Periodicity: Melting Points and Conductivity

The melting points and electrical conductivity of the Period 3 elements tell a fascinating story about how structure and bonding change across a period. By examining these trends, we can connect macroscopic physical properties to the microscopic world of atoms, bonds, and intermolecular forces.

Melting Points Across Period 3

The melting points of Period 3 elements vary dramatically:

Element	Na	Mg	Al	Si	P4	S8	Cl2	Ar
Melting Point (degrees C)	98	650	660	1414	44	115	-101	-189
Structure	Metallic	Metallic	Metallic	Giant covalent	Simple molecular	Simple molecular	Simple molecular	Monatomic
Bonding to Overcome	Metallic	Metallic	Metallic	Covalent	London forces	London forces	London forces	London forces

The pattern shows a rise from sodium to silicon, then a dramatic drop at phosphorus, with continued low values through to argon. This reflects three fundamentally different types of structure.

graph LR
    subgraph "Metals (Metallic bonding)"
        Na["Na<br/>98°C"]
        Mg["Mg<br/>650°C"]
        Al["Al<br/>660°C"]
    end
    subgraph "Giant Covalent"
        Si["Si<br/>1414°C"]
    end
    subgraph "Simple Molecular (London forces)"
        P["P4<br/>44°C"]
        S["S8<br/>115°C"]
        Cl["Cl2<br/>-101°C"]
        Ar["Ar<br/>-189°C"]
    end
    Na --> Mg --> Al --> Si --> P --> S --> Cl --> Ar

The Metals: Na, Mg, Al

Sodium, magnesium, and aluminium are metallic. In a metal, a lattice of positive ions is surrounded by a sea of delocalised electrons. The melting point depends on the strength of the metallic bonding.

Factors Determining Metallic Bond Strength

Factor	Na	Mg	Al
Ion charge	+1	+2	+3
Ionic radius (pm)	95	65	50
Delocalised electrons per atom	1	2	3
Charge density (charge/radius)	Low	Medium	High
Melting point (degrees C)	98	650	660

The trend: Metallic bonding strength increases from Na to Al because:

The charge on the metal ion increases (+1 -> +2 -> +3)
The ionic radius decreases (95 -> 65 -> 50 pm)
The number of delocalised electrons per atom increases (1 -> 2 -> 3)
The charge density of the ion increases

All of these increase the electrostatic attraction between the cations and the sea of electrons.

Note: Aluminium's melting point (660 degrees C) is only slightly higher than magnesium's (650 degrees C), despite the much stronger bonding. This is partly because Al has a different crystal structure (face-centred cubic vs hexagonal close-packed for Mg). Crystal structure can affect how efficiently atoms pack and how bonding forces translate into melting point.

Why Not Just "Number of Electrons"?

A common oversimplification is to say "more delocalised electrons = higher melting point." In reality, it is the combination of ion charge, ion size, and number of delocalised electrons that matters. The key concept is electrostatic attraction between positive ions and delocalised electrons -- stronger when charges are higher and distances are shorter.

Silicon: Giant Covalent

Silicon has a giant covalent structure (diamond-type structure). Each silicon atom forms four strong covalent bonds in a tetrahedral arrangement extending throughout the entire crystal.

To melt silicon, you must break many of these strong Si-Si covalent bonds. This requires enormous amounts of energy, giving silicon the highest melting point of any Period 3 element: 1414 degrees C.

Key characteristics of giant covalent structures:

Very high melting and boiling points
Hard solids
Insoluble in water and organic solvents
Generally do not conduct electricity (except graphite, which has delocalised electrons)

Comparison with Other Giant Covalent Structures

Substance	MP (degrees C)	Bond Type	Structure
Diamond (C)	3550	C-C covalent	Tetrahedral, each C bonded to 4 others
Silicon (Si)	1414	Si-Si covalent	Tetrahedral, each Si bonded to 4 others
SiO2 (quartz)	1713	Si-O covalent	Tetrahedral, each Si bonded to 4 O atoms

Silicon has a lower melting point than diamond because Si-Si bonds are longer and weaker than C-C bonds (silicon is larger, so bond overlap is less effective).

