Solubilty Product Constant of Baso4
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
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