Question

Explain why (a) negative ions are larger than their corresponding atoms. (b) scandium, a transition metal, forms an ion with a noble-gas structure. (c) electronegativity decreases down a group in the periodic table.

Short answer

Answer: Negative ions have a larger overall size compared to their corresponding neutral atoms because the addition of extra electron(s) increases electron-electron repulsion, causing the energy level to expand, resulting in a larger ion.

Step-by-step solution

Understanding negative ions and atoms

Negative ions are formed when an atom gains one or more electrons. This additional electron effectively increases the size of the ion compared to the neutral atom. The additional electron(s) will be added to an existing energy level, causing an increase in electron-electron repulsion. This will, in turn, cause the energy level to expand, resulting in a larger overall size of the ion when compared to its corresponding atom.

Scandium and noble-gas structure

Scandium, a transition metal, forms an ion with a noble-gas structure. This is because transition metals have incompletely filled d orbitals, and in the case of scandium, the electron configuration is [Ar] 3d^1 4s^2. When scandium forms a positive ion, it loses those three outermost electrons, becoming Sc^3+ with the electron configuration [Ar], which is the noble-gas configuration of argon. The noble gas configuration represents a stable and full set of valence electrons, making it energetically favorable for transition metals like scandium to form ions with a noble-gas structure.

Electronegativity and its trend down a group

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Electronegativity generally decreases as we move down a group in the periodic table. This is due to the increase in atomic size as we go down the group. With each lower row in a group, an additional electron shell is added, which increases the distance between the nucleus (which has a positive charge due to protons) and the valence electrons that participate in bonding (which are negatively charged). As a result, the nucleus can't attract these valence electrons as effectively as if they were closer. Additionally, more electron shells result in greater shielding of the nucleus' positive charge by the inner shell electrons. This shielding effect further weakens the attraction between the nucleus and bonding electrons, ultimately causing a decrease in electronegativity down a group in the periodic table.