Testing Conductivity to Classify Bonding

You can’t see a chemical bond, but you can see a bulb light up. That is the whole idea behind this practical. By pushing a small current through a set of unknown solids, and then through the same substances dissolved in water, you collect evidence that points to how each one is bonded. It takes one lesson, uses cheap equipment, and gives students a result table they can reason from rather than simply memorise.

Purpose

To test a range of solids and solutions for electrical conductivity using a low-voltage circuit, and to use the pattern of results to classify each substance as metallic, ionic, simple molecular or giant covalent.

The principle behind the test

An electric current is a flow of charge. A substance can only carry a current if it contains charged particles that are free to move. There are two kinds of charged particle you will meet in this lab:

• Delocalised electrons. Metals have a shared pool of outer electrons that can drift through the whole structure. Graphite, a form of carbon, also has delocalised electrons, spread across its flat layers.
• Ions. Ionic compounds are made of positive and negative ions. In the solid they are locked in a lattice and cannot travel. When the compound is dissolved in water (or melted), the ions are released and can move towards the electrodes.

Simple molecular substances, such as sugar and wax, are built from neutral molecules. Neither the solid nor a solution of sugar contains mobile charges, so neither conducts. Most giant covalent substances, such as sand (silicon dioxide), have every outer electron fixed in a bond, so they don’t conduct either. Graphite is the famous exception.

This gives a simple decision tree:

1. Does the solid conduct? If yes, it is almost certainly metallic (or graphite).
2. If not, does it conduct when dissolved in water? If yes, it is ionic (it has released mobile ions).
3. If it dissolves but the solution doesn’t conduct, it is simple molecular.
4. If it neither conducts nor dissolves, and it is very hard with a high melting point, it is likely giant covalent.

The full background is in which substances conduct electricity, and why, and the four families of solid are compared in four types of solids.

Equipment

• Low-voltage DC power supply set to 6 V, or a battery pack of four 1.5 V cells
• A 6 V filament bulb in a holder, or an LED with a suitable series resistor (an LED is more sensitive to small currents)
• An ammeter (0–1 A, or a milliammeter for weak currents) as an alternative or extra indicator
• Three connecting leads and two crocodile clips
• Two carbon (graphite) rod electrodes held in a bung or a plastic electrode holder
• Small beakers (50–100 cm³)
• Distilled or deionised water, a wash bottle and paper towels
• Spatulas and a stirring rod
• Eye protection

Samples

• Solids to test directly: a strip of copper, a strip of aluminium foil, a graphite rod or a thick pencil “lead”, a lump of paraffin wax, a sugar cube or crystals, a large crystal or compressed block of table salt (sodium chloride), dry sand
• Substances to dissolve: table salt, sugar, and a little sand for comparison

Never connect any part of this circuit to a mains socket, and never use a mains-powered device other than a proper school low-voltage supply. Everything here runs at 6 V DC or less.

Method

Part A: testing the solids

1. Build a series circuit: power supply, bulb (or LED and resistor), ammeter, and two crocodile clips on the free ends of the leads.
2. Check the circuit works by touching the two clips together briefly. The bulb should light. Then separate them.
3. Clip or press both crocodile clips onto the first solid, about 2 cm apart. Switch on for a few seconds.
4. Record whether the bulb lights (or the LED glows) and the ammeter reading.
5. Switch off. Repeat for each solid. For powders like sand or salt crystals, place the sample in a small dry beaker and push both graphite electrodes into it, 1–2 cm apart.
6. Wipe the clips or electrodes dry between samples.

Part B: testing solutions

1. Replace the crocodile clips with the two graphite electrodes.
2. Half-fill a clean beaker with distilled water and test it first. This is your control. It should give little or no current.
3. Add a spatula of salt to fresh distilled water, stir until dissolved, then test. Switch on for no more than about ten seconds.
4. Rinse the electrodes with distilled water and dry them. Repeat with sugar dissolved in distilled water.
5. Stir a spatula of sand into distilled water and test the mixture.
6. Record observations, including anything happening at the electrodes (bubbles, smell).

Keeping the test fair

• Same voltage, same electrode spacing and the same depth of immersion for every solution.
• Roughly the same amount of solute in the same volume of water.
• Rinse the electrodes between tests so no ions carry over from one beaker to the next.

Risk assessment notes

Hazard	Risk	Control
Electricity	Shock from mains	Use only a low-voltage DC supply or batteries (6 V). No mains connections at any point.
Hot components	Bulb and resistor can get warm; a short circuit heats wires	Switch off between tests; don’t hold clips together for long
Electrolysis products	Passing current through salt solution releases small amounts of hydrogen and chlorine at the electrodes	Keep each test short (about ten seconds), work in a well-ventilated room, don’t sniff the beaker. Anyone with asthma should stand back.
Water near electrics	Spills onto the power supply	Keep the supply away from beakers; wipe spills at once
Chemicals	Salt, sugar, sand and wax are low hazard	Eye protection; no tasting in the lab

Your teacher will adapt this to the school’s own risk assessment. Heating substances until they melt is not part of this practical. Molten ionic compounds do conduct, but melting them needs temperatures and materials that belong in a teacher demonstration or a video, not a student bench.

