Question

In going across a row of the periodic table, electrons are added and ionization energy generally increases. In going down a column of the periodic table, electrons are also being added but ionization energy decreases. Explain.

Short answer

In going across a row (period) of the periodic table, the ionization energy generally increases due to the increase in nuclear charge and relatively constant electron shielding, which results in a higher effective nuclear charge (Z_eff) and stronger attraction between electrons and the nucleus. Conversely, in going down a column (group), the ionization energy decreases because the atomic radius increases, weakening the attraction between electrons and the nucleus, and increased electron shielding reduces the effective nuclear charge experienced by outer electrons, making them easier to remove.

Step-by-step solution

Definition of Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is an important property of atoms that can help us understand their reactivity and bonding behavior.

Electron Configuration and Atomic Structure

As we move across a row (period) of the periodic table, electrons are added to the same energy level (shell) and the atomic number increases. Since electrons are added to the same shell, they experience a similar distance from the nucleus. In contrast, as we move down a column (group) of the periodic table, electrons are added to a new, higher energy level that is further from the nucleus.

Nuclear Charge and Atomic Radius

Nuclear charge is the charge of the nucleus, which is determined by the number of protons in the nucleus. As we move across a period, the nuclear charge increases, which means the nucleus attracts the electrons more strongly. On the other hand, as we move down a group, the atomic radius increases, meaning the electrons are further away from the nucleus and are less attracted to it.

Effective Nuclear Charge and Electron Shielding

Effective nuclear charge (Z_eff) is the actual positive charge experienced by an electron, considering the shielding effect of other electrons between the nucleus and the electron in question. As we move across a period, the shielding effect remains relatively constant, whereas the nuclear charge increases. This results in a higher effective nuclear charge that attracts electrons more effectively. Conversely, as we move down a group, the increased electron shielding reduces the effective nuclear charge experienced by the outer electrons, allowing them to be removed more easily.

Ionization Energy across a Period

In going across a row of the periodic table, electrons are being added to the same energy level while the nuclear charge increases, making the electrons experience a higher effective nuclear charge (Z_eff). This causes the electrons to be more strongly attracted to the nucleus, and thus, more energy is required to remove them, resulting in an overall increase in ionization energy.

Ionization Energy down a Group

In going down a column of the periodic table, electrons are being added to a new, higher energy level, further from the nucleus. This results in a larger atomic radius and a weaker attraction between the electrons and the nucleus. Additionally, increased electron shielding reduces the effective nuclear charge experienced by the outer electrons, making them easier to remove. Consequently, ionization energy decreases as we move down a group in the periodic table.