Periodicity of physical properties of the elements in Period 3

Periodicity is the repeating pattern of properties of elements across different periods. As you move from left to right across Period 3 (Na to Ar), the physical properties change in a predictable way. Atomic radius is the first key trend. It consistently decreases across the period. This is because, from sodium to argon, a proton is added to the nucleus and an electron is added to the same outer shell (the third shell). The increasing nuclear charge pulls the electrons in this outer shell more strongly towards the nucleus, shrinking the atom. The shielding effect from the inner shells remains relatively constant, so it doesn't counteract the stronger pull. Ionic radius also shows clear trends. Metal atoms (Na, Mg, Al) lose their outer shell electrons to form positive ions (cations). These cations are much smaller than their parent atoms because they have lost an entire electron shell and the remaining electrons are pulled more tightly by the unchanged nuclear charge. Non-metal atoms (P, S, Cl) gain electrons to form negative ions (anions). These anions are larger than their parent atoms because the addition of electrons increases the repulsion between the electrons in the outer shell, causing it to expand.

First ionisation energy (FIE) is the energy needed to remove one electron from each atom in one mole of gaseous atoms. The general trend across Period 3 is an increase. This is because as you move from Na to Ar, the nuclear charge increases and the atomic radius decreases. This means the outermost electron is held more strongly by the nucleus and requires more energy to be removed. However, there are two important 'dips' in this trend. The first dip is from Magnesium (Mg) to Aluminium (Al). Mg's electron configuration is [Ne]3s², while Al's is [Ne]3s²3p¹. The electron removed from Al is in a 3p orbital, which is at a slightly higher energy level and further from the nucleus than the 3s orbital. The 3s electrons also provide some shielding. Therefore, less energy is needed to remove this 3p electron. The second dip is from Phosphorus (P) to Sulfur (S). P's electron configuration is [Ne]3s²3p³, with one electron in each 3p orbital. S's is [Ne]3s²3p⁴, which has one 3p orbital containing a pair of electrons. The repulsion between these two paired electrons in the same orbital in sulfur makes it easier to remove one of them compared to removing an electron from a half-filled orbital in phosphorus.

The structure and bonding of the elements across Period 3 change dramatically, which directly explains the trends in their melting points and electrical conductivity.

• Sodium (Na), Magnesium (Mg), and Aluminium (Al): These are metals with giant metallic lattice structures. They are held together by strong metallic bonds, which is the electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons. From Na to Mg to Al, the number of delocalised electrons per atom increases (1, 2, 3 respectively) and the ionic charge increases (+1, +2, +3). This leads to a stronger metallic bond, so the melting point increases. These delocalised electrons are mobile charge carriers, so all three are good electrical conductors, with conductivity increasing from Na to Al.

• Silicon (Si): This is a metalloid with a giant molecular (or giant covalent) structure, similar to diamond. Each Si atom is covalently bonded to four others in a tetrahedral arrangement. To melt silicon, these numerous strong covalent bonds must be broken, which requires a vast amount of energy. This gives silicon a very high melting point, the highest in the period. It has very few free electrons, making it a semiconductor.

• Phosphorus (P), Sulfur (S), and Chlorine (Cl): These are non-metals with simple molecular structures. They exist as small, discrete molecules (P₄, S₈, Cl₂). The covalent bonds within the molecules are strong, but the forces between the molecules are weak van der Waals' forces. Melting only requires overcoming these weak intermolecular forces, not breaking the covalent bonds. Therefore, they have low melting points. The melting point increases from Cl₂ to P₄ to S₈ because the size of the molecule and number of electrons increases, leading to stronger van der Waals' forces. As there are no mobile charge carriers, they are electrical insulators.

• Argon (Ar): This is a noble gas that exists as individual atoms (monatomic). The only forces between the atoms are extremely weak van der Waals' forces, so it has a very low melting point.

The properties of Period 3 oxides transition from basic to amphoteric to acidic across the period, which is explained by the change in bonding from ionic to covalent.

• Na₂O and MgO: These are basic oxides. They have giant ionic lattice structures due to the large electronegativity difference between the metal and oxygen. This results in high melting points. They react with water to form alkaline solutions (e.g., Na₂O + H₂O → 2NaOH) and react with acids to form a salt and water (e.g., MgO + 2HCl → MgCl₂ + H₂O).
• Al₂O₃: This is an amphoteric oxide, meaning it can act as both an acid and a base. It has a giant ionic lattice with significant covalent character. It is insoluble in water but reacts with both strong acids and strong bases. Reaction with acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O. Reaction with base: Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄].
• SiO₂: This is an acidic oxide with a giant covalent structure. It has a very high melting point. It is insoluble in water but will react with strong bases like hot, concentrated sodium hydroxide: SiO₂ + 2NaOH → Na₂SiO₃ + H₂O.
• P₄O₁₀ and SO₂/SO₃: These are acidic oxides. They have simple molecular structures, resulting in low melting and boiling points. They react vigorously with water to form strong acids. P₄O₁₀ + 6H₂O → 4H₃PO₄ (phosphoric acid). SO₃ + H₂O → H₂SO₄ (sulfuric acid).

The trend in bonding for Period 3 chlorides mirrors that of the oxides, changing from ionic to covalent. This directly impacts their melting points and, most importantly, their reaction with water.

• NaCl and MgCl₂: These are white crystalline solids with giant ionic lattice structures and high melting points. They do not react with water but simply dissolve to form hydrated ions. The resulting solutions are approximately neutral (NaCl gives pH 7, MgCl₂ gives pH ~6.5 due to slight polarisation of water by the Mg²⁺ ion).
• AlCl₃: This is a complex case. As a solid, it has an ionic lattice. However, it has significant covalent character and sublimes at 180°C, forming a dimer, Al₂Cl₆, which is a simple molecule. It reacts vigorously and exothermically with water in a reaction called hydrolysis. White, steamy fumes of hydrogen chloride gas are produced, and the remaining solution is strongly acidic (pH ~3). A simplified equation is AlCl₃(s) + 3H₂O(l) → Al(OH)₃(s) + 3HCl(g).
• SiCl₄ and PCl₅: These are covalent substances that are liquids or low melting point solids at room temperature. They have simple molecular structures. They are rapidly and violently hydrolysed by water, producing dense white fumes of HCl and an acidic solution. SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g). PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(g).
• The trend is clear: ionic chlorides dissolve, while covalent chlorides are hydrolysed by water.