Final answer:
Ionization energy decreases down a group and increases across a period in the periodic table, influenced by atomic size and nuclear charge respectively. Option 2 is correct.
Step-by-step explanation:
The subject of this question is the periodic trend in ionization energy within the periodic table, which is a key concept in Chemistry, particularly when discussing the properties of elements and their behavior in forming chemical compounds. Ionization energy, by definition, refers to the amount of energy needed to remove an electron from an atom. As reflected in the provided materials and summarized in the following trends:
Ionization energy decreases from top to bottom within a group because as we move down a group, the valence electrons are further away from the nucleus due to the larger atomic radius, and thus, experience less electrostatic pull to the nucleus.
Ionization energy increases from left to right across a period because as more protons are added to the nucleus without a change in the energy level of the valence electrons, their attraction to the nucleus increases, leading to higher ionization energy.
This knowledge is essential for understanding both the chemical reactivity of elements and the formation of ions. It's important to note that while these trends generally hold true, there are exceptions due to electron shielding and the energy of different subshells, as explained in the provided materials.