Investigating Melting Points and Bond Types

Candle wax softens in a warm hand. Table salt sits unchanged in a flame that would melt a candle in seconds. Both are white solids, both are common, and both are made of well-understood particles, yet their melting points are separated by roughly 700 °C. This practical turns that contrast into data. You measure (or bracket) the melting points of five solids with safe electrical heating, then explain the spread using the forces that must be overcome in each one.

Purpose

To compare the melting behaviour of a set of solids using an electric water bath and a melting-point apparatus, and to use the results, together with data-book values, to decide whether melting each substance breaks intermolecular forces, ionic bonds, covalent bonds or metallic bonds.

The principle

Melting happens when particles gain enough energy to leave their fixed positions and slide past one another. The melting point therefore reflects how strongly the particles are held, and that depends on what the particles are:

• Simple molecular solids (wax, stearic acid, sugar). The covalent bonds inside each molecule are strong, but melting doesn’t break them. It only overcomes the much weaker forces between molecules. Melting points are low: usually below about 300 °C, and often below 100 °C.
• Ionic solids (sodium chloride). Every ion is attracted to all its oppositely charged neighbours in a giant lattice. Melting has to loosen a huge number of strong electrostatic attractions, so melting points are high: sodium chloride melts at 801 °C.
• Giant covalent solids (silicon dioxide in sand). Melting means breaking strong covalent bonds that run through the entire structure. Melting points are very high: silica melts at roughly 1,700 °C.
• Metallic solids. Positive ions are held in a sea of delocalised electrons. Melting points vary widely, from below room temperature to above 3,000 °C.

Within the simple molecular group, bigger molecules and molecules that can hydrogen bond usually melt higher. Sugar (sucrose) has many –OH groups and hydrogen bonds strongly, which is part of why it melts well above wax, whose chains attract each other only through London forces. There is a fuller discussion in how bonding affects melting point.

Equipment

• Electric water bath or a beaker of water on a hot plate (no flame), with a thermometer (0–100 °C)
• Electric melting-point apparatus with capillary tubes (most school models reach about 300 °C)
• Boiling tubes and a test-tube rack
• Spatulas, a watch glass, and a mortar and pestle for grinding
• Heat-resistant mat and tongs or a test-tube holder
• Eye protection

Samples (all school-safe)

• Paraffin wax (small shavings)
• Stearic acid (octadecanoic acid), a few grams
• Sucrose (granulated sugar), finely ground
• Sodium chloride (table salt)
• Clean dry sand (silicon dioxide)

Not used in this method: sulfur, iodine, naphthalene, lead or lead compounds, or any metal that would need strong heating. These appear below only as data-book values. They give off harmful vapours or need temperatures and equipment beyond a student bench.

Method

Part A: the water bath (up to about 90 °C)

1. Set the water bath to about 90 °C, or heat a beaker of water on a hot plate. Do not use a Bunsen burner for this part: wax vapour is flammable.
2. Put a small pea-sized amount of each solid into a separate labelled, dry boiling tube: wax, stearic acid, sugar, salt, sand.
3. Stand the tubes in the water bath with a thermometer in the water.
4. Watch as the temperature rises. Record the temperature at which each sample starts to melt (becomes glassy or liquid at the edges) and when it is completely liquid.
5. Wax and stearic acid melt in this range. Sugar, salt and sand do not change at all. Record “no change up to 90 °C” for them.
6. Optional: remove the stearic acid tube, put a thermometer into the liquid, and record the temperature every 30 seconds as it cools. A flat section (the plateau) on the graph marks the freezing point, which is the same temperature as the melting point.

Part B: the melting-point apparatus (up to about 300 °C)

1. Grind a little sugar finely. Press the open end of a capillary tube into the powder and tap it down until you have 2–3 mm of packed solid. Do the same for salt and sand.
2. Place the capillaries in the apparatus and heat quickly to about 150 °C, then slowly (around 2 °C per minute).
3. Watch through the lens. Record the temperature at which the sugar begins to melt, and note any browning.
4. Continue to the apparatus’s safe maximum. Salt and sand will show no change. Record “does not melt below [maximum temperature]”.
5. Switch off and let the block cool before removing the capillaries. They stay hot for a long time.

Why not a Bunsen for salt and sand?

It’s tempting to “prove” that salt has a high melting point by heating it in a test tube over a Bunsen burner. It doesn’t work. Heated in a test tube or on a metal lid, the salt never gets close to 801 °C. It may crackle and jump as trapped water escapes from the crystals, but it won’t melt, and glass tubes can crack or soften under that kind of heating. A result of “didn’t melt in the Bunsen” is still useful evidence that the melting point is very high, but electrical apparatus gives cleaner, safer data. The same goes for sand.

Risk assessment notes

Hazard	Risk	Control
Hot water and hot glass	Scalds and burns	Use tongs or a holder; place tubes on a heat-resistant mat; don’t overfill the bath
Hot melting-point block	Burns long after switching off	Leave to cool; don’t touch the metal block
Molten wax	Burns; wax is flammable	No flames nearby; electric heating only
Stearic acid	Mild eye and skin irritant	Eye protection; wash hands after
Sugar heated strongly	Caramelises, then chars and smokes	Heat only in the capillary apparatus; ventilate the room
Broken capillaries	Cuts	Handle gently; dispose of in the glass bin

Follow your school’s risk assessment, which may change amounts or equipment.

