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

Verified step by step guidance

1. Understand that atomic radius is the distance from the nucleus of an atom to the outermost electron shell.
2. Recognize that as you move to the right across a period in the periodic table, the number of protons in the nucleus increases, which increases the positive charge of the nucleus.
3. Acknowledge that the increased nuclear charge attracts the electrons more strongly, pulling them closer to the nucleus and resulting in a smaller atomic radius for main-group elements.
4. Note that for transition elements, the additional electrons are added to an inner d-subshell rather than the outermost shell, which does not significantly increase the shielding effect.
5. Understand that the effective nuclear charge experienced by the outer electrons in transition elements does not increase as much as in main-group elements, so the atomic radius does not decrease significantly across a period for transition elements.

Key Concepts

Atomic Radius

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 (Z_eff)

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 and Electron Shielding

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.