Chapter 16 -- Equilibrium (Slightly Soluble Salts and Complex Ions)
Monday, October 01, 2007
12:28 PM
Chapter 16 -- Equilibrium (Slightly Soluble Salts and Complex Ions)
Many important ionic compounds are only slightly soluble in water. Chemists often refer to these
compounds as being "insoluble". This term isn't always accurate, though. In reality, many "insoluble"
compounds should be referred to as "slightly soluble". Barium Sulfate is a typical example of a slightly
soluble compound. Only 0.000246 g of BaSO4 will dissolve per 100 g of water. The decomposition of
Barium Sulfate does not involve water (due to insolubility) and therefore water is left out of the reaction:
We can derive an equilibrium constant for the solubility equilibrium based on this equation. This new
constant is called the solubility product constant (Ksp). It's formula is:
Ksp and Molar Solubility
Molar solubility is the molarity of a solute in a saturated aqueous solution. Molar solubility and Ksp are not
the same thing but we can use one to find the other.
Ksp values are not precisely known and its calculations are, therefore, subject to a higher amount of error.
There are two general types of Ksp problem types: determining a value of Ksp from experimental data and
calculating equilibrium concentrations from a known Ksp value.
At 20.0˚C, a saturated aqueous solution of silver carbonate contains 32mg of Ag2CO3 per liter of solution.
Calculate Ksp for Ag2CO3 at 20˚C. Calculate Ksp for the balanced equation is:
Ag2CO3 (s) ↔ 2 Ag+ + CO32-
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Many reference books list only Ksp values (and not solubilities for slightly soluble substances). It is helpful,
therefore, to be able to calculate the molar solubility of a solute from a given Ksp.
Calculate the molar solubility of silver sulfate at 25˚C using the following given information:
Ksp = 1.4 x10-5 at 25˚C
Ag2SO4 ↔ 2 Ag+ + SO4-2
The Common Ion Effect and Equilibrium
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The common ion effect appears when the same ion is created by two different compounds within the
same solution. For example, sodium sulfate dissociates completely in water to form Na+ and SO4-2 ions
according to this equation:
Na2SO4 → 2 Na+1 + SO4-2
If sodium sulfate is also added to the solution, the sulfate ion concentration increases due to the
combination of the two equations:
Ag2SO4 ↔ 2 Ag+ + SO4-2
Because the solution now has two sources of sulfate ion a stress is created on the original equilibrium.
According to Le Chatelier's Principle the solution will respond by shifting the equilibrium to the left. To
create this new equilibrium, several things must happen:
○ Some Ag2SO4 precipitates
○ [Ag+] is smaller than its original value
○ [SO4-2] is larger than its original value
Of course, a similar shift in equilibrium could be obtained by adding additional sources of the positive cation
(in this case it would be another source of Ag+).
All of this means that if a second source of a common ion is introduced the solubility of the originally
slightly soluble compound will be reduced. In plain terms, additional common ions reduce the solubility of a
compound.
An example calculation can be seen if we calculate the molar solubility of Ag2SO4 in 1.00 M Na2SO4.
Solubility, Activity, and Precipitation
In the previous section we saw that the solubility of a compound can be effected by a common ion. This
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solubility can also be affected by a non-common ion as well. CaF2 is an example of this phenomenon. Its
solubility is 50% greater in a solution of Na2SO4 than it is in water. This is puzzling since there are no
common ions present in a solution of CaF2 and Na2SO4. The reason behind such solubility is the increased
abundance of negative anions and positive cations in the solution compared to a solution of water and
CaF2. These interionic interactions are collectively called activity and can result in an "effective"
concentration. This is in contrast to "stoichiometric" concentration which is based only on the number of
particles present. We should use activity when performing equilibrium calculations.
A common lab mistake is using tap water instead of deionized water. Sometimes this would make no
difference, but often the solutions being used will react with some of the tap water ions and create
unintended products and precipitates. An example is a solution of tap water and AgNO3. The presence
of Cl- ions from tap water cause a precipitate to form (probably AgCl). What concentration of Cl- can be
tolerated without forming a precipitate?
Let's try to prepare 0.100 M AgNO3 using tap water with a [Cl-] = 1 x10-6 M.
We can write a reaction quotient in the form of an ion product. This is indicated by the symbol Qip. This
will allow us to compare Qip to Ksp.
Qip is much greater than Ksp (1 x10-7 > 1.8 x10-10). The overall meaning of this is that when the reaction
quotient is larger than K a reaction should occur in the reverse direction. In our example this means that
the Ag and Cl ions will combine to form AgCl. Once again, this situation means that a precipitation should
occur and that the precipitation will continue until Qip falls to a value equal to K.
A contrasting example involves the production of 0.100 M AgNO3 in deionized water with a [Cl-] = 1
x10-10.
The first step is to establish a Qip value.
In this case, Qip is smaller than Ksp (1 x10-11 < 1.8 x10-10). This scenario means that the reaction should
proceed in the forward direction however, this isn't possible because there is no AgCl to dissolve. There
will be no AgCl produced.
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Based on these examples we can establish several rules that allow us to predict what should happen when
we mix solutions containing ions capable of precipitating as a slightly soluble ("insoluble") solid.
Qip > Ksp = precipitation should occur
Qip < Ksp = precipitation should not occur
Qip = Ksp = saturated solution
A slightly soluble solid never totally precipitates but we consider precipitation to be generally complete if
99.9% of the target ion is precipitated. With this knowledge, we can set up some rules to determine if
precipitation is complete.
Conditions that are favorable for precipitation completeness:
○ A very small value of Ksp
○ A high initial concentration of the target ion
○ A concentration of a common ion that greatly exceeds the concentration of the target ion
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