Explain the trend in melting points of the period 3 elements

There are four main areas we need to consider in answering this question. The first is Sodium, Magnesium and Aluminium. These are all metals and so have a metallic bonding structure. As you go across the period, the atomic radii decreases. This is due to an increased positive charge of the nucleus (increased electronegativity) pulling the outer electrons closer to the nucleus. Shielding does not charge significantly as the electron is added to the outer shell, so atomic radii falls. Metallic bonding strength relies on number of delocalised electrons (increase increases attraction), charge of ions (increase increases attraction) and radius size (decrease increases attraction). Going from Na to Al delocalised electrons increases (1 to 3 per ion), charge increases (1+ to 3+) and radius falls so melting point increases as it takes more energy to overcome the electrostatic forces of attraction between positive ions and delocalised electrons. Silicon has strong covalent bonds linking the atoms together in a tetrahedral macromolecular structure. This is a giant covalent structure. Lots of energy is required to overcome the strong covalent bonds, so silicon has the highest melting point. Phosphorous (P₄ Mr = 124), Sulfur (S₈ Mr = 256) and chlorine (Cl₂ Mr = 71) all are molecular substances and so are bonded to each other by van der Waal forces. Larger molecules have larger electron clouds leading to stronger van der Waals forces so sulphur has the largest melting point with phosphorus having a lower melting point and chlorine having the lowest due to differing energy required to overcome the van der Waals forces. Finally, Argon exists as a single atom (monoatomic Ar = 40). Argon relies on van der Waals forces however due to Argon’s low mass the van der Waals forces are very weak. This means very little energy is required to overcome the forces, so the melting point is the lowest across the period.