Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Across a period, from left to right, the ionization energy generally increases. This increase is due to the higher nuclear charge of the atoms as we move across the period; the additional protons in the nucleus attract the electrons more strongly, making it harder to remove an electron. Moreover, the atomic radius decreases across a period, meaning the electrons are closer to the nucleus and more strongly attracted to it. Conversely, down a group, the ionization energy decreases. This is because the atomic radius increases with each successive element due to the addition of electron shells. The increased distance of the valence electrons from the nucleus, along with the increased shielding effect of the inner electrons, makes it easier to remove an outer electron.

The position of an element in the Periodic Table is closely linked to its electronegativity, which is the tendency of an atom to attract a shared pair of electrons towards itself in a chemical bond. Generally, electronegativity increases across a period (from left to right) and decreases down a group (from top to bottom). This trend is due to the increase in nuclear charge (number of protons) as we move across a period, drawing electrons closer to the nucleus, thereby increasing the atom's ability to attract electrons. Down a group, although the nuclear charge increases, the effect is offset by the additional shells of electrons that increase atomic radius and shield the valence electrons from the nucleus's pull. As a result, elements towards the top right of the Periodic Table (excluding the noble gases) are usually the most electronegative, with fluorine being the most electronegative element.

The elements in Group 18, known as noble gases, are unreactive due to their unique electronic configurations. Atoms of these elements have complete outer electron shells, which makes them highly stable. For instance, helium (He) has a complete outer shell with two electrons, while other noble gases like neon (Ne), argon (Ar), and krypton (Kr) have eight electrons in their outermost shell, satisfying the octet rule. This complete valence shell configuration means that noble gases have no tendency to gain, lose, or share electrons, which is the basis for most chemical reactions. As a result, they do not readily form compounds under normal conditions. Their chemical inertness is also reflected in their very high ionization energies and virtually zero electronegativities, further contributing to their lack of reactivity.

Elements in the same group have similar reactivities primarily due to their identical valence electron configurations. Despite having different numbers of electron shells, the elements in a group share the same number of electrons in their outermost shell. For example, all alkali metals in Group 1 have one electron in their outer shell, which is relatively loosely held and easily lost. This similarity in valence electron configuration leads to comparable chemical behaviors, particularly in how these atoms gain, lose, or share electrons during chemical reactions. The number of inner electron shells affects factors like atomic size and shielding effect but does not significantly alter the fundamental reactivity pattern of a group. This is because chemical reactivity is largely determined by the valence electrons, which are involved in chemical bonding and reactions.

The atomic radius of elements within a group in the Periodic Table increases as we move down the group. This increase is primarily due to the addition of shells (energy levels) as we go from one element to the next. Each shell represents a larger orbit for the electrons, thereby increasing the atomic size. For instance, in Group 1, lithium has only two shells, whereas sodium has three, making sodium's atomic radius larger than lithium's. Additionally, the increased number of electrons results in greater repulsion between electron shells, further contributing to an increase in atomic size. However, the effective nuclear charge (the net positive charge experienced by valence electrons) does not increase significantly within a group, as the additional electrons are added to the new outermost shell. This relatively constant effective nuclear charge across a group allows the atomic size to increase without much resistance from the nucleus.