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Unit – IV
Solubility equilibria
Solubility:
The solubility of a substance is the amount of that substance that will dissolve in a
given amount of solvent. “Solubility” may be considered to be equilibrium; the
equilibrium is between solid and ions in solution. Any ionic solid is 100% ionized in aqueous solution; once it actually dissolves.
NaCl(s) NaCl(aq) Na+(aq) + Cl–(aq)
The solubility of a substance fundamentally depends on the physical and chemical properties of the solute and solvent as well as on temperature, pressure and the pH of the solution. The extent of the solubility of a substance in a specific solvent is measured as the saturation concentration, where adding more solute does not increase the concentration of the solution and begin to precipitate the excess amount of solute.
Solubility product
Solubility product (Ksp), is the mathematical product of its dissolved ion concentrations raised to the power of their stoichiometric coefficients. This statement is called the solubility product principle.
MyXz (s) ⇌ yMZ+^ (aq) + zXY-(aq)
Example: Ca 3 (PO 4 ) 2 ⇌ 3Ca2+^ + 2 PO 4 3– Ksp = [Ca2+]^3 [PO 4 3–]^2
The dissolving of a precipitate is an equilibrium process. For example the dissolving of barium sulphate is represented as
BaSO 4 ⇌ Ba2+^ + SO 4 2– The dissolving of BaSO 4 to form Ba2+^ and SO 4 2–^ ion occurs at the same rate as the recombination of the Ba2+^ and SO 4 2–^ ion and thus we have equilibrium. A special form of the equilibrium constant is useful here. Because the amount of undissolved BaSO 4 does not affect the equilibrium system and it doesn’t make sense to refer to the concentration of an undissolved chemical, the equilibrium expression is modified and only the product of the concentration of the two ions appears. This constant is referred to as the solubility product constant Ksp.
Ksp = [Ba2+] [SO 4 2–] ---------- (1)
Example 2: For a solution of AgCl in water, the equilibrium is:
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H ^2 O
K sp Mz y Xyz
AgCl(s) Ag+(aq) + Cl–^ (aq)
The solubility product constant (Ksp) = [Ag+] [Cl–]
Solubility and solubility product:
Let us consider the reaction
BaSO 4 ⇌ Ba2+^ + SO 4 2–
The BaSO 4 that dissolves result in the formation of one Ba2+^ ions and one SO 4 2–^ ions.
(a)The solubility of BaSO 4 , which can be expressed as the number of moles of BaSO 4 that dissolve per litre, is equal to the concentration of the Ba2+^ ion and also the concentration of the SO 4 2–^ ion in these saturated solutions. (b)The Ba2+^ concentration and the SO 4 2–^ concentration are equal since for each BaSO 4 unit that dissolves one Ba2+^ and one SO 4 2–^ ion form. The above two points can be summarised as follows Solubility = s = [Ba2+] = [SO 4 2–] From equation (1) we know that Ksp = [Ba2+] [SO 4 2–] Ksp = s × s = s^2 Thus, given the solubility of solid ionic compound, the Ksp can be calculated.
Problem 1: Calculate the Ksp value of BaSO 4 which has a solubility of 3.9×10 ‒^5 mol/L at 25°C.
Solution: BaSO4(s) ⇌ Ba2+(aq)+ SO 4 2-(aq) Ksp = [Ba2+] [SO 4 2–] Ksp = s × s = s^2 = 3.9×10 ‒^5 × 3.9×10 ‒^5 = 1.52×10 ‒^9
Problem 2: What is the solubility of AgCl if the Ksp is 1.6 x 10-
Solution: AgCl(s) ⇌ Ag+(aq) + Cl-(aq) K sp = [Ag+][Cl-] If s is the solubility of AgCl, then: [Ag+] = s and [Cl-] = s K sp = s × s = s^2 = 1.6 x 10- s = s = 1.3 x 10-5^ mol/L.
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1.6x 10 -
A special type of titrimetric procedures involves the formation of precipitates during the course of a titration. The titrant reacts with the analyte forming an insoluble material and the titration continues till the very last amount of analyte is consumed. The first drop of titrant in excess will react with an indicator resulting in a color change and announcing the termination of the titration.
Argentometric Titrations
The most widely applicable precipitation titrations involve the use of silver nitrate with chlorides, bromides, iodides, and thiocyanate. Since silver is always there, precipitation titrations are referred to as Argentometric titrations. This implies that this type of titration is relatively limited.
Conditions for precipitation titration:
Any precipitation reaction could in theory be adopted to a volumetric technique, provided that (1) The precipitation reaction reaches equilibrium very rapidly after each addition of titrant. (2) No interfering situations like co-precipitation, occlusion or adsorption of foreign ions occur. (3) An indicator capable of locating the stochiometric equivalence point accurately is available.
