| Contents for this page | Related topics |  |
|---|---|---|
|
The ionization of water Acid-base indicators Additional questions |
The common ion effect The strength of acids The hydrolysis of ions |
 Data Glossary |
| Learning Outcomes | ||
| After studying this section, you will (a) know the definition of pH and be able to apply it, (b) be able to calculate the pH of solutions of known hydrogen ion concentrations, and the hydrogen ion concentration of solutions of known pH, and (c), understand how indicators work. | ||
Water is an AMPHIPROTIC substance, in that it can act both as an acid and as a base:
The above equilibrium reaction will have an equilibrium constant K
Since the concentration of water can be considered to be a constant,
The constant Kw is known as the IONIC PRODUCT of water. It has an experimentally determined value of 1.0 x 10-14 at 25 ºC. The implication of the above is that in any aqueous solution (at 25 ºC) the product of [H3O+] and [OH-] will have the value of 1.0 x 10-14 and
Note that [H3O+] is often abbreviated as [H+], but always remember that in aqueous solution, hydogen ions or protons (H+) always exist bound to a water molecule, thus forming the H3O+ ion.
The function pH is defined as
The relationship between [H3O+] and pH is shown on the scale below:
At pH = 7.0, [H3O+] = [HO-], and the aqueous solution is said to be NEUTRAL. pH values less than 7.0 are ACIDIC, while those with a pH value greater than 7.0 are ALKALINE or BASIC.
| In the laboratory, the pH of solutions is measured with a pH meter, such as the one shown on the right. |  |
Note that theoretically, pH values may be less than 0 or greater than 14. Concentrated hydrochloric acid is about 10 mol.dm-3. If it were completely dissociated at that concentration (which it is not!), it would give a hydrogen ion concentration, [H+], of 10, corresponding to a pH = -1. Measurements of pH at high acid or base concentrations are unreliable.
| Approximate pH values of some aqueous solutions | |
| Solution | pH |
| gastric juice | 0.8 |
| household ammonia | 11.9 |
| human blood | 7.4 |
| lemon juice | 2.3 |
| milk | 6.4 |
| orange juice | 3.5 |
| sea water | 8.5 |
| urine | 6.0 |
| vinegar | 2.8 |
| Titrations are normally carried out by adding a solution of an acid (or alkali) from a burette to a accurately known volume of alkali (or acid). In the example shown here, the titration of a solution of acetic acid (ethanoic acid, CH3COOH) (in the conical flask) with sodium hydroxide NaOH (in the burette) is simulated.
We have added a drop of an indicator called phenolphthalein to the acid . Phenolphthalein is colourless at low pH. At pH 8, its colour changes to pink, and at pH 10, it changes to red. Now, neutralization occurs when all the acid in the flask has been converted to the sodium salt, which should have a pH of 9. So, when the colour is pink, the reaction is complete and we are at the "end-point". An additional drop will cause the colour to change to red, when too much NaOH will have been added! |  |
Acid-base indicators are weak organic acids that have a different colour from their conjugate bases. Taking phenolphthalein as example:
At low pH values, the concentrations of the red base form are so low that we cannot see any colour. Near pH 8, the concentration of the red base form becomes significant, to the point that the human eye can detect it as a pink tinge. This deepens in colour as the pH and the concentration of the red base form increase.
Indicators must be carefully chosen so that their colour changes take place at the pH values expected for an aqueous solution of the salt produced in the titration.