Expected results

Substance	Solid conducts?	Dissolved in water conducts?	Likely bonding / structure
Copper	Yes, bulb bright	Doesn’t dissolve	Metallic
Aluminium foil	Yes, bulb bright	Doesn’t dissolve	Metallic
Graphite rod	Yes, bulb lights (a little dimmer than copper)	Doesn’t dissolve	Giant covalent, with delocalised electrons
Paraffin wax	No	Doesn’t dissolve	Simple molecular
Sugar	No	No (current near zero)	Simple molecular
Table salt	No	Yes, bulb lights; bubbles at both electrodes	Ionic
Sand	No	No; sand doesn’t dissolve	Giant covalent
Distilled water (control)	n/a	Very little current, bulb dark	Simple molecular (a tiny number of ions)

Two results usually surprise people. First, the pencil lead lights the bulb even though carbon is a non-metal. Graphite’s layers of carbon atoms each share one electron per atom in a delocalised system, so electrons can move along the layers. Second, the dry salt crystal doesn’t conduct even though it is full of ions. The ions are there but they are held in place.

A metal like copper conducts so well that you are really only measuring the resistance of the bulb. With an ammeter, graphite typically gives a noticeably lower current, because it is a poorer conductor than the metals.

If tap water is used by mistake, it will conduct slightly, because it contains dissolved ions such as calcium, magnesium and hydrogencarbonate. That is exactly why distilled water is the control.

Analysis questions

1. Explain why solid sodium chloride does not conduct but sodium chloride solution does.
2. Graphite and diamond are both pure carbon. Predict whether diamond would light the bulb, and explain why in terms of electrons.
3. Sugar dissolves in water, yet the solution doesn’t conduct. What does this tell you about the particles sugar forms when it dissolves?
4. Why was distilled water tested before any solutions?
5. Bubbles appeared at the electrodes in salt solution but not in sugar solution. Why is this extra evidence that the salt solution contains ions?
6. A student finds that a grey solid conducts as a solid and is shiny when scratched. Which bonding type is most likely? What further test could rule out graphite?
7. Suggest why an LED is a better indicator than a filament bulb for weakly conducting samples.

Answer pointers

1. In the solid, ions are fixed in the lattice. In solution they are free to move and carry charge.
2. Diamond would not conduct. Each carbon uses all four outer electrons in four covalent bonds, so no electrons are delocalised.
3. Sugar dissolves as whole neutral molecules. It doesn’t split into ions.
4. It is the control: it shows that any current in later tests comes from the dissolved substance, not the water itself.
5. Bubbles show that a chemical change (electrolysis) happens at the electrodes, which requires ions to arrive there and gain or lose electrons.
6. Metallic. Graphite is not malleable and leaves a mark on paper; a metal bends or flattens when hammered.
7. An LED needs a much smaller current to glow visibly, so it shows weak conduction that would leave a bulb dark.

Sources of error and limitations

• Contaminated electrodes or beakers. A trace of salt left on an electrode can make the sugar solution appear to conduct slightly. Rinse with distilled water, not tap water.
• Damp solids. A salt crystal that has absorbed moisture from the air, or has been handled with wet fingers, may carry a small surface current. Keep samples dry.
• Poor contact. Oxide layers on metals (especially aluminium) or loose crocodile clips can make a good conductor look weak. Scratch the surface lightly and check the clips.
• Insensitive indicator. A filament bulb needs a relatively large current. Weak conductors can pass a current too small to light it. The ammeter or LED reveals the difference.
• Varying concentration. Adding different amounts of solute changes the current. The test is qualitative, so exact amounts don’t matter much, but big differences can confuse comparisons between solutions.
• Classification is evidence, not proof. Conductivity alone doesn’t separate every case. Graphite behaves like a metal in this test. Some covalent substances, such as hydrogen chloride, react with water to form ions, so their solutions conduct. Combine this test with others, such as melting points and solubility, which are covered in investigating melting points and bond types.

Extensions

• Use the ammeter to rank solutions by current at the same concentration. Compare salt with a weaker electrolyte provided by your teacher, such as a dilute solution of ethanoic acid (vinegar).
• Vary the distance between the graphite electrodes in salt solution and plot current against separation.
• Research why graphite conducts along its layers much better than across them.

Key takeaways

• A substance conducts only if it has charged particles that are free to move: delocalised electrons or mobile ions.
• Metals conduct as solids. Graphite is the non-metal exception, thanks to delocalised electrons in its layers.
• Ionic compounds don’t conduct as solids but do when dissolved (or molten), because the ions are released.
• Simple molecular substances, like sugar and wax, don’t conduct as solids or in solution.
• Giant covalent substances such as sand don’t conduct and don’t dissolve.
• Keep the circuit at low voltage (6 V DC), keep tests short, and use distilled water as the control.