Saturated solution
If the solute’s concentration is less than its solubility, the solution is said to be unsaturated.
A solution that contains a relatively low concentration of solute is called dilute, and one with a relatively high concentration is called concentrated.
A saturated solution is a chemical solution containing the maximum concentration of a solute dissolved in the solvent. Additional solute will not dissolve in a saturated solution.
Solutions may be prepared in which a solute concentration exceeds its solubility. Such solutions are said to be supersaturated, and they are interesting examples of nonequilibrium states.
For example, the carbonated beverage in an open container that has not yet “gone flat” is supersaturated with carbon dioxide gas; given time, the CO2 concentration will decrease until it reaches its equilibrium value.
Unsaturated Solution: Q (reaction quotient) < Ksp
Saturated Solution: Q = Ksp
Supersaturated Solution: Q > Ksp
Solubility equilibrium - a dynamic equilibrium between a solute and a solvent in a saturated solution in a closed system.
The Solubility Product Constant (Ksp)
A solubility equilibrium is a heterogeneous equilibrium system between a solid ionic compound and its ions dissolved in a saturated aqueous solution.
The equilibrium law equation for the solubility equilibrium between solid silver iodide and its aqueous ions
AgI (s) ⇌ Ag+ (aq) + I- (aq)
can be written as:
K = [C]c[D]d / [A]a[B]b
K = [Ag+ (aq)] [I- (aq)] / [AgI (s)]
The concentration of silver and iodide ions may vary from one saturated solution to the next at the same temperature.
However, the concentration of solid silver iodide remains constant. Therefore, we can ignore the concentration of the solid silver iodide in the equilibrium law equation for the solubility equilibrium.
The equilibrium constant equation for silver iodide therefore becomes:
K = [Ag+ (aq)] [I- (aq)]
This equation gives us the solubility product constant (Ksp), which is the value of the equilibrium law equation for a solubility equilibrium.
The subscript “sp” distinguishes this constant from the general equilibrium constant.
The general equilibrium constant, K, is sometimes written as Kc or Keq to distinguish it from other equilibrium constants, such as Ksp.
The Ksp of AgI(s) at 25 °C is 8.5 3 10-17.
Temperature is always included when reporting values for the solubility product constant.
Solubility product constant (Ksp) values vary with temperature because the solubility of substances varies with temperature.
Solubility and the Solubility Product Constant
The solubility of an ionic compound varies with the ions it contains.
In general, the more highly charged the anions and cations are, the less soluble the ionic compound will be, because it takes more energy for a solvent to break the ionic bonds in the crystal lattice.
Therefore, ionic compounds containing divalent ions, such as calcium ions (Ca2+) and carbonate ions (CO32-), are generally less soluble than ionic compounds containing monovalent ions, such as sodium ions (Na+) and nitrate ions (NO3-).
Predicting Precipitation
Solubility is a complex property and difficult to predict. The only reliable way to know if an ionic compound is soluble is to do the experiment.
Solubility of Some Ionic Compounds at SATP
SATP - Standard Ambient Temperature and Pressure (298.15K (25°C), 101.3kPA (1 atmosphere))
To predict if a precipitate will form when two aqueous solutions are mixed, it is necessary to use a solubility table to predict if any ionic compounds will form that will be of low solubility.
Example:
Mixing an aqueous solution of sodium chloride, NaCl(aq), with an aqueous solution of silver nitrate, AgNO3 (aq).
These two ionic compounds are both highly soluble in water.
NaCl (s) → Na+ (aq) + Cl- (aq)
AgNO3 (s) → Ag+ (aq) + NO3- (aq)
Combined solution therefore will contain aqueous sodium ions, chloride ions, silver ions, and nitrate ions.
Predicting of precipitation
In a solution only the combination of silver and chloride ions is likely to form a precipitate.
Whether a precipitate actually forms depends on the concentrations of the silver ions and the chloride ions in the solution.
When we know the concentrations of ions in aqueous solution, we can use a quantitative method to predict whether a precipitate will form.
The trial ion product (Q) is the reaction quotient applied to the initial ion concentrations of a slightly soluble ionic compound.
We get a value for the trial ion product (Q) by multiplying the concentrations of ions in a specific solution raised to powers equal to their coefficients in a balanced chemical equation.
Then, we compare the value of the trial ion product, Q, with the solubility product constant, Ksp.
The following rules are used to predict whether precipitation will occur:
• If Qsp > Ksp, the dissolution equilibrium system shifts to the left. Precipitation occurs and will continue until the solution reaches a new equilibrium.
• If Qsp < Ksp, the dissolution equilibrium system will shift to the right, and no precipitation occurs. The solution is unsaturated, so more solid can dissolve.
• If Qsp = Ksp, the solution is at equilibrium. No precipitation or overall change in concentration will occur.
The Common Ion Effect
The position of a solubility equilibrium system can be shifted by changing the temperature, the concentration of the dissolved ions, or both.
Common ion effect - a reduction in the solubility of an ionic compound due to the presence of a common ion in solution.
Adding a common ion to a solution increases the concentration of that ion in solution. As a result, equilibrium shifts away from the ion.
The Common Ion Effect
Adding chromate ions to an equilibrium system of lead(II) chromate (A) causes the equilibrium position to shift to the left.
As a result, more solid lead(II) chromate precipitates (B).
