Hydroxides, solubility effects

The OH concentration increases (decreases) by one order of magnitude for every unit increase (decrease) in pH. This means that the formation of a metal hydroxide (whether as a colloid or as a precipitate) in aqueous solution will be strongly dependent on temperature when the product of the free metal ions and OH ions is close to the hydroxide solubility product, although increase in Ksp with temperature may partially offset this effect. 

Commonly, different metals exhibit different solution pH of zero net charge. For this reason, different metals exhibit minimum solubility at different pH values, which makes it difficult to precipitate effectively two or more metals, as metal-hydroxides, simultaneously. Thus metal-hydroxide solubility as a function of pH displays a U-shaped behavior. The lowest point in the U-shaped figure signifies the solution pH of zero net charge and is demonstrated below. Consider the solid Fe(OH)2s,... 

Since pH = 14 - pOH" (where pOH" denotes the negative log of OH"), the pH of minimum solubility for Fe(OH)2s would be 11.21. The example above is only for demonstration purposes since only two of the many potentially forming Fe2+-hydroxy species were employed. A graphical representation of the solubility of Fe(OH)2s (Eq. 2.47) and Fe(OH)3s as a function of pH are shown in Figure 2.7. The data in Figure 2.9 show the solubility of various heavy metals as a function of pH, whereas the data in Figure 2.10 show the decrease in metal-hydroxide solubility as pH increases (common ion effect). They do not, however, show the expected increase in metal-hydroxide solubility as pH increases. 

The solubility of TPP is much greater in benzene than in mineral oil, and It is therefore likely that its average location (10. 11)is nearer to the Interface and the copper does not have to be transported (e.g., as a complex) into the droplet Interior. Since the microdroplet has a net negative surface charge, it is expected that the local concentration of hydroxide Is lower, and that hydroxide cannot effectively penetrate very deeply into the surface region. This is consistent with the effect of hydroxide on an alkylation reaction, to be discussed below. This can account for its failure to Increase the rate of the base removal component, but Its role In promoting the dependence of k on copper ion remains unexplained. 

Not all solubility curves are smooth, as can be seen in Figure 3.1b. A discontinuity in the solubility curve denotes a phase change. For example, the solid phase deposited from an aqueous solution of sodium sulphate below 32.4 °C will consist of the decahydrate, whereas the solid deposited above this temperature will consist of the anhydrous salt. The solubility of anhydrous sodium sulphate decreases with an increase in temperature. This negative solubility effect, or inverted solubility as it is sometimes called, is also exhibited by substances such as calcium sulphate (gypsum), calcium, barium and strontium acetates, calcium hydroxide, etc. These substances can cause trouble in certain types of crystallizer by causing a deposition of scale on heat-transfer surfaces. 

Gnanaprakash G, Philip RB (2007) Effect of divalent metal hydroxide solubility product on the size of ferrite nanoparticles. Mater Lett 61 4545-4548... 

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). 

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. 

It is frequently advisable in the routine examination of an ester, and before any derivatives are considered, to determine the saponification equivalent of the ester. In order to ensure that complete hydrolysis takes place in a comparatively short time, the quantitative saponi fication is conducted with a standardised alcoholic solution of caustic alkali—preferably potassium hydroxide since the potassium salts of organic acids are usuaUy more soluble than the sodium salts. A knowledge of the b.p. and the saponification equivalent of the unknown ester would provide the basis for a fairly accurate approximation of the size of the ester molecule. It must, however, be borne in mind that certain structures may effect the values of the equivalent thus aliphatic halo genated esters may consume alkali because of hydrolysis of part of the halogen during the determination, nitro esters may be reduced by the alkaline hydrolysis medium, etc. 

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. 

In petroleum and oxygenate finish removers, the major ingredient is normally acetone, methyl ethyl ketone [78-93-3], or toluene. Cosolvents include methanol, / -butanol [71-36-3], j -butyl alcohol [78-92-2], or xylene [1330-20-7]. Sodium hydroxide or amines are used to activate the remover. Paraffin wax is used as an evaporation retarder though its effectiveness is limited because it is highly soluble in the petroleum solvents. CeUulose thickeners are sometimes added to liquid formulas to assist in pulling the paraffin wax from the liquid to form a vapor barrier or to make a thick formula. Corrosion inhibitors are added to stabili2e tbe formula for packaging (qv). 

Cupric chloride or copper(II) chloride [7447-39 ], CUCI2, is usually prepared by dehydration of the dihydrate at 120°C. The anhydrous product is a dehquescent, monoclinic yellow crystal that forms the blue-green orthohombic, bipyramidal dihydrate in moist air. Both products are available commercially. The dihydrate can be prepared by reaction of copper carbonate, hydroxide, or oxide and hydrochloric acid followed by crystallization. The commercial preparation uses a tower packed with copper. An aqueous solution of copper(II) chloride is circulated through the tower and chlorine gas is sparged into the bottom of the tower to effect oxidation of the copper metal. Hydrochloric acid or hydrogen chloride is used to prevent hydrolysis of the copper(II) (11,12). Copper(II) chloride is very soluble in water and soluble in methanol, ethanol, and acetone. 

Technology Description To achieve precipitation, acid or base is added to a solution to adjust the pH to a point where the constituents to be removed have their lowest solubility. Chemical precipitation facilitates the removal of dissolved metals from aqueous wastes. Metals may be precipitated from solutions as hydroxides, sulfides, carbonates, or other soluble salts. A comparison of precipitation reagents is presented in Table 7. Solid separation is effected by standard flocculation/ coagulation techniques. 

The classical methods used to separate the lanthanides from aqueous solutions depended on (i) differences in basicity, the less-basic hydroxides of the heavy lanthanides precipitating before those of the lighter ones on gradual addition of alkali (ii) differences in solubility of salts such as oxalates, double sulfates, and double nitrates and (iii) conversion, if possible, to an oxidation state other than -1-3, e g. Ce(IV), Eu(II). This latter process provided the cleanest method but was only occasionally applicable. Methods (i) and (ii) required much repetition to be effective, and fractional recrystallizations were sometimes repeated thousands of times. (In 1911 the American C. James performed 15 000 recrystallizations in order to obtain pure thulium bromate). 

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