From WolfWikis

Dissociation

In aqueous solutions many solutes like salts, acids and bases fall apart into ionic species, this is known as dissociation and the solutes are known as electrolytes. For some electrolytes the dissociation process essentially goes to completion or is at least extensive. They are known as strong electrolytes, for others the dissociation is only partial, the weak electrolytes.

Strong electrolytes

Dissociation of strong electrolytes is essentially a chemical reaction going to completion and can be written as such, e.g.:

NaCl_(s) ==(in water)==> Na⁺_(aq) + Cl^-_(aq)

Al₂(SO₄)_3(s) ==(in water)==> 2Al³⁺_(aq) + 3SO₄^2-_(aq)

As with all reaction equations dissolution equations need to be balanced to be of any use. It should be noted that the equation is necessary to compute proper concentrations in solution, because due to the stoichiometric coefficients the concentration of the resulting species is not necessarily the same.

If we start e.g. with 0.5 [mmol] of Al₂(SO₄)_3(s) and dissolve and dilute it to a solution of 1 l, the resulting concentrations will be [Al³⁺]= 1mM and [SO₄^2-]= 1.5 mM respectively. They are related by a ratio of 2:3, but are not identical.

The subscript aq means more than just that the ion in now in an aqueous solution. Most ions are enveloped in a covering of more or less tightly bound water molecules. These molecules are known as ligands. Often it is better to think of the ion plus its ligands as the solute than just the naked ion. However, to keep track of the stoichiometric bookkeeping this usually does not matter much. The subscript aq indicates that we choose to skip over such complicating details.

Strong (Brønsted) acids

Some protic acids dissociate almost entirely.

Some of them produce one proton: monoprotic

nitric acid: HNO₃ ==> H⁺+ NO₃^-

Others are diprotic

sulfuric acid: H₂SO₄ ==> 2H⁺+ SO₄^2-

There are also triprotics etc., but often the second or third proton do not dissociate as readily. They are better discussed under weak electrolytes.

The proton is actually hardly ever 'naked' in water. It is often better to associate it with at least one water molecule and speak of H₃O⁺ hydronium ions.

The [H⁺] (or [H₃O⁺]) concentration is often given on a logarithmic scale as pH= -¹⁰log[H⁺].

Notice that what we said about the aluminum and sulfate concentrations above also holds for the [H⁺] concentration. Dissolving 1 mol of acid into a solution of 1 liter (1M xxxx acid) will produce [H⁺]= 1 [mol/lit] or [H⁺]=2 [mol/l] depending on whether the acid xxx is nitric or sulfuric!

This is why in the older literature people sometimes used a different measure called normality. This was based on that amount of acid needed to produce a certain [H⁺]. So for example 1N sulfuric acid requires 0.5 moles of the acid dissolved into 1 liter. This measure has fallen in disuse and is not recommended.

Strong (Brønsted) bases. Hydroxides

These bases produce OH^- ions. There are not that many real strong ones, the most familiar ones are the hydroxides of the alkali metals like sodium and potassium. (The vernacular name is lye)

sodium hydroxide: NaOH ==> Na⁺+ OH^-

potassium hydroxide: KOH ==> K⁺+ OH^-

Some produce more than one hydroxide ion:

barium hydroxide: Ba(OH)₂ ==> Ba²⁺+ 2OH^-

Simple acid-base neutralization

When acidic and basic solutions are brought together the OH^- and H⁺ react to form water. This reaction is known as neutralization:

OH^- + H⁺ ==> H₂O

This reaction does not entirely go to completion. It goes to an equilibrium, but it is pretty extensive and for the moment we will pretend it runs its course completely.

Exercises

Exercise 1

Assuming full dissociation and complete solubility, what is [Cl^-], [La³⁺] and [Y³⁺] in the following aqueous solutions:

1. 0.1 mmol of LaCl₃ in 1 l
2. 0.05 mmol of LaCl₃ and 0.05 mmol of YCl₃ in 1 l
3. 0.05 mmol of LaCl₃ and 0.02 mmol of YCl₃ in 2 l
4. 0.05 mmol of LaCl₃ and 0.05 mmol of YCl₃ and 0.05 mmol of BaCl₂ in 1 l
5. 0.01 mmol of LaCl₃ and 0.05 mmol of YCl₃ and 0.05 mmol of La(NO₃)₃ in .5 l

Exercise 2

Assuming full dissociation and complete solubility, what is [SO₄^2-], [Al³⁺] and [K⁺] in the following aqueous solutions:

1. 0.05 mol of Al metal and 0.02 mol of KCl dissolved in 1 l of .2M sulfuric acid
2. 0.05 mol of Al₂(SO₄)₃ dissolved in 1 l of .1M nitric acid
3. 0.05 mol of Al₂(SO₄)₃ and 0.01 mol of K₂SO₄ in 500cm³ of aqueous solution.

Exercise 3

Assuming full dissociation and solubility and neglecting the water equilibrium, what is [H⁺] and pH of the following aqueous systems:

1. A solution of .05M nitric acid
2. A solution of .2M sulfuric acid
3. 200ml solution of .05M nitric acid added to 200ml solution of .2M sulfuric acid
4. .2 mol of HCl gas dissolved to form 500 ml of solution

Exercise 4

Assuming full dissociation and solubility and neglecting the water equilibrium, what is [OH^-] of the following aqueous systems:

1. 60 mmol of NaOH, Ba(OH)₂ and KOH each, together in 333 ml solution
2. 622 mg of anhydrous NaOH and 31 mg of anhydrous Ba(OH)₂ in 540 ml of aqueous solution
3. 50 gram of a 50% (by weight) solution of NaOH diluted to 1 liter.

Exercise 5

Assuming full dissociation and solubility and neglecting the water equilibrium, what is [H⁺] of the following aqueous systems after partial neutralization:

1. 500 ml of .1M sulfuric acid to which 30 mmol anhydrous NaOH is added.
2. 500 ml of .1M nitric acid to which 10 mmol anhydrous Ba(OH)₂ is added.
3. a mixture of 250 ml of .1M nitric acid and 250 ml of .2 sulfuric to which 1 g of anhydrous KOH is added.

Exercise 6

1. 5 g of anhydrous NaOH is exposed to an excess of dry HCl gas. What is the mass of water formed and the [Na⁺]. You may neglect the formation of water vapor, the solubility of HCl in water. Also take the volume of the final solution to be the volume of the water formed and assume complete solubility of the NaCl formed.
2. Are the above assumptions realistic?