Common-ion concentration

When solving common-ion-effect problems, calculations like the ones above involving finding concentrations and Ksp s can still be done, but the concentration of the additional common ion will have to be inserted into the solubility product constant expression. Sometimes, if the Ksp is very small and the common ion concentration is large, we can simply approximate the concentration of the common ion by the concentration of the ion added. 

In the presence of a common ion of the precipitate, the contribution of the solubility to the total common ion concentration often may be neglected. If excess M" is present to the extent Cmj the solubility 5 may be calculated from... 

The concentration of AgCl in the solution is thus 1.7 x 10 M this is the solubility under these conditions. In view of equation (7.67), the solubility will decrease as the common ion concentration increases. 

We express now the conductance as the specific conductivity Xjp divided by the common ion concentration ci (molar conductivity),... 

Addition of excess H ions to this solution will cause the equilibrium to move towards undissociated acid thereby decreasing the concentration of Ac . This effect is known as the common-ion effect and is of considerable practical importance. Thus, e.g. in the precipitation of metal ions as insoluble sulphides. 

Sodium sulphate crystallises out in hydrated form (common ion effect) and is filtered off on concentration, sodium dichromate is obtained. For analytical purposes, the potassium salt. K2Cr20-. is preferred potassium chloride is added and the less soluble potassium dichromate obtained. 

The high sodium ion concentration results in facile crystallisation of the sodium salt. This process of salting out with common salt may be used for recrystallisation, but sodium benzenesulphonate (and salts of other acids of comparable molecular weight) is so very soluble in water that the solution must be almost saturated with sodium chloride and consequently the product is likely to be contaminated with it. In such a case a pure product may be obtained by crystallisation from, or Soxhlet extraction with, absolute alcohol the sul-phonate is slightly soluble but the inorganic salts are almost insoluble. Very small amounts of sulphones are formed as by-products, but since these are insoluble in water, they separate when the reaction mixture is poured into water ... 

A particular concentration measure of acidity of aqueous solutions is pH which usually is regarded as the common logarithm of the reciprocal of the hydrogen-ion concentration (see Hydrogen-ION activity). More precisely, the potential difference of the hydrogen electrode in normal acid and in normal alkah solution (—0.828 V at 25°C) is divided into 14 equal parts or pH units each pH unit is 0.0591 V. Operationally, pH is defined by pH = pH(soln) + E/K, where E is the emf of the cell ... 

Fig. 10-11. The pH scale is a measure of hydrogen ion concentration. The pH of common substances is shown with various values along the scale. The Adirondack Lakes are located in the state of New York and are considered to be receptors of acidic deposition. Source U.S. Environmental Protection Agency, Acid Rain—Research Summary," EPA-600/8-79-028, Cincinnati, 1979.

It was pointed out earlier that the low nucleophilicity of fluoride ion and its low concentration in HF solutions can create circumstances not commonly observed with the other halogen acids. Under such conditions rearrangement reactions either of a concerted nature or via a true carbonium ion may compete with nucleophilic attack by fluoride ion. To favor the latter the addition of oxygen bases, e.g., tetrahydrofuran, to the medium in the proper concentration can provide the required increase in fluoride ion concentration without harmful reduction in the acidity of the medium. 

The rate constant k is measured as a function of (added) concentration of the common ion X , and from the plot according to Eq. (4-76) the ratio k i/k2 = slope/intercept is evaluated. 

Except for those reactions whose characteristic rate constants vary linearly with the hydronium or hydroxide ion concentration, the most effective presentation of pH-rate data is a graphical one. Two kinds of plot pH-rate profiles) are commonly seen ... 

These substances accelerate the reaction, and their effectiveness increases in the order given. This suggestion was questioned by Pocker, who found that the effects of such added substances were not directly proportional to their concentrations and could easily be explained by macro effects on the solvent character. He also found that common-ion effects were small in the reaction, the effect of added 1-methylpyridinium bromide was negligible, and that there was no evidence for surface catalysis on the walls of the vessel. There is an exact parallel between the relative rates of the Finkelstein reactions... 

