The trend in melting points of Period 3 elements shows an increase from sodium to silicon, followed by a decrease.

This trend can be understood by examining the nature of the bonding in these elements. Sodium and magnesium exhibit metallic bonding, characterized by a sea of delocalized electrons that move freely around positively charged metal ions. The strength of metallic bonds correlates with the number of delocalized electrons; as we progress from sodium to magnesium, the number of these delocalized electrons increases. This accounts for magnesium’s higher melting point compared to sodium.

Aluminium features a giant metallic lattice structure, where each atom is bonded to eight neighboring atoms. This extensive bonding contributes to a very strong metallic bond, resulting in a high melting point.

Silicon, on the other hand, possesses a giant covalent lattice structure, with each silicon atom bonded to four neighboring atoms. The strength of these covalent bonds also leads to a high melting point.

In contrast, phosphorus, sulfur, chlorine, and argon have simple molecular structures that are held together by relatively weak intermolecular forces. As we move from silicon to argon, the number of electrons in the outer shell decreases, leading to weaker intermolecular forces and consequently lower melting points.

In summary, the melting point trend of Period 3 elements is largely influenced by their bonding characteristics. Metallic bonding results in high melting points, while covalent and intermolecular bonding lead to lower melting points.