Determination of chloride by Volhard’s method:
Volhard’s method is used in the estimation of chloride ions with standard solution of AgNO 3. It is an example of titration in which indicator forms a coloured complex ion with the titrant (i.e) Silver nitrate. In this method, excess of AgNO 3 , solution is added in order to precipitate out the anions to be estimated. Then excess (unreacted) Ag+^ ions are determined by back titration with standard potassium or ammonium thiocyanate solution. The end point is detected by adding indicator like ferric ammonium sulphate or ferric alum which forms a soluble red complex with excess of titrant (i.e) NH 4 SCN or KSCN.
Determination of Cl–^ ion involves: (a) Preparation of standard ammonium thiocyanate: By weighing appropriate quantity of ammonium thiocyanate, its 250ml of 0.1M solution is prepared in distilled water.
(b) Determination of unknown chloride ions: Given chloride solution is diluted to 250ml using distilled water. From this 25ml solution is pipette out into a titration flask, to which 25ml of 0.1M AgNO 3 is added. After shaking the solution, 2ml of nitrobenzene and 5ml of 50% HNO 3 and ferric alum is added as an indicator. It is then titrated with
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standard NH 4 SCN till reddish brown colour is obtained to supernatant solution. Suppose the reading is ‘X’ ml. Then, (25–X) ml will be the amount of NH 4 SCN required to precipitate Cl–^ ions. Using following chemical reactions amount of Cl–^ ions in given solution can be calculated. Cl–^ + AgNO 3 → AgCl ↓ + NO 3 –^ + AgNO 3 Ag+^ + SCN–^ → AgSCN↓
Fe3+^ + SCN–^ → Fe(SCN)2+
Adsorption indicators:
Commonly used indicators for titration of chloride against silver nitrate are Fluorescein, Dichlorofluorescein and Eosin.
(a) Fluorescein: It is a suitable indicator for very dilute solution of chloride with silver nitrate in neutral or faintly acidic solution. At the end point there is a sudden change of white precipitate in greenish yellow medium to distinctly red. (b) Dichlorofluorescein: It is a suitable indicator for every dilute solution of chloride (say drinking water) and works even in the presence of acetic acid and weakly acid solution. The end point is from yellowish green to red. (c) Eosin: It is used for the titration of bromide and iodide with silver nitrate(from burette) in the presence of acetic acid. The end point is the change from pink to reddish violet.
Adsorption indicator Technique :
Titration of chloride by silver ion:
In fajans method of titrating Cl–^ with Ag+, the adsorption indicator fluorescein is used. During the titration and up to the equivalence point the precipitate of AgCl adsorbs Cl–^ ions present in excess in solution. The fluorescein ion being negative is repelled by the negatively charged precipitate surface. The solution at this stage has the yellow-green colour imparted by fluorescein. At the equivalence point, all the adsorbed Cl–^ ions are removed by Ag+. On the addition of the first excess of Ag+, AgCl surface adsorbs Ag+^ ions and the surface becomes positively charged.
The fluorescein ions now strongly get adsorbed on this positive surface. The silver fluorescein formed has a deep red colour and the formation of this colour signals the end point. A little excess of Ag+^ is required to form the silver fluorescein and to indicate the endpoint.
Gravimetric methods of analysis:
Gravimetric procedures always involve the separation of the analyte constituent from the sample so that it can be isolated and weighed. This separation can be a physical separation, such as through solubility or volatilization or it can be a chemical separation (i.e) by chemical reaction, so as to form a precipitate that can be weighed.
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Factors which affect solubility of precipitate:
Temperature: heat is absorbed as most solids dissolve. Therefore the solubility of precipitates increases with increasing temperature. Solubility products for most sparingly soluble compounds increase at higher temperatures.
Solvent composition: The nature of solvent influences the solubility of most inorganic substances. It is markedly less in mixtures of water and organic solvents than in pure water.
Common ion effect and Diverse ion effect: The solubility of a slightly soluble salt is decreased by the presence of an ion which is in common with one of the ions of the slightly soluble salt. For example, in the precipitation of AgCl from a solution of AgNO 3 using KCl as the precipitant, the addition of an excess of KCl will serve due to common ion effect to reduce the molar solubility of AgCl. However the solubility of AgCl becomes significantly large in solutions containing a high chloride concentration, such as 1.0M or above owing to the formation of AgCl2.–^ and AgCl 3 2–^ species. The benefit of the common ion effect in suppressing the solubility of a precipitate is lost in the need to wash the precipitate throroughly to remove the common ion.
Diverse ion effect: Many slightly soluble salts show an increased solubility in the presence of increased concentration of certain salts having no ion in common with those of the slightly soluble salt. Thus, AgCl is more soluble in KNO 3 solution than in water. This effect is called the diverse ion or neutral salt effect.