Expected results

Substance	Structure	Observation in this lab	Approximate melting point
Paraffin wax	Simple molecular (long hydrocarbon chains)	Softens and melts over a range in the water bath	About 50–65 °C, depending on grade
Stearic acid	Simple molecular	Melts sharply in the water bath	About 69 °C
Sucrose	Simple molecular, extensive hydrogen bonding	No change in water bath; melts in apparatus with browning	About 185 °C (breaks down as it melts)
Sodium chloride	Ionic lattice	No change in either piece of equipment	801 °C (data book)
Sand (silicon dioxide)	Giant covalent	No change in either piece of equipment	About 1,700 °C (data book)

Notice that wax melts over a range rather than at one temperature. Commercial wax is a mixture of chains of different lengths, and mixtures melt over a range. Stearic acid, being a fairly pure single compound, gives a sharper result. The same idea is used to check purity in organic chemistry; see melting point determination.

Sugar is a useful warning. It does melt at around 185 °C, but it immediately starts to decompose, turning yellow and then brown as caramel forms. A melting point that comes with a colour change is really a decomposition point, and it depends on how fast you heat.

Data-book values to extend the comparison

These were not heated in the method; they are for discussion only.

Substance	Type	Melting point
Mercury	Metallic	−38.83 °C
Sodium	Metallic	97.80 °C
Iodine	Simple molecular (I₂)	113.70 °C
Sulfur	Simple molecular (S₈)	115.21 °C
Lead	Metallic	327.46 °C
Aluminium	Metallic	660.3 °C
Copper	Metallic	1,084.62 °C
Silicon	Giant covalent	1,414 °C
Iron	Metallic	1,538 °C
Tungsten	Metallic	3,422 °C

The metals refuse to follow one simple rule: mercury is a liquid at room temperature and sodium melts below the boiling point of water, while tungsten survives temperatures that would melt almost anything. Metallic melting points depend on how many electrons each atom contributes to the delocalised sea, the size of the ions and the way they are packed.

Analysis questions

1. Explain why stearic acid melts below 100 °C even though it contains strong C–C and C–H covalent bonds.
2. Sucrose and stearic acid are both simple molecular. Suggest why sucrose melts at a much higher temperature.
3. Why does paraffin wax melt over a range of temperatures?
4. Put the five tested substances in order of melting point and label each with the type of force broken on melting.
5. Sodium chloride and silicon dioxide both failed to melt in the apparatus. Use data-book values to explain which has the stronger overall structure, and why.
6. A student claims that “covalent substances have low melting points”. Use your results to criticise this statement.
7. Suggest why silicon’s melting point is so much higher than sulfur’s, though both are non-metals next to each other in Period 3.

Answer pointers

1. Melting separates whole molecules; only weak intermolecular forces are overcome. The covalent bonds within each molecule stay intact.
2. Sucrose has many –OH groups, so its molecules form many hydrogen bonds with one another, which are stronger than the London forces between stearic acid molecules. (Stearic acid has one –COOH group that can hydrogen bond, but most of the molecule is a hydrocarbon chain.)
3. It is a mixture of hydrocarbons of different chain lengths, each melting at a slightly different temperature.
4. Wax < stearic acid < sucrose (intermolecular forces) < sodium chloride (ionic bonds) < silicon dioxide (covalent bonds).
5. Silicon dioxide, about 1,700 °C against 801 °C. Melting it breaks a continuous network of strong Si–O covalent bonds.
6. It is only true for simple molecular covalent substances. Giant covalent substances like sand have some of the highest melting points known.
7. Silicon has a giant covalent structure, so melting breaks covalent bonds. Sulfur is made of separate S₈ molecules held by weak intermolecular forces.

Sources of error and limitations

• Heating too fast. The thermometer or sensor lags behind the sample, so readings come out high. Slow down near the expected melting point.
• Thermometer position. In the water bath the thermometer measures the water, not the solid inside the tube, which is slightly cooler while it is warming.
• Judging the start of melting. “Softening” and “melting” can look similar for wax. Agree a definition, such as the first visible liquid.
• Impurities. Impure samples melt lower and over a wider range.
• Decomposition. For sugar, what you see depends on heating rate, and the result is not a true melting point.
• Upper limit of equipment. For salt and sand you can only say the melting point is above the maximum temperature of the apparatus. The actual values come from data books.

To combine this with a second line of evidence, run testing conductivity to classify bonding on the same substances. The two sets of results together are much more convincing than either alone. For a broader structural comparison, see simple molecular vs giant covalent.

Key takeaways

• Melting point tells you how strongly particles are held, which depends on the type of bonding.
• Simple molecular solids melt low because only intermolecular forces are overcome; hydrogen bonding and larger molecules raise the value.
• Ionic and giant covalent solids melt very high because strong bonds throughout the structure must be broken.
• Metals span an enormous range, from mercury at −38.83 °C to tungsten at 3,422 °C.
• Use electric water baths and melting-point apparatus. Salt won’t melt over a Bunsen, and hazardous substances such as sulfur, iodine and lead compounds stay as data-book values.