Final answer:
The shielding effect is caused by inner electrons that decrease the effective nuclear charge felt by outer electrons in an atom. As we move across a period in the periodic table, Zeff increases because the increase in shielding is not enough to counteract the added nuclear charge, resulting in smaller atomic radii.
Step-by-step explanation:
The shielding effect in an element is caused by the inner electrons, which create a 'shield' that reduces the effective nuclear charge felt by the outer electrons. The effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom. The concept is summarized by the equation: Zeff = Z − shielding, where 'Z' is the atomic number and 'shielding' represents the reduction of nuclear attraction due to the inner electrons. As a result, outer electrons feel a smaller attractive force from the nucleus than they would if there were no inner electrons.
As we move from left to right across a period in the periodic table, the nuclear charge (Z) of the atoms increases by one each time a new element is encountered. However, the increase in shielding is relatively small in comparison. This slight increase in shielding does not completely counterbalance the increase in nuclear charge. Therefore, the effective nuclear charge increases across a period, causing electrons to be pulled closer to the nucleus, which in turn decreases the covalent radii of these atoms.