The atomic radius decreases across a period primarily because as one moves from left to right across a period in the periodic table, the number of protons in the nucleus increases, resulting in a greater positive nuclear charge. Electrons are added to the same principal energy level (same shell) without increasing shielding effect significantly. This stronger effective nuclear charge pulls the electron cloud closer to the nucleus, reducing the size of the atom and thus decreasing the atomic radius across the period.
Key Points:
- Electrons are added to the same energy level across a period, so the shielding effect remains almost constant.
- Increasing number of protons increases the positive charge of the nucleus.
- The stronger attraction between the nucleus and the electrons pulls the electrons closer.
- This leads to a decrease in atomic size from left to right across the period.
- There are minor exceptions due to electron-electron repulsions and subshell filling but the general trend holds.
This trend explains why elements like lithium have larger atomic radii compared to fluorine in the same period, as fluorine’s nucleus pulls its electrons more strongly owing to a higher effective nuclear charge.