Simple Molecular: P4, S8, Cl2, Ar

The remaining elements form simple molecular structures. The molecules are held to each other by London (dispersion) forces -- weak intermolecular forces caused by temporary induced dipoles.

When these substances melt, it is the intermolecular forces that are overcome, not the covalent bonds within molecules. Since London forces are weak, the melting points are low.

The strength of London forces depends on the number of electrons (and hence the size and polarisability) of the molecule:

Substance	Formula	Electrons per Molecule	Melting Point (degrees C)
Sulfur	S8	128	115
Phosphorus	P4	60	44
Chlorine	Cl2	34	-101
Argon	Ar	18	-189

Why Does Molecular Size Matter?

Larger molecules have more electrons and a larger electron cloud that is more polarisable (more easily distorted into a temporary dipole). Greater polarisability leads to stronger London forces:

A temporary dipole forms when electrons randomly shift to one side of the molecule
This temporary dipole induces a dipole in a neighbouring molecule
The two dipoles attract each other
Larger electron clouds produce larger temporary dipoles, creating stronger attractions

This is why S8 (128 electrons) has much stronger London forces than Ar (18 electrons), despite both being non-polar.

The trend among molecular substances: S8 > P4 > Cl2 > Ar, which follows the number of electrons per molecule (128 > 60 > 34 > 18).

Summary of the Overall Trend

The melting point pattern across Period 3 can be summarised:

Na -> Al: Increasing melting points due to stronger metallic bonding (higher charge, more delocalised electrons, smaller ionic radius)
Si: Very high melting point due to strong covalent bonds in a giant covalent structure
P4 -> Ar: Low melting points because only weak London forces hold molecules together; decreases as molecular size decreases

Common Misconception

Misconception: "Silicon has a high melting point because it has strong London forces between its molecules."

Correction: Silicon does NOT exist as simple molecules. It has a giant covalent structure -- a continuous 3D network of Si-Si bonds. There are no individual Si molecules and no intermolecular forces. The high melting point is entirely due to the need to break strong covalent bonds.

Electrical Conductivity Across Period 3

Element	Na	Mg	Al	Si	P4	S8	Cl2	Ar
Conductivity	Good	Good	Very good	Semi-conductor	None	None	None	None
Reason	Delocalised e-	Delocalised e-	Delocalised e-	Band gap	No mobile charges	No mobile charges	No mobile charges	No mobile charges

Metals (Na, Mg, Al)

Good conductors because they have delocalised electrons that are free to move through the lattice and carry charge. Aluminium is a particularly good conductor due to its three delocalised electrons per atom.

Important: Metals conduct in both solid and liquid states because the delocalised electrons remain mobile in both phases.

Conductivity decreases with increasing temperature in metals because increased lattice vibrations scatter the moving electrons, impeding their flow.

Silicon -- Semiconductor

A semiconductor. Pure silicon has very low conductivity at room temperature, but it increases with temperature (opposite to metals). This is because at higher temperatures, more electrons gain enough energy to jump from the valence band to the conduction band, creating mobile charge carriers.

Silicon's conductivity also increases enormously when doped with small amounts of other elements:

n-type doping (e.g., adding phosphorus, which has 5 outer electrons): provides extra free electrons
p-type doping (e.g., adding boron, which has 3 outer electrons): creates "holes" that act as positive charge carriers

Non-metals (P4, S8, Cl2, Ar)

Do not conduct electricity because all electrons are localised -- either in covalent bonds within molecules or as lone pairs. There are no free-moving charged particles to carry current.

Worked Example: Predicting Properties from Structure

Question: Predict which would have the higher melting point: P4 or I2. Both are simple molecular.

Answer: I2 has 106 electrons per molecule, while P4 has 60. Both are held by London dispersion forces. Since I2 has more electrons, it has a larger, more polarisable electron cloud and therefore stronger London forces. I2 should have the higher melting point.

