Precipitates, 290 comparing solubility

Note any differences in the color of the precipitates compared to AgCl and in the solubility in the ammonia solution. 

The final volume is 20 mL. The millimoles Ag added equals 0.20 X 10 = 2.0 mmol. The millimoles Cl taken equals 0.10 X 10 = 1.0 mmol. Therefore, the millimoles excess Ag" equals (2.0 — 1.0) = 1.0 mmol. From Example 10.6, we see that the Ag" concentration contributed from the precipitate is small, that is, on the order of 10 mmol/mL in the absence of a common ion. This will be even smaller in the presence of excess Ag" since the solubility is suppressed. Therefore, we can neglect the amount of Ag contributed from the precipitate compared to the excess Ag . Hence, the final concentration of Ag is 1.0 mmol/20 mL = 0.050 M, and... 

Note that in the various cases the K p is equated to the solubility raised to different powers, depending on the cWge type of the precipitate. Hence, once cannot directly compare solubilities of salts of different charge types by examination of the numerical values of the respective solubility product constants. 

So far, two processes involving membrane separation for canola protein isolation are considered promising, and both are currently under commercial development Researchers at the University of Toronto developed a membrane-based process for canola protein isolation from defatted meal [29,42], in which, after precipitation, the soluble proteins were ultraflltered to be concentrated and diaflltered for purification. Two protein isolates were produced precipitated and soluble, with a combined protein recovery of more than 70% of total meal protein. Both products were high in protein (>85%), low in phytates (<1%), essentially free of glucosinolates (<2 (xmol/g) and had desirable functional properties comparable to those of soy protein. While the methionine content of both protein isolates was similar to the reported values, the soluble product was found to have a higher level of lysine than any canola proteins obtained before, and the precipitated protein isolate was, on the other hand, farther enriched with leucine [43]. This amino acid composition makes them suitable for nse in infant formulae (Table 4.6). Despite their excellent nutritive quality. 

Confirmation of the acidic nature of the original substance may be obtained, provided that it is sparingly soluble in water, by filtering the alkaline solution from undissolved material and acidifying the filtrate with concentrated hydrochloric acid, when the acidic substance should be precipitated (compare it with the original material). Under these conditions ampho-... 

The sodium soaps of fatty acid form calcium soaps of such low solubdity that they act as their own budders. Initial soap additions precipitate the calcium ion and the soap added thereafter functions in soft water. At high temperatures, the calcium soaps are relatively soluble compared to calcium tripolyphosphate. Thus sodium tripolyphosphate (STEP) can budd (revert) soaps in a hot water wash. However, at low temperatures the relative affinity of STEP for calcium decreases so that STEP cannot budd soaps in a cold water wash. 

Urea Enzymatic Dialysis Method. This method (16) uses 8 M urea [57-13-6] to gelatinize and facUitate removal of starch and promote extraction of the soluble fiber at mild (50°C) temperatures. EoUowing digestion with heat-stable a-amylase and protease, IDE is isolated by filtration or I DE is obtained after ethanol precipitation. Values for I DE are comparable to those obtained by the methods described eadier, and this method is less time-consuming than are the two AO AC-approved methods. Corrections for protein are required as in the AO AC methods. 

Some metals are soluble as atomic species in molten silicates, the most quantitative studies having been made with Ca0-Si02-Al203(37, 26, 27 mole per cent respectively). The results at 1800 K gave solubilities of 0.055, 0.16, 0.001 and 0.101 for the pure metals Cu, Ag, Au and Pb. When these metal solubilities were compared for metal alloys which produced 1 mm Hg pressure of each of these elements at this temperature, it was found drat the solubility decreases as the atomic radius increases, i.e. when die difference in vapour pressure of die pure metals is removed by alloy formation. If the solution was subjected to a temperature cycle of about 20 K around the control temperamre, the copper solution precipitated copper particles which grew with time. Thus the liquid metal drops, once precipitated, remained stable thereafter. 

Similar measurements were made for the heat of precipitation of silver iodide,5 which is even less soluble in water than silver chloride. As shown in Table 33 in Sec. 102, a saturated solution of Agl at 25°C contains only 9.08 X 10-9 molcs/liter, as compared with 1.34 X 10-6 for AgCl. By calorimetric measurement the heat of precipitation of Agl at 25°C was found to be 1.16 electron-volts per ion pair, or 20,710 cal/mole. 

The negative voltage shows that the state of equilibrium favors the reactants more than the products for the reaction as written. For standard conditions, the reaction will not tend to occur spontaneously. However, if we place Ag(s) in 1 M H+, the Ag+ concentration is not 1 M— it is zero. By Le Chatelier s Principle, this increases the tendency to form products, in opposition to our prediction of no reaction. Some silver will dissolve, though only a minute amount because silver metal releases electrons so reluctantly compared with H2. It is such a small amount, in fact, that no silver chloride precipitate forms, even though silver chloride has a very low solubility. 

Discussion. These anions are both determined as silver bromide, AgBr, by precipitation with silver nitrate solution in the presence of dilute nitric acid. With the bromate, initial reduction to the bromide is achieved by the procedures described for the chlorate (Section 11.56) and the iodate (Section 11.63). Silver bromide is less soluble in water than is the chloride. The solubility of the former is 0.11 mg L 1 at 21 °C as compared with 1.54 mg L 1 for the latter hence the procedure for the determination of bromide is practically the same as that for chloride. Protection from light is even more essential with the bromide than with the chloride because of its greater sensitivity (see Section 11.57). 

Sometimes it is important to know under what conditions a precipitate will form. For example, if we are analyzing a mixture of ions, we may want to precipitate only one type of ion to separate it from the mixture. In Section 9.5, we saw how to predict the direction in which a reaction will take place by comparing the values of J, the reaction quotient, and K, the equilibrium constant. Exactly the same techniques can be used to decide whether a precipitate is likely to form when two electrolyte solutions are mixed. In this case, the equilibrium constant is the solubility product, Ksp, and the reaction quotient is denoted Qsp. Precipitation occurs when Qsp is greater than Ksp (Fig. 11.17). 

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