The pH is one of the most important characteristics of an electrolyte, commonly expressed as a number between zero and fourteen, and is the negative logarithm of the hydrogen ion concentration [195]. 

The predictions of the pH/potential diagram are generally fulfilled, but in very concentrated acid solutions, attack may diminish, owing to the relative insolubility of the relevant salt in the acid. Thus, lead nitrate, although soluble in water, has (owing to common ion effect) only slight solubility in concentrated nitric acid, and the corrosion rate is reduced. Similarly, lead chloride is less soluble in moderately concentrated hydrochloric acid than... 

An increase in carbonate-ion concentration moves the equilibrium in favour of calcium carbonate deposition. Thus one secondary effect of cathodic protection in seawater is the production of OH , which favours the production of CO, , which in turn promotes the deposition of CaCOj. Cathodically protected surfaces in seawater will often develop an aragonite (CaCOj) film. This film is commonly referred to as a calcareous deposit. 

The acidity of a solution has pronounced effects on many chemical reactions. It is therefore important to be able to learn and control the hydrogen ion concentration. This control is obtained through application of the Equilibrium Law. Common types of calculation, based on this law, are those needed to determine KA from experimental data and those using KA to find [H+], We will illustrate both of these types, using benzoic acid, QH6COOH, as an example. 

The great importance of the solubility product concept lies in its bearing upon precipitation from solution, which is, of course, one of the important operations of quantitative analysis. The solubility product is the ultimate value which is attained by the ionic concentration product when equilibrium has been established between the solid phase of a difficultly soluble salt and the solution. If the experimental conditions are such that the ionic concentration product is different from the solubility product, then the system will attempt to adjust itself in such a manner that the ionic and solubility products are equal in value. Thus if, for a given electrolyte, the product of the concentrations of the ions in solution is arbitrarily made to exceed the solubility product, as for example by the addition of a salt with a common ion, the adjustment of the system to equilibrium results in precipitation of the solid salt, provided supersaturation conditions are excluded. If the ionic concentration product is less than the solubility product or can arbitrarily be made so, as (for example) by complex salt formation or by the formation of weak electrolytes, then a further quantity of solute can pass into solution until the solubility product is attained, or, if this is not possible, until all the solute has dissolved. 

As shown above the sulphide ion concentration of a saturated aqueous solution of hydrogen sulphide may be controlled within wide limits by suitably changing the concentration of hydrogen ions—a common ion—of the solution. In a like manner the hydroxide ion concentration of a solution of a weak base, such as aqueous ammonia (Kb = 1.8 x 10-5), may be regulated by the addition of a common ion, e.g. ammonium ions in the form of the completely dissociated ammonium chloride. The magnitude of the effect is best illustrated by means of an example. In a 0.1M ammonia solution, the degree of dissociation is given (Section 2.13) approximately by. 

Calcium oxalate monohydrate has a solubility of 0.0067 g and 0.0140 g L 1 at 25° and 95 °C respectively. The solubility is less in neutral solutions containing moderate concentrations of ammonium oxalate owing to the common-ion effect (Section 2.7) hence a dilute solution of ammonium oxalate is employed as the wash liquid in the gravimetric determination. 

To measure the hydrogen ion concentration of a solution the glass electrode must be combined with a reference electrode, for which purpose the saturated calomel electrode is most commonly used, thus giving the cell ... 

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. 

FIGURE 11.15 If the concentration of one of the ions of a slightly soluble salt is increased, the concentration of the other decreases to maintain a constant value of Ksp. (a) The cations (pink) and anions (green) in solution, (b) When more anions are added (together with their accompanying spectator ions, which are not shown), the concentration of cations decreases. In other words, the solubility of the original compound is reduced by the presence of a common ion. In the insets, the blue background represents the solvent (water). 

We can gain a quantitative understanding of the common-ion effect by considering how a change in the concentration of one of the ions affects the solubility product. Suppose we have a saturated solution of silver chloride in water ... 

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