Solubilty Product Constant of Baso4

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  • Topic: Solubility, Solubility equilibrium, Solutions
  • Pages : 9 (2660 words )
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  • Published : March 11, 2013
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ubility * 1
To measure the molar solubility of a sparingly soluble salt in water. * 2
To prepare a calibration curve based on complex ion formation for absorbance enhancement. * 3
To calculate the solubility product constant (Ksp) of a sparingly soluble salt from its molar solubility. * 4
To confirm the common ion effect on the molar solubility of a sparingly soluble salt. Introduction
In previous introductory chemistry courses, you learned some basic solubility rules that are useful in determining if an ionic solid will dissolve in water. Solids that dissolve completely, such as NaCl and NH4NO3, were referred to as "soluble" and others that did not dissolve completely, such as AgCl and BaSO4, were referred to as "insoluble". In fact, very few ionic solids are completely insoluble, meaning that they will not form any ions when placed in aqueous solution. Most solids that are commonly referred to as "insoluble" are actually slightly soluble and will produce an equilibrium between undissolved solid and ions in solution. For example, when copper (II) iodate (Cu(IO3)2) is placed in water, the following equilibrium is established. ( 1 )

Cu(IO3)2(s) Cu2+(aq) + 2 IO3-(aq)
The equilibrium constant associated with this reaction is called the solubility product constant and is given the symbol Ksp. ( 2 )
Ksp = [Cu2+][IO3− ]2
It is important to emphasize that the equilibrium in equation 1 is only true if some solid is present. If the solid completely dissolves in solution, then the product of the ions as shown in equation 2 is not equal to the Ksp. However, as long as some solid is in contact with solution, the solution will become saturated with the ions according to equation 1. The molar solubility of a solid is the maximum number of moles of the solid that will dissolve in one liter of solution. Molar solubility is measured in moles/liter and has units of molarity (M). Molar solubility can be determined by measuring the concentration of the ions formed in a solution saturated with the solid. For example, if the molarity of Cu2+ in a solution saturated with Cu(IO3)2 can be determined, the stoichiometry in equation 1 indicates that this is also equal to the molarity of Cu(IO3)2 that dissolved in solution or its molar solubility. Alternatively, if the molarity of IO3- in a solution saturated with Cu(IO3)2 can be determined, the stoichiometry in equation 1 indicates that the molar solubility of Cu(IO3)2 in water is one half the [IO3-] in solution.In a saturated solution, if all the ions came from the solid, then the ratio of cations to anions is known. In the case of Cu(IO3)2, it is known that 2[Cu2+] = [IO3-]. Thus, if the Cu2+ concentration can be measured, the IO3- concentration can be calculated and vice versa. These concentrations can then be entered into the Ksp expression (equation 2) to solve for the solubility product constant of Cu(IO3)2. The preceding discussion referred to dissolving the sparingly soluble salt in pure water. However, if some of the ions that are to be produced by the solid are present in solution from another source, Le Châtelier's Principle predicts that the equilibrium in equation 1 will shift to the left and less solid will dissolve. This would result in a lower molar solubility. This is referred to as the common ion effect. In the example of Cu(IO3)2, the presence of either Cu2+ or IO3- in solution should result in a lower molar solubility than in pure water. Earlier this semester in the Solutions and Spectroscopy lab, you prepared a calibration curve for Cu2+ over the range of approximately 0.1 M to 0.4 M. As you may recall, Cu2+ in solution has a pale blue color and you were able to measure its absorbance at 620 nm. For this experiment, it is necessary to prepare a calibration curve for Cu2+ over the range of 5 x 10-4 M to 0.01 M. At these concentrations, the absorbance of Cu2+ is too low to detect directly by spectroscopy. To improve the absorbance of the Cu2+ solutions,...
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