How do acids and alkalis react?

How do acids and alkalis react?

Demonstration

This activity links the neutralisation of an acid by an alkali to the changes in ionic concentrations that result from the reaction between hydrogen ions and hydroxide ions.
 

Lesson organisation

The practical involves a mildly toxic alkaline solution, barium hydroxide. It needs careful preparation and manipulation, involving a conductimetric titration. A clear demonstration is more likely to lead to successful learning than a class activity that will need a lot of teacher support.

A good demonstration should take about 25 mins.

Chemicals

Each demonstration requires:

Barium hydroxide solution, 0.10 M, (HARMFUL, IRRITANT), about 200 cm³ 

Dilute sulfuric acid, 1.0 M, (IRRITANT), about 25 cm³

Purified (distilled or deionised) water

Phenolphthalein indicator solution (HIGHLY FLAMMABLE)

Apparatus

Eye protection for the teacher and for any members of class who assist at the bench

Test-tube (150 x 25 mm)

Test-tube rack

Beakers (100 cm³), 2

Measuring cylinder (100 cm³)

Glass rod

Burette (50 cm³)

Clamps (2) and stand

Small funnel (for filling the burette)

White tile (for standing beaker on during titration) and white card background (for class visibility)

Pair of carbon electrodes (Note 1) in holder, with 4 mm plug adapters

Plug leads (4 mm plug at each end), 4

Bulb (12 V) in holder

AC demonstration ammeter

Low voltage AC supply (Note 2) 

Health & Safety and Technical notes

Wear eye protection. 

Barium hydroxide solution, Ba(OH)₂(aq), (HARMFUL, IRRITANT at concentration used) - see CLEAPSS Hazcard. Solid barium hydroxide (CORROSIVE) contains water of crystallisation, and reacts with carbon dioxide from the air while in storage.  It is only slightly soluble in water (maximum 4 g in 100 cm³) but the solution is much more alkaline than limewater. Make 250 cm³ of the solution. Purified water should be boiled to remove carbon dioxide before being added to solid barium hydroxide. Once prepared the solution is very sensitive to carbon dioxide and immediately goes cloudy (barium carbonate) when exposed to the atmosphere. All fresh solutions, which will be alkaline and may irritate sensitive skin, must be protected with a soda lime guard tube. For all these reasons, it is important to check the concentration of the solution before the demonstration to ensure it is reasonably close to 0.10 M – an exact concentration is not necessary. To do this titrate 50 cm³ of the solution with the dilute sulfuric acid (1.0 M), to be used in the experiment, using phenolphthalein indicator (two drops). A titre value between 4 cm³ and 6 cm³ is acceptable.

Dilute sulfuric acid, H₂SO₄(aq), (IRRITANT at concentration used) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.

Phenolphthalein indicator solution (HIGHLY FLAMMABLE) - see CLEAPSS Hazcard and CLEAPSS Recipe Book. Phenolphthalein indicator solution should be provided in a dropper bottle.

1 The two carbon electrodes need to be mounted securely in a holder that keeps them parallel. If available, a 4 mm plug adapter should be fitted to the top of each electrode. If not available, crocodile clips can be used instead, but you will need a cardboard or plastic separator between the clips to avoid accidental short-circuiting.

2 The low voltage power supply should be a variable low voltage unit capable of supplying alternating current (AC) at about 12 V when connected through a 12 V bulb and the electrodes dipped in solution in series. A demonstration AC ammeter should be included if available, and if the 12 V bulb fails to light brightly enough for the class to see when tested as in Stage 1 below, the ammeter is essential.

 

Procedure

Stage 1
a Mix equal volumes of dilute sulfuric acid and barium hydroxide solution in a test-tube to observe what happens.

b Add 50 cm³ of barium hydroxide solution to one beaker, and add 2–4 drops of phenolphthalein indicator solution to show the solution is alkaline.

c Dip the electrodes into this solution to demonstrate that it conducts electricity. The AC supply should ensure there is no electrolysis.

d Rinse the electrodes with purified water. Now test the sulfuric acid in a second beaker to show that it conducts electricity.

Stage 2
e Fill the burette with 1.0 M sulfuric acid to the zero mark. Fix the burette securely over the beaker containing 50 cm³ of 0.1 M barium hydroxide solution, ready for titration.

f Clamp the electrode assembly firmly at one side of the beaker so the electrodes dip into the full depth of the solution, and connect it to the rest of the test circuit. Place the stirring rod in the solution.

g Switch on the supply, note the ammeter reading and the bulb brightness.

h Add sulfuric acid from the burette, 0.5 cm³ at a time, with stirring. After each addition, note the ammeter reading and the bulb brightness, and look for a permanent change in the indicator colour.

i At the indicator end-point note the volume of acid added, the ammeter reading and the bulb brightness.

j Continue to add portions of acid and note the ammeter reading and bulb brightness until the change on further addition is minimal.

Teaching notes

This experiment provides important evidence for the simple hydrogen ion theory of acidity, and for the ionic nature of the neutralisation reaction with hydroxide ions. Thus it forms a natural part of a sequence of experiments in which this theoretical model can be built up for students. This experiment is not likely to be useful on its own.

It also depends on students’ understanding of ionic theory in general, and their appreciation that the conductivity of a solution depends on the concentration of ions in the solution. Alternating current is used rather than direct current to avoid electrolysis taking place, at least to any extent that would affect the outcome of the experiment.

Although this is likely to be a demonstration for most students, some teachers may wish to use it as a class experiment, possibly with older students. Safety issues are relatively minor, essentially using dilute sulfuric acid (1.0 M) and barium hydroxide solution (0.1 M). The latter should be treated as HARMFUL for students, even at this low concentration.

The reason for using such different concentrations in a conductimetric titration is to minimise the decrease in conductivity caused by increasing the volume of water present as the titration proceeds, apart from the changes in the total number of ions present.

Here are answers to the questions
1 This revises the students’ understanding of ions, and their ability to identify the ions that are present (barium cations and hydroxide anions in barium hydroxide, and hydrogen cations and sulfate anions in sulfuric acid). The symbols for these ions are:

Ba²⁺, OH^-, H⁺, SO₄^2-

2 The ions are the current carriers in solution, and conductivity depends on the concentration of ions. In asking which ions are removed in the formation of barium sulfate, there is an opportunity to write the ionic equation for the precipitation reaction

Ba²⁺(aq) + SO₄^2-(aq) → BaSO₄(s)

and hence to identify a fall in the total ion concentration which is reflected in a fall in conductivity.

3 Having dealt with the barium and sulfate ions, the students are then in a position to focus on what else is happening, starting with identifying hydrogen cations and hydroxide anions as potential reactants. These ions do react and the reaction is simple.

4 The ionic equation for the reaction between these ions is:

H⁺(aq) + OH^-(aq) → H₂O(l)

5 Any reactant added once an end-point has been reached is ‘surplus’, and so ions are again present and the conduction of the solution rises again.

It is unlikely that the end-point will be marked by zero conductivity. The titration method used is too crude to find the exact point at which almost no ions are present, and even a drop of sulfuric acid added nearly at the end-point will take the reaction past the end-point.