F.6/7 Chemistry Practical
Acid-base equilibrium in aqueous solution
Objective:
(a) To study acid-base equilibrium in aqueous solution
(b) To contrast buffer solution and non-buffer solution
Group size: Individual
Introduction
Aqueous solutions of hydrogen chloride and ammonia will be readily recognized as examples
of an acidic and a basic solution, respectively, but substances regarded as salts may act as acids or
bases. For example, the hydrated aluminium ion, of aluminium chloride, may donate a proton to a
water molecule, thereby acting as an acid.
[Al(H2O)6]3+ + H2O <===> [Al(H2O)5OH]2+ + H3O+
The carbonate ion of sodium carbonate may accept a proton from a water molecule, thereby acting
as a base.
CO32- + H2O <===> HCO3-+ OHIn pure water, there is a small, but measurable concentration of hydronium and hydroxide ions.
Consideration of the above equation shows that, in pure water, the concentration of the hydronium
ions is exactly equal to that of the hydroxide ions.
Since the ionization of water is very slight, the concentration of water is very much greater
than that of the hydronium and hydroxide ions, and is therefore virtually constant. Thus the
equilibrium constant can be simplified to give a new constant, Kw, known as the ionic product for
water:
Kw = [H3O+][OH-]
A solution which has a higher concentration of hydronium ions than pure water is said to be
acidic; a solution with a higher concentration of hydroxide ions than pure water is said to be basic.
Since the product of the concentration of hydronium and of hydroxide ions is a constant, at a given
temperature, the concentration of only one of theses ions need be stated to indicate the
concentration of both. It is usually the concentration of hydronium ions which is used to specify the
acidity or basicity of a solution and because the values may be very low, and subject to great
variation, the pH scale is used.
Pre-laboratory work
1. Define 'pH' of a solution.
2.
How could a basic solution be defined in terms of its pH?
Material: one pH meter for one bench
1.0 M sodium ethanoate solution, 0.1 M aluminium chloride solution,
0.10 M iron(III) chloride solution, 0.10 M iron(II) chloride solution,
0.10 M magnesium sulphate solution, 0.10 M sodium sulphate solution,
0.10 M tin(IV) chloride solution
5 cm3 graduated pipette(8), 50 cm3 measuring cylinder (4),
50 cm3 pipettes (4), one additional 100 cm3 beaker for each student
50 cm3 burette (4)
Students should label the above apparatus before you start the experiment.
Chemicals per group
F.6/7 Chem.Prac./AB3_saltSolutions/p.1(2)
0.10 M sodium hydroxide solution (75 cm3), 0.10 M ethanoic acid (75 cm3),
0.10 M hydrochloric acid (75 cm3)
Procedure
1. Determine the pH of a 1.00 M aqueous sodium ethanoate solution, using a pH meter. Wash
the bulb of the pH electrode in distilled water immediately after use.
2. Determine the pH of 0.10 M aqueous solutions of the following substances,
Aluminium chloride, magnesium sulphate, iron(III) chloride, sodium
sulphate, iron(II) chloride and tin(IV) chloride
recording your results in a tabular form and make out an order of decreasing acidity.
3. Prepare a buffer solution (solution A) as follows: Pipette 25.0 cm3 of 0.10 M sodium
hydroxide solution into a 100 cm3 beaker and add 50.0 cm3 of 0.10 M ethanoic acid. Mix
thoroughly and determine the pH of the solution.
4. Prepare a solution of the same pH, which does not act as a buffer (solution B), as follows:
Place about 25 cm3 of 0.10 M sodium hydroxide solution in a 100 cm3 beaker and run in
0.10 M hydrochloric acid from a burette until the pH of this solution is approximately the
same as that of the buffer solution, prepared in step 3.
5. Compare the behavior of solutions prepared in step 3 and 4.
(a) By using 10 cm3 measuring cylinder take 5 cm3 of each solution in two separate 100
cm3 beakers. Dilute the solution to about 50 cm3 by adding distilled water. Determine
the pH of the two diluted solutions.
(b) By using 100 cm3 measuring cylinder transfer 25 cm3 of each solution in two separate
100 cm3 beaker. Add 0.5 cm3 of 0.10 M hydrochloric acid in each of the beaker, using
5 cm3 graduated pipette. Determine the pH of each solution. Add a further 4.0 cm3 of
acid to each of the solutions and again determine the pH in each case.
(c) Repeat experiment (b) but add 0.10 M sodium hydroxide instead of the hydrochloric
acid.
Results for step 3,4,5
pH value
Original
(a)
(b)
(c)
value
+ 0.5 cm3 + 4.5 cm3
+ 0.5 cm3 + 4.5 cm3
Solution A
Solution B
Questions for discussion
3. Explain the behaviour of the sodium ethanoate in water.
4.
5.
6.
7.
Would the pH of the solution be different if you used (a) potassium ethanoate, (b) ammonium
ethanoate?
Will the chloride or sulphate ion have much influence on the pH of the solution?
Relate the acidity of the solution to the charge on the metal ion.
Define the term 'buffer solution', incorporating the characteristics of the buffer solution you
have prepared and studied.
END
F.6/7 Chem.Prac./AB3_saltSolutions/p.2(2)