Purity of Precipitates:
Purity of a precipitate is a vital aspect for successful gravimetric determination. The precipitate may be contaminated with one or more other substances because those other substances are themselves insufficiently soluble in the mother liquid. A precipitated phase can become contaminated with substances from its mother liquid even when the equilibrium solubilities of those other substances are not exceeded. The surface of a precipitate has positive and negative centres and these may attract negative and positive ions form the mother liquid. Such attracted ions would get adsorbed on the surface and thus contaminate the precipitate.
General rules of precipitation:
(a) Precipitation should be carried out in dilute solution to minimize the error due to co-precipitation. (b) The reagents should be mixed slowly with constant stirring. This will keep the degree of super saturation low, promoting the growth of large crystals. (c) A slight excess of the reagent is generally required to ensure complete precipitation. (d) Precipitation is effected in hot solutions, provided the product is thermally stable. At higher temperatures the solubility is increased with a consequent reduction in the degree of super saturation.
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(e) The completion of precipitation should be ascertained by adding a few drops of the precipitant through the side of the beaker and observing for absence of fresh precipitate formation in the supernatant liquid. (f) Crystalline precipitates should be digested preferably overnight, except in those cases, where post precipitation may occur. Digestion on the stream bath is desirable. Digestion has no effect upon amorphous or gelatinous precipitate. (g) The precipitate should be washed with the appropriate dilute solution of an electrolyte (h) If appreciable co-precipitation is a possibility, the precipitate should be dissolved in a suitable solvent and reprecipitated.
Types, care and use of crucibles:
Different types of crucibles are used for collecting and weighing the precipitates.
They are mainly of two types.
(1) Crucibles made of silica or porcelain are used when the precipitates are ignited to constant weights. (2) Sintered crucibles are used for precipitates which are weighed after drying in air- oven. Crucibles are expensive. They should be handled properly for obtaining reliable gravimetric data. They should never be touched but must always be handled with a pair of tongs. They should not be placed on the table but only on a clean tile. Porcelain crucibles: These are usually of 3 to 4 cm in diameter. These should be carefully hated with a small flame in order to avoid fracture before being heated to a high temperature. These are suited for heating the precipitates to very high temperatures. Porcelain crucibles are suitable for use up to a temperature of 1000 – 1200 °C. Silica crucibles: These are more expensive than porcelain crucibles. However, these may be safely exposed to sudden changes of temperature without any risk of fracture. These must not be used for alkalis or hydrofluoric acid, because these chemicals attack silica. Certain precipitates tend to adhere strongly to these crucibles and hence these crucibles should be carefully cleaned immediately after use. Platinum Crucible: Platinum crucibles are marketed in various shapes and sizes. The choice for an experiment is decided by the quantity of precipitate handled. These are very expensive and hence their use is very much restricted to cases where they are indispensable. These crucibles can be more readily and more uniformly heated to redness than porcelain crucibles. The Gooch crucible: A Gooch crucible is made of porcelain with the bottom perforated with a number of small holes. It is usually fitted into a glass adaptor (Gooch funnel) by means of narrow rubber ring, the adapter passing through the rubber stopper of a filter flask. The bottom of the crucible is covered with asbestos. This covering is done by shaking the asbestos fibre bits with water and pouring the mixture into the crucible. The water drains leaving the asbestos as a
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by a rapidly growing crystal. Contaminants are located within the crystal and therefore washing will not remove these impurities form the precipitate. (b) Post-precipitation: After the formation of the precipitate, the impurities will get soluble. This is another type of precipitate contamination. Example: When calcium is determined as its oxalate and if the solution contains magnesium ions, magnesium oxalate will be precipitated slowly on calcium oxalate. This post precipitation is greater when the precipitate is kept in contact with the mother liquid longer. It occurs with sparingly soluble substances which form supersaturated solutions. These substances usually have an ion in common with the primary precipitate.
Ca2+^ (NH 4 ) 2 C 2 O 4 CaC 2 O 4
If Mg2+^ is present as an impurity in Ca2+, then we can have a post precipitation of MgC 2 O 4 over CaC 2 O 4. The post precipitation can be avoided by keeping a high pH.
Differentiate post-precipitation and co-precipitation:
S.No Post-precipitation Co-precipitation
Increases with time and increases with higher temperature
When we allow the main precipitate with another, liquor, then they will go into solution
Magnitude of error/ contamination in post precipitation is greater
Contamination will be less
Post precipitation can be avoided
(a) By adding some masking agents
(b)By maintaining a high pH.
Examples of co-precipitation are
When estimating barium
Ba2+^ + SO 4 2–^ BaSO 4
Contaminated with CrO 4 2–^ BaCrO 4
The contaminants will go and get occluded inside the crystal lattice of BaSO 4. These are the ways the co-precipitation can cause erroneous results. In order to minimise this co-precipitation we digest the precipitate. Digesting the precipitate in suitable solvent and reprecipitation gives good yield on co-precipitation.
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