Electronegativity is a measure of the ability of an atom in a chemical bond to attract a shared pair of electrons towards itself. Unlike ionization energy or electron affinity, electronegativity is not a directly measurable energy value but rather a relative scale, most commonly the Pauling scale, where Fluorine (the most electronegative element) is assigned a value of 4.0.

Periodic Trends in Electronegativity:
- Across a period (left to right): Electronegativity generally increases. As we move across a period, the effective nuclear charge experienced by the valence electrons increases, and the atomic radius decreases. Both these factors contribute to a stronger pull on shared electrons in a bond.

Down a group (top to bottom): Electronegativity generally decreases. As we move down a group, the principal energy level of the valence electrons increases, leading to larger atomic radii. The increased distance from the nucleus and enhanced shielding by inner electrons reduce the attractive force the nucleus exerts on bonding electrons.

The difference in electronegativity between two bonded atoms helps predict the type of bond:
- Large difference (>1.7-1.8): Predominantly ionic bond.
- Intermediate difference (0.4-1.7): Polar covalent bond.
- Small difference (<0.4): Nonpolar covalent bond.
- Zero difference: Pure covalent bond.

Atomic Radius is a measure of the size of an atom. It is often defined as one-half the distance between the nuclei of two identical atoms that are bonded together (e.g., in a diatomic molecule like Cl₂).

Periodic Trends in Atomic Radius:
- Across a period (left to right): Atomic radius generally decreases. As we move from left to right across a period, the number of protons in the nucleus increases, leading to a greater nuclear charge. Electrons are added to the same principal energy level, meaning they do not significantly increase the shielding effect. The stronger attraction from the nucleus pulls the electron shells closer, resulting in a smaller atomic radius.

Down a group (top to bottom): Atomic radius generally increases. As we move down a group, electrons are added to new, successively higher principal energy levels. These new shells are further from the nucleus. Furthermore, the inner electrons effectively shield the outer electrons from the full nuclear charge, allowing the electron cloud to expand and leading to a larger atomic radius.

Ionic Radius refers to the radius of an ion. It differs from the atomic radius because the gain or loss of electrons significantly alters the balance between nuclear attraction and electron-electron repulsion.
- Cations (positive ions): Cations are always smaller than their parent atoms. This is due to two main reasons:
1. The complete removal of the outermost electron shell (if all valence electrons are lost, e.g., Na to Na⁺).
2. Even if a shell is not completely removed, the loss of electrons reduces the total electron-electron repulsion among the remaining electrons. With fewer electrons, the remaining electrons are pulled more tightly by the same positive nuclear charge. For example, Li⁺ is much smaller than Li.
- Anions (negative ions): Anions are always larger than their parent atoms. This is because the addition of electrons increases the electron-electron repulsion within the electron cloud. This increased repulsion causes the electron cloud to spread out, leading to a larger effective radius. The nuclear charge remains the same but now has to attract more electrons, resulting in a weaker average pull per electron. For example, Cl⁻ is larger than Cl.

Isoelectronic Series: An isoelectronic series consists of atoms and ions that have the same number of electrons (and therefore the same electron configuration). In an isoelectronic series, the ionic radius decreases as the nuclear charge (number of protons) increases. For example, consider the isoelectronic series N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺. All these species have 10 electrons (the electron configuration of Neon). Their radii decrease in this order because the increasing number of protons from N (Z=7) to Al (Z=13) exerts a stronger pull on the same number of electrons, thereby contracting the electron cloud.