How do acids and alkalis
react?
Demonstration
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.
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.
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:
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+, SO42-
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) + SO42-(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.