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.

Understanding the size of atoms and ions is key to grasping many concepts in chemistry.
The atomic radius is the distance from the nucleus to the outer edge of the electron cloud. As we move across a period in the periodic table, the atomic radius decreases because each subsequent element has an additional proton in the nucleus, pulling the electron cloud closer.
Conversely, when an atom becomes an ion, the size changes significantly. Anions, or negative ions, are larger than their corresponding neutral atoms. This increase in size occurs because the addition of negatively charged electrons to an atom leads to increased electron-electron repulsion in the electron cloud, causing it to expand.
On the other hand, cations, or positive ions, are smaller than their atoms because the loss of electrons can lead to a decrease in electron shielding and the remaining electrons are drawn closer to the nucleus by its unaltered positive charge.

Transition metals are known for their ability to form ions with various oxidation states. They often achieve stability by losing electrons and obtaining an electronic arrangement similar to a noble gas.
For example, scandium, a transition metal, tends to lose electrons from its 3d and 4s subshells. This results in a Sc3+ ion with the same electron configuration as the noble gas argon. By losing electrons, scandium attains a stable, lower energy state, fulfilling the noble-gas electron configuration. This tendency is common among transition metals, which often form positive ions by shedding their outermost electrons.

Electronegativity is a measure of an atom's ability to attract and bond with electrons. It's essential when predicting how atoms will bond and the type of bond they will form.
Atoms with high electronegativity, such as fluorine, attract electrons strongly. However, as you move down a group in the periodic table, the electronegativity decreases. This is because the atoms at the bottom of a group have more electron shells, which increase the distance between the nucleus and valence electrons. This distance, coupled with the electron shielding effect, weakens the nucleus's pull on the valence electrons, thereby reducing the atom's electronegativity. This trend explains why halogens at the top of Group 17 are more electronegative than those at the bottom.

Anions play a critical role in chemical reactions and bonding. These negatively charged ions form when an atom gains extra electrons, leading to a size larger than the original atom.
Anions like chloride (Cl-) or oxide (O2-) have more electrons than protons, which increases electron-electron repulsion within the electron cloud, making the ion's radius larger than that of its atom. This expansion impacts properties such as ionic radius and polarity. Furthermore, anions are attracted to positively charged ions (cations), resulting in ionic bonds, which are crucial in the formation of ionic compounds.

Noble gases are the group 18 elements known for their lack of reactivity, and this inertness is due to their electron configuration. Every noble gas has a complete set of electrons in its outer shell, making it stable and unreactive.
Other elements strive to achieve this stable configuration, referred to as the 'octet rule' for many elements. When atoms form ions, they often do so by losing or gaining electrons to attain a noble-gas electron configuration. For instance, metals lose electrons to mirror the electron configuration of the noble gas preceding them in the periodic table, while nonmetals gain electrons to resemble the configuration of the noble gas following them. This pursuit of electron configuration stability is a driving force in chemical bonding.

The electron shielding effect is a phenomenon that occurs within an atom as it has multiple layers of electrons.
The inner layers of electrons provide a 'shield' that lessens the pull of the nucleus on the outer electrons. This effect becomes more pronounced in larger atoms with more electron shells. Because the inner electrons repel the outer electrons, they experience less attraction to the positive charge of the nucleus, effectively increasing the atomic size and decreasing the electronegativity. This shielding influences various atomic properties, including ionization energy, atomic radius, and the atom's ability to attract electrons in bonding situations.