric acid. Other double salts have been used, e.g. cuprous salts, but are not as convenient as the above salts. 

Solubility Product — The solubility product constant commonly referred to as the solubility product provides a convenient method of predicting the solubility of a material in water at equilibrium. Copper hydroxide, for example, dissolves according to the following equilibrium ... 

MSH + MOH). Accordingly, solubilities depend sensitively not only on temperature but also on pH and partial pressure of H2S. Thus, by varying the acidity. As can be separated from Pb, Pb from Zn, Zn from Ni, and Mn from Mg. In pure water the solubility of Na2S is said to be 18.06g per 100 g H2O and for Ba2S it is 7.28 g. In the case of some less-basic elements (e.g. AI2S3, Cr2S3) hydrolysis is complete and action of H2S on solutions of the metal cation results in the precipitation of the hydroxide likewise these sulfides (and SiS2, etc.) react rapidly with water with evolution of H2S. 

The solubility of barium hydroxide in water at 20°C is 1.85 g/100 g water. A solution is made up of256 mg in 35.0 g of water. Is the solution saturated If not, how much more barium hydroxide needs to be added to make a saturated solution ... 

Ammonium chloride solutions are slightly acidic, so they are better solvents than water for insoluble hydroxides such as Mg(OH)2. Find the solubility of Mg(OH)2 in moles per liter in 0.2 M NH4CI and compare with the solubility in water. Hint Find K for the reaction... 

The similarities among the hydroxides are obvious. Let s compare sodium carbonate and ammonia. Sodium carbonate, Na2C03, dissolves in water to give a solution with the properties that identify a base. Quantitative studies of the solubilities of carbonates show that carbonate ion, C03-2, can react with water. The reactions are... 

Exercise 21-10 demonstrates that there is a regular trend in the solubilities of the alkaline earth hydroxides. 

Although the hydroxides of the alkaline earth elements become more soluble in water as we go down the column, the opposite trend is observed in the solubilities of the sulfates and carbonates. For example, Table 21-VII shows the solubility products of the alkaline earth sulfates. 

Aluminum sulfate, A12(S04)3 H20, is the commonest alum used. Hydration is typically 14 to 16 H20. It hydrolyzes and polymerizes in water and typically is used within a narrow window of pH levels of 5.5 to 6.5 to minimize the solubility of aluminum in the treated water. If alkalinity is present (say, due to calcium bicarbonate), the following reaction occurs, producing insoluble aluminum hydroxide [Al(OH)3]. 

We can use Le Chatelier s principle as a guide. This principle tells us that, if we add a second salt or an acid that supplies one of the same ions—a common ion —to a saturated solution of a salt, then the equilibrium will tend to adjust by decreasing the concentration of the added ions (Fig. 11.15). That is, the solubility of the original salt is decreased, and it precipitates. We can conclude that the addition of excess OH- ions to the water supply should precipitate more of the heavy metal ions as their hydroxides. In other words, the addition of OH ions reduces the solubility of the heavy metal hydroxide. The decrease in solubility caused by the addition of a common ion is called the common-ion effect. 

The solubility of a solid can be increased by removing one of its ions from solution acid can be used to dissolve a hydroxide, sulfide, sulfite, or carbonate precipitate and nitric acid can be used to oxidize metal sulfides to sulfur and a soluble salt. 

Use the data in Table 11.4 to calculate the solubility of each of the following sparingly soluble substances in its respective solution aluminum hydroxide at (a) pH = 7.0 ... 

---