Actual values: P4 = 44 degrees C, I2 = 114 degrees C -- confirming the prediction.

A-Level Deep Dive: Periodicity: Melting Points and Conductivity

Spec mapping

Edexcel 9CH0 specification Topic 1, sub-topic 1.5 (alongside content from Topic 2.7 on metallic and giant covalent structure) covers the trends in melting and boiling points across period 3 elements (Na, Mg, Al, Si, P, S, Cl, Ar) in terms of structure (metallic, giant covalent, simple molecular) and bonding strength; the trend in electrical conductivity across period 3 (high for metals, very high then non-conductor for Si, zero for non-metals); and the relationship between physical properties and underlying structure type (refer to the official Pearson Edexcel specification document for exact wording). This is examined in Paper 1 (Advanced Inorganic and Physical Chemistry) with strong synoptic connections to Topic 2 (Bonding and structure).

Worked example with full mark scheme

Question (8 marks):

The melting points (°C) of period 3 elements are: Na 98, Mg 650, Al 660, Si 1414, P 44, S 119, Cl −101, Ar −189.

(a) Identify the element with the highest melting point and explain its bonding/structure. (3) (b) Explain why P, S, Cl and Ar have low melting points despite their increasing atomic mass. (3) (c) Explain why aluminium has a higher melting point than sodium and magnesium, despite all three being metals. (2)

Solution with mark scheme:

(a) B1 — silicon (Si, m.p. 1414 °C). B1 — Si forms a giant covalent (giant atomic) structure analogous to diamond, with each Si atom covalently bonded to four neighbours in a tetrahedral lattice. B1 — to melt Si requires breaking many strong covalent Si–Si bonds, requiring very high temperatures. Common error: candidates write "Si has strong covalent bonds" without specifying the giant structure. The macroscopic property follows from the combination of strength and extent.

(b) B1 — P, S, Cl and Ar all have simple molecular (or in Ar's case, monatomic) structures: P₄, S₈, Cl₂, Ar. B1 — the only forces between molecules are weak London dispersion forces (induced-dipole induced-dipole interactions). B1 — although these forces increase with molecular size (S₈ has more electrons than P₄, Cl₂ less than P₄, etc.), they are far weaker than the covalent or metallic bonding holding atoms together in earlier elements. Thus melting points are low. Common error: candidates conflate breaking the covalent bonds within molecules with the much weaker intermolecular forces — only the latter break on melting.

(c) B1 — all three are metals with delocalised "sea of electrons", but Al has 3 valence electrons per atom (3s² 3p¹) compared to Mg's 2 (3s²) and Na's 1 (3s¹). B1 — more delocalised electrons per atom → stronger metallic bonding → higher melting point. Al³⁺ also has a smaller radius and higher charge than Mg²⁺ or Na⁺, increasing the electrostatic attraction with the electron sea.

Total: 8 marks (B3 + B3 + B2).

Specimen question modelled on the Edexcel 9CH0 Paper 1 format

Question (5 marks): A student is given samples of Na, Si, S₈ and Ar at room temperature.

(a) Predict the electrical conductivity of each sample as a solid, and explain. (3) (b) Predict the conductivity of molten S₈ and explain. (2)

Mark scheme decomposition by AO:

Part	AO1	AO2	Marks
(a)	1	2	3
(b)	0	2	2
Total	1	4	5

(a) B1 (AO1.1) — Na conducts (high) because of delocalised mobile electrons in metallic bonding. M1 (AO2.1) — Si is a semiconductor (limited conductivity, increasing with temperature) because in the giant covalent lattice all valence electrons are localised in Si–Si bonds, but the band gap is small enough for thermal excitation. A1 (AO2.1) — S₈ and Ar are non-conductors: in S₈ all valence electrons are in covalent bonds within the S₈ ring; in Ar the electrons are tightly held within the atom. No mobile carriers.