Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for transition elements.

Atomic radius refers to the size of an atom, typically measured from the nucleus to the outermost electron shell. As you move across a period in the periodic table, the atomic radius generally decreases due to the increasing positive charge of the nucleus, which pulls the electrons closer, resulting in a smaller atomic size.

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. As you move across a period, the number of protons increases while the shielding effect from inner electrons remains relatively constant, leading to a higher Z_eff. This increased attraction between the nucleus and the valence electrons causes the atomic radius to decrease.

Transition elements exhibit unique electron configurations that involve d-orbitals. As you move across the transition metals, the addition of electrons to the d-orbitals does not significantly increase the effective nuclear charge felt by the outermost s-electrons due to increased electron shielding. This results in a less pronounced decrease in atomic radius compared to main-group elements.