Why do metallic properties decrease across a period?

Metallic properties decrease across a period due to an increase in nuclear charge and decrease in atomic radius.

In the periodic table, as you move from left to right across a period, the number of protons in the nucleus (nuclear charge) increases. This increase in nuclear charge attracts the electrons more strongly, pulling them closer to the nucleus and decreasing the atomic radius. As a result, the outermost electrons are closer to the nucleus and more tightly held, making it harder for them to be lost and form positive ions, a key characteristic of metals.

Additionally, across a period, the number of energy levels (shells) remains the same while the number of electrons increases. These additional electrons go into the same energy level, not a new one further from the nucleus. This means that the shielding effect, where inner shell electrons shield outer shell electrons from the attractive force of the nucleus, remains constant. However, the increasing nuclear charge outweighs the constant shielding effect, pulling the outermost electrons closer to the nucleus.

This combination of decreasing atomic radius and increasing nuclear charge results in a decrease in metallic properties. Metals tend to have large atomic radii and low ionisation energies, meaning they can easily lose electrons to form positive ions. However, as you move across a period, atoms become smaller and their ionisation energies increase, making it harder for them to lose electrons and exhibit metallic properties. Therefore, elements transition from metals on the left of the period to non-metals on the right.

In summary, the decrease in metallic properties across a period is due to the increase in nuclear charge and decrease in atomic radius, which make it harder for atoms to lose electrons and form positive ions.