Solubility 1

Solubility

A. Solubility Product

In this chapter, we will be discussing the solubility of ionic compounds (salts) in water. However, solubility, at least qualitatively, is very roughly defined. A salt is�

o Soluble if at least 0.1 mol dissolves in 1 L water

o Slightly soluble if about 0.001 � 0.1 mol dissolves in 1 L

o Insoluble if less than 0.001 mol dissolves in 1 L

Realize that, with ionic compounds, they are completely ionized in water, but ONLY WHEN THEY ARE DISSOLVED. Thus, the only the portion that dissolves is ionized!

In other words, we have a heterogenous equilibrium between the undissolved solid and the ions that are in solution. e.g.

BaSO4(s) Ba2+(aq) + SO42 (aq)

][SO ][BaK 24

2sp = 1.1×10 10 at 25 °C

We use the symbol Ksp for the equilibrium constant, and it is referred to as the solubility product.

Note that for an equilibrium to be established, we must have some solid present (doesn�t matter how much, just some).

Application: Ba2+ is very toxic, yet BaSO4 is used in humans for medical imaging because very little BaSO4 dissolves.

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Calculations are quite straightforward and no different from those of other equilibria, such as weak acids and bases.

Example: How much barium sulfate dissolves in a litre of water? (Ksp = 1.1×10 10)

BaSO4(s) Ba2+(aq) + SO42 (aq)

Note that if you�re given the solubility, either in moles or grams per litre, you can do the reverse calculation to determine Ksp.

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B. Common Ion Effect

The presence of common ions will reduce the solubility of the compound, much like how the ionization of a weak acid is decreased if a common ion were present.

If we have barium sulfate already in equilibrium, what happens to the equilibrium we add more sulfate?

BaSO4(s) Ba2+(aq) + SO42 (aq) Ksp = 1.1×10 10

Calculate the solubility of barium sulfate in 0.01 M Na2SO4. How does this compare to the solubility on the last page?

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Example: What is the concentration of Ca2+ present in a solution made by equilibrating solid CaF2 (Ksp = 4.1×10 11) with 0.20 potassium fluoride?

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C. Unsaturated, Saturated, and Supersaturated

A solution is deemed to be saturated if the equilibrium between the solid and the dissolved ions is established.

o In this case, reaction quotient Q = Ksp(we also call Q the ion product)

What if the value of Q < Ksp? Then the solution is considered to be unsaturated, and more BaSO4(s) will dissolve, provided that there is some present.

Yet, if Q > Ksp the solution is supersaturated, and it contains more of the product ions than it could hold. Thus, when Q is larger than the solubility product, some precipitation will always occur!

Example: If solid Ba(NO3)2 is added to a 0.1 M solution of Na2SO4, assuming no volume change, how much Ba(NO3)2 can be added before barium sulfate starts to precipitate?

o Note: we assume that precipitation will occur at the point where Q just exceeds Ksp

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Example: PbCl2 has Ksp = 1.7×10 5. If equal volumes of 0.030 M Pb(NO3)2 and 0.030 M KCl are mixed, will precipitation occur?

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D. pH and Complexation

The solubility of many substances influences, or is influenced by, the pH of the solution.

1. Metal hydroxides (usually have low solubility)

Example: What is the pH of a solution made by equilibrating water with solid Ca(OH)2, which has Ksp = 7.9 × 10 6?

o This solution is quite basic. Some other metal hydroxides have very low solubilities, so their solutions are not basic.

Another example: Cu(OH)2 has 1.6 × 10 19. What is the maximum [Cu2+] possible in neutral solution? What is the maximum pH of a solution in which Cu2+ = 0.50 M?

o Reasonable [Cu2+] are only possible in acidic solution!

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2. Solubilities of salts of weak acids

Solubilities of such salts are affected by pH. However, we will only look at them from a qualitative point of view.

Example: AgCl and AgBr are not very soluble; when HNO3 is added, nothing happens. Silver acetate (CH3COOAg) also has a low solubility in water, but when HNO3 is added, it dissolves. Explain this behaviour.

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3. Complexation

As we saw on the last page, the consumption of one of the products results in the dissolving of more solid.

In complexation, we consume one of the products by complexing it with a species known as a complexant. Again, we will only treat this qualitatively.

Example: The addition of NH3 to a solution of silver chloride (equilibrated in water) causes the solid to dissolve. Explain.

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E. Applications of Solubility: Limestone

Limestone is used to construct many historically important buildings, statues, monuments, and tombstones.

o Greek Parthenon

o Government buildings

o UWO buildings

o $270,000 monoliths welcoming visitors to London at the corner of Oxford St and Airport Rd

Limestone contains CaCO3, which has a Ksp of 3.8 × 10 9.

CaCO3(s) Ca2+(aq) + CO32 (aq)

With the small Ksp, it is virtually insoluble in pure water. However, rain water gradually dissolves it away. Why?

o Rain water contains both carbonic and sulfuric acids, both of which react with carbonate.

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1. Carbonic Acid

CO2 in the air dissolves in water to form carbonic acid.

H2O + CO2 H2CO3

This reacts with the carbonate (CO32 ) in an acid-base

neutralization reaction, thereby decreasing [CO32 ].

H2CO3 + CO32 2 HCO3

So, more CaCO3 dissolves to compensate for the amount of carbonate that was used up. The dissolved products are washed away by the rain.

2. Sulfuric Acid

SO2 is produced from the burning of fossil fuels that contain sulfur. Power plants, steel mills, paper mills, refineries, and smelters are the largest producers, as are auto emissions.

SO2 is oxidized to SO3 in the atmosphere. This then combines with water to form H2SO4.

The sulfuric acid reacts with the carbonate from the limestone, which causes more limestone to dissolve.

H+ + CO32 HCO3

H+ + HCO3 H2CO3 H2O + CO2

(If limestone is dropped in acid, it dissolves as CO2 is evolved)