18.3 Electrolytes, ionisation and conductivity | Reactions in aqueous solution

18.3 Electrolytes, ionisation and conductivity (ESAFP)

You have learnt that water is a polar molecule and that it can dissolve ionic substances in water. When ions are present in water, the water is able to conduct electricity. The solution is known as an electrolyte.

Electrolyte: An electrolyte is a substance that contains free ions and behaves as an electrically conductive medium.

Because electrolytes generally consist of ions in solution, they are also known as ionic solutions. A strong electrolyte is one where many ions are present in the solution and a weak electrolyte is one where few ions are present. Strong electrolytes are good conductors of electricity and weak electrolytes are weak conductors of electricity. Non-electrolytes do not conduct electricity at all. Conductivity in aqueous solutions, is a measure of the ability of water to conduct an electric current. The more ions there are in the solution, the higher its conductivity. Also the more ions there are in solution, the stronger the electrolyte.

Factors that affect the conductivity of electrolytes (ESAFQ)

The conductivity of an electrolyte is therefore affected by the following factors:

The concentration of ions in solution. The higher the concentration of ions in solution, the higher its conductivity will be.
The type of substance that dissolves in water. Whether a material is a strong electrolyte (e.g. potassium nitrate, \(\text{KNO}_{3}\)), a weak electrolyte (e.g. acetic acid, \(\text{CH}_{3}\text{COOH}\)) or a non-electrolyte (e.g. sugar, alcohol, oil) will affect the conductivity of water because the concentration of ions in solution will be different in each case. Strong electrolytes form ions easily, weak electrolytes do not form ions easily and non-electrolytes do not form ions in solution.
Temperature. The warmer the solution, the higher the solubility of the material being dissolved and therefore the higher the conductivity as well.

Electrical conductivity

Aim

To investigate the electrical conductivities of different substances and solutions.

Apparatus

Solid salt (\(\text{NaCl}\)) crystals
different liquids such as distilled water, tap water, seawater, sugar, oil and alcohol
solutions of salts e.g. \(\text{NaCl}\), \(\text{KBr}\), \(\text{CaCl}_{2}\), \(\text{NH}_{4}\text{Cl}\)
a solution of an acid (e.g. \(\text{HCl}\)) and a solution of a base (e.g. \(\text{NaOH}\))
torch cells
ammeter
conducting wire, crocodile clips and 2 carbon rods.

Method

Set up the experiment by connecting the circuit as shown in the diagram below. In the diagram, X represents the substance or solution that you will be testing.
When you are using the solid crystals, the crocodile clips can be attached directly to each end of the crystal. When you are using solutions, two carbon rods are placed into the liquid and the clips are attached to each of the rods.
In each case, complete the circuit and allow the current to flow for about 30 seconds.
Observe whether the ammeter shows a reading.

Results

Record your observations in a table similar to the one below:

Test substance	Ammeter reading

What do you notice? Can you explain these observations?

Conclusions

Solutions that contain free-moving ions are able to conduct electricity because of the movement of charged particles. Solutions that do not contain free-moving ions do not conduct electricity.

Remember that for electricity to flow, there needs to be a movement of charged particles e.g. ions. With the solid \(\text{NaCl}\) crystals, there was no flow of electricity recorded on the ammeter. Although the solid is made up of ions, they are held together very tightly within the crystal lattice and therefore no current will flow. Distilled water, oil and alcohol also don't conduct a current because they are covalent compounds and therefore do not contain ions.

The ammeter should have recorded a current when the salt solutions and the acid and base solutions were connected in the circuit. In solution, salts dissociate into their ions, so that these are free to move in the solution. Look at the following examples:

Dissociation of potassium bromide:

\[\text{KBr (s)} \rightarrow \text{K}^{+}\text{(aq)} + \text{Br}^{-}\text{(aq)}\]

Dissociation of table salt:

\[\text{NaCl (s)} \rightarrow \text{Na}^{+}\text{(aq)} + \text{Cl}^{-}\text{(aq)}\]

Ionisation of hydrochloric acid:

\[\text{HCl (l)} + \text{H}_{2}\text{O (l)} \rightarrow \text{H}_{3}\text{O}^{+}\text{(aq)} + \text{Cl}^{-}\text{(aq)}\]

Dissociation of sodium hydroxide:

\[\text{NaOH (s)} \rightarrow \text{Na}^{+}\text{(aq)} + \text{OH}^{-}\text{(aq)}\]

temp text

18.2 Ions in aqueous solution

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18.4 Precipitation reactions