Periodicity of physical properties of the elements in Period 3

· Period 3 elements: Na, Mg, Al, Si, P, S, Cl, Ar.
· Key physical properties to know: atomic radius, ionic radius, melting point, electrical conductivity.
· Periodicity = repeating patterns in properties across periods, caused by changes in nuclear charge, electron arrangement, structure and bonding.
· Exam answers must link: structure → bonding/forces → property.

Atomic radius across Period 3

· Atomic radius decreases from Na → Cl.
· Across the period, proton number increases but electrons are added to the same principal shell.
· Shielding remains similar, because no new inner shell is added.
· Increased nuclear charge attracts the outer electrons more strongly.
· Therefore, the outer shell is pulled closer to the nucleus, so atomic radius decreases.
· Argon is often excluded from covalent/atomic radius comparisons because it exists as single atoms and does not form typical covalent bonds.

This graph shows the decrease in atomic radius across Period 3. The trend is caused by increasing nuclear charge with similar shielding, pulling outer electrons closer to the nucleus. Source

Ionic radius across Period 3

· Cations are smaller than their atoms because they have lost the outer shell electrons.
· Na⁺, Mg²⁺, Al³⁺ decrease in radius across the period due to increasing nuclear charge attracting the remaining electrons more strongly.
· Anions are larger than their atoms because they have gained electrons, increasing electron–electron repulsion.
· There is a large jump in radius from Al³⁺/Si⁴⁺ to P³⁻, because the comparison changes from positive ions to negative ions.
· P³⁻, S²⁻, Cl⁻ decrease in radius across the period because nuclear charge increases while the ions have the same electron shell pattern.
· Exam phrase: “ionic radius decreases across isoelectronic ions as nuclear charge increases.”

This graph shows the sharp contrast between small Period 3 cations and much larger Period 3 anions. It helps students explain the jump in ionic radius using loss or gain of electrons and changes in electron–electron repulsion. Source

Melting point trend across Period 3

· Overall pattern: Na < Mg < Al < Si, then a sharp fall to P, S, Cl, Ar.
· Na, Mg and Al have giant metallic lattices.
· Melting point increases from Na → Al because:
· charge on metal ion increases: Na⁺ → Mg²⁺ → Al³⁺
· ionic radius decreases
· number of delocalised electrons per atom increases: 1 → 2 → 3
· metallic bonding becomes stronger.
· Silicon has a giant covalent structure, so many strong covalent bonds must be broken.
· Therefore, Si has the highest melting point in Period 3.
· P, S, Cl are simple molecular substances with weak London forces between molecules.
· Argon is monatomic, with weak London forces between atoms.
· Melting point of simple molecular substances depends on strength of intermolecular forces, not covalent bonds inside molecules.
· Sulfur has a higher melting point than phosphorus and chlorine because S₈ molecules are larger and have more electrons, so they have stronger London forces.

This graph shows the increase in melting point from the metals to silicon, followed by a sharp decrease for simple molecular elements. It is useful for linking the melting point trend to metallic bonding, giant covalent bonding and London forces. Source

Electrical conductivity across Period 3

· Na, Mg and Al conduct electricity because they contain mobile delocalised electrons.
· Conductivity generally increases from Na → Al because the number of delocalised electrons per atom increases.
· Aluminium is the best conductor among Na, Mg and Al because each atom contributes three delocalised electrons.
· Silicon is a semiconductor: it has limited conductivity because it does not have freely mobile delocalised electrons like a metal.
· P, S, Cl and Ar do not conduct electricity because they have no mobile ions or electrons.
· Exam phrase: “Electrical conductivity requires mobile charged particles.”
· Do not say simple molecular substances cannot conduct because they are “covalent”; the better reason is no mobile charge carriers.

Structure and bonding summary

· Na, Mg, Al: giant metallic structure; held by metallic bonding between positive ions and delocalised electrons.
· Si: giant covalent structure; strong covalent bonds throughout a 3D network.
· P₄, S₈, Cl₂: simple molecular structures; weak London forces between molecules.
· Ar: monatomic; weak London forces between atoms.
· For melting point questions, focus on the particles being separated during melting:
· metals: overcome metallic bonding
· silicon: break many covalent bonds
· simple molecules/atoms: overcome intermolecular forces only.
· For conductivity questions, focus on whether there are mobile electrons or mobile ions.

Common exam comparisons

· Why does atomic radius decrease across Period 3?
· Nuclear charge increases, electrons are added to the same shell, shielding is similar, so outer electrons are attracted more strongly.

· Why does melting point increase from Na to Al?
· Metallic bonding strengthens because metal ion charge increases, ionic radius decreases and more delocalised electrons are present.

· Why is Si’s melting point much higher than Al’s?
· Si is giant covalent, so many strong covalent bonds must be broken; Al has metallic bonding.

· Why does melting point drop sharply from Si to P?
· Si is giant covalent, but P₄ is simple molecular, so only weak London forces are overcome when phosphorus melts.

· Why does sulfur have a higher melting point than phosphorus/chlorine?
· S₈ molecules are larger and have more electrons, so they have stronger London forces than P₄ or Cl₂.