The chemical reactivity of metals and nonmetals follows distinct trends in the periodic table, influenced mainly by their atomic structure and electron configurations.
For metals, as you move from left to right across a period, the atomic number increases. This increase in protons leads to a greater positive charge in the nucleus, which pulls the electrons closer and increases the effective nuclear charge. As a result, the outermost electrons become more tightly held, making it harder for the metal to lose electrons and react. Hence, metal reactivity decreases from left to right.
Conversely, moving down a group increases the number of electron shells, which creates a greater distance between the nucleus and the outermost electrons. This results in weaker attraction between the nucleus and the valence electrons due to increased electron shielding. Therefore, it becomes easier for these outer electrons to be lost, leading to an increase in reactivity as you go down the group.
Now, looking at nonmetals, their reactivity trends are different. As you move from left to right across a period, nonmetals gain electrons to fill their valence shells, which makes them more reactive. Typically, they need to gain fewer electrons to achieve a stable electron configuration, hence their reactivity increases from left to right.
On the other hand, as you move down a group of nonmetals, the additional electron shells mean that the valence electrons are further from the nucleus and experience more shielding. This makes it harder for nonmetals to attract additional electrons, thus decreasing their reactivity as you move down the group.
In summary, the trends in chemical reactivity for metals and nonmetals across periods and groups can be attributed to the interplay of nuclear charge, electron shielding, and the energies associated with gaining or losing electrons.