CIE Syllabus focus:
'Interpret given data for melting and boiling temperatures of Period 2 and Period 3 elements, explaining trends in terms of structure and bonding.'
Melting and boiling temperature trends across Periods 2 and 3 show how strongly particles attract each other. To explain the data, always identify the element’s structure first, then the bonding or intermolecular forces involved.
How to interpret the trend
Across both periods, the pattern is not random. Values usually rise through the metallic elements, reach a maximum at an element with a giant covalent structure, and then fall sharply for simple molecular elements and the noble gas.
A strong answer links each part of the trend to the kind of structure present:
Metallic structure: strong attraction between positive ions and delocalized electrons
Giant covalent structure: many strong covalent bonds throughout a network
Simple molecular structure: weak intermolecular forces between molecules
Monatomic noble gas: only very weak forces between single atoms
Boron, carbon, and silicon are especially important because they do not fit the same pattern as the surrounding metals or molecules.
Boron, carbon, and silicon have a giant covalent structure.
Giant covalent structure: A continuous network of atoms joined by strong covalent bonds throughout the whole structure.
Because many strong covalent bonds must be overcome, these elements have very high melting and boiling temperatures.
Period 2: lithium to neon
Lithium to carbon
From lithium to beryllium, melting and boiling temperatures increase because metallic bonding becomes stronger. Beryllium forms smaller, more highly charged ions than lithium and contributes more delocalized electrons, so the attraction in the metallic lattice is stronger.
Boron has a much higher temperature than the metals before it because it is giant covalent, not metallic. A large amount of energy is needed to overcome the strong covalent bonding in its structure.
Carbon has the highest melting and boiling temperatures in Period 2. This is also because of its giant covalent structure. The strong covalent bonds extend throughout the lattice, so a very large amount of energy is required.
Nitrogen to neon
There is a very sharp drop from carbon to nitrogen because nitrogen exists as simple molecules. When nitrogen melts or boils, the strong covalent bond inside each molecule is not broken; only the attractions between molecules are overcome.
These attractions are London dispersion forces.
London dispersion forces: Weak intermolecular attractions caused by temporary dipoles that induce dipoles in nearby particles.
Oxygen and fluorine are also simple molecular substances, so their melting and boiling temperatures remain low. Neon is monatomic, so it also has low values.
Among these simple molecular and monatomic elements, differences are much smaller than those seen for metals or giant covalent structures. In general, larger electron clouds give stronger London dispersion forces, but the exact order of the melting temperatures can be irregular because packing in the solid also matters.
Period 3: sodium to argon
Sodium to silicon
From sodium to magnesium to aluminum, melting and boiling temperatures generally increase because metallic bonding becomes stronger across the metals. The metal ions have higher charge, and there are more delocalized electrons, so the electrostatic attraction in the lattice is stronger.
Silicon has a very high melting and boiling temperature because it has a giant covalent structure. As with carbon, many strong covalent bonds must be overcome, so much more energy is needed than for the metals or molecular elements.
Phosphorus to argon
There is a large drop from silicon to phosphorus because phosphorus is simple molecular, existing as molecules. The same idea applies as in Period 2: melting or boiling only overcomes intermolecular forces, not the covalent bonds within each molecule.
Sulfur has higher melting and boiling temperatures than phosphorus because sulfur exists as molecules, which are larger and have more electrons. This gives stronger London dispersion forces between sulfur molecules.
Chlorine exists as molecules, which are smaller than and , so its London dispersion forces are weaker and its temperatures are lower.
Argon is monatomic. It has only weak attractions between individual atoms, so it has the lowest melting and boiling temperatures in Period 3.
Key comparisons and exam focus
The most important skill is to connect the shape of the graph or data trend to a change in structure.
A steady rise among metals suggests strengthening metallic bonding.
A very high peak usually indicates a giant covalent element such as carbon or silicon.
A sudden drop after that peak shows a change from a giant structure to simple molecules.
Low values for molecular elements and noble gases show that only weak intermolecular forces are being overcome.
When comparing Period 2 with Period 3, the overall pattern is similar:
metals first
then a giant covalent maximum
then simple molecular elements
ending with a noble gas
However, the exact values depend on the specific structure and particle size. For example, the larger molecular size of makes sulfur stand out among the Period 3 nonmetals.
In exam answers, avoid saying that simple molecular substances have low temperatures because they have “weak covalent bonds.” The covalent bonds inside the molecules are strong. The low melting and boiling temperatures are due to weak forces between molecules, not weak bonds within molecules.
Practice Questions
Silicon has a much higher melting temperature than phosphorus. Explain why. [2 marks]
Silicon has a giant covalent structure / giant covalent lattice. (1)
Phosphorus is simple molecular / exists as molecules, so silicon requires much more energy to overcome strong covalent bonds, while phosphorus only requires intermolecular forces to be overcome. (1)
Describe and explain the trend in boiling temperatures across Period 3 from sodium to argon. [6 marks]
Boiling temperatures increase from sodium to magnesium to aluminum because metallic bonding becomes stronger. (1)
Stronger metallic bonding is due to higher ionic charge and more delocalized electrons. (1)
Silicon has a very high boiling temperature because it has a giant covalent structure. (1)
There is a large drop from silicon to phosphorus because phosphorus is simple molecular, so only intermolecular forces are overcome. (1)
Sulfur has a higher boiling temperature than phosphorus because molecules are larger / have more electrons, so London dispersion forces are stronger. (1)
Chlorine has a lower boiling temperature than sulfur because molecules are smaller, and argon is lowest because it is monatomic with very weak intermolecular forces. (1)
FAQ
Boiling mainly depends on how much energy is needed to separate particles completely in the liquid.
Melting can be less regular because the way molecules pack in a solid crystal also affects the energy needed. Two substances with similar intermolecular forces can have noticeably different melting temperatures if their solids pack differently.
Both have giant covalent structures, but the bonds are not equally strong.
Carbon forms shorter covalent bonds.
Shorter bonds usually mean stronger bonding.
Silicon atoms are larger, so the overlap between bonding orbitals is less effective.
As a result, less energy is needed to overcome the bonding in silicon than in carbon.
Carbon exists in different allotropes, mainly graphite and diamond, and they do not have identical physical data.
Also, at very high temperatures carbon may sublime rather than form a normal liquid under standard conditions. Because of this, different sources may quote slightly different values or may report a sublimation temperature instead of a true melting temperature.
Room temperature lies between their different melting and boiling temperatures.
Sulfur exists as relatively large $S_8$ molecules, so the attractions between molecules are strong enough for sulfur to remain solid at room temperature.
Chlorine exists as much smaller $Cl_2$ molecules, so the intermolecular forces are weaker. That means chlorine has a much lower boiling temperature and is a gas at room temperature.
Sulfur can exist in more than one solid form, called allotropes, such as rhombic sulfur and monoclinic sulfur.
These forms have slightly different crystal arrangements, which can lead to small differences in measured melting data. Data books may also round values differently or quote measurements taken under slightly different conditions.
For exam purposes, the key idea is unchanged: sulfur has higher temperatures than phosphorus and chlorine because its molecules are larger and intermolecular forces are stronger.
