Question

Why is the ionization energy of \(\mathrm{Cl}\) lower than \(\mathrm{F}\), but higher than \(\mathrm{S}\) ?

Short answer

The ionization energy of \(\text{Cl}\) is lower than \(\text{F}\) due to greater electron shielding but higher than \(\text{S}\) due to a higher nuclear charge.

Step-by-step solution

- Define Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

- Explain Periodic Trends in Ionization Energy

In general, ionization energy increases across a period from left to right and decreases down a group in the periodic table. This is due to the increasing nuclear charge and the decreasing atomic radius across a period, and the increasing electron shielding down a group.

- Compare \(\text{Cl}\) and \(\text{F}\)

Chlorine (\(\text{Cl}\)) is in the same group as Fluorine (\(\text{F}\)), group 17, and is below it in the periodic table. Therefore, chlorine has more electron shells than fluorine, which increases electron shielding. This makes it easier to remove an outer electron from chlorine, resulting in a lower ionization energy compared to fluorine.

- Compare \(\text{Cl}\) and \(\text{S}\)

Chlorine (\(\text{Cl}\)) is to the right of Sulfur (\(\text{S}\)) in period 3. Chlorine has a higher nuclear charge than sulfur, which results in a greater attraction between the nucleus and the outer electrons. Therefore, more energy is required to remove an outer electron from chlorine than from sulfur, giving chlorine a higher ionization energy compared to sulfur.

Key concepts

Periodic Trends

Periodic trends are patterns in the periodic table that show how certain properties of elements change across a period (from left to right) and down a group (from top to bottom). One key periodic trend is ionization energy. As you move from left to right across a period, ionization energy generally increases. This is because atoms have more protons as you go across a period, making their nuclear charge stronger. A stronger nuclear charge means electrons are held more tightly, so more energy is needed to remove one. However, as you move down a group, ionization energy generally decreases. This is due to electrons being further from the nucleus and experiencing more electron shielding.

Electron Shielding

Electron shielding occurs when inner electrons block the attraction between the nucleus and the outer electrons. The more shells of electrons an atom has, the more shielding occurs, making it easier to remove an outer electron. In the case of chlorine (Cl) and fluorine (F), chlorine has more electron shells since it is lower down in the periodic table. This increased shielding in chlorine reduces the effective nuclear charge felt by the outermost electrons, making it easier to remove them compared to fluorine. Therefore, chlorine has a lower ionization energy than fluorine.

Nuclear Charge

Nuclear charge refers to the total charge of the protons in the nucleus. A higher nuclear charge means a stronger attraction between the nucleus and the electrons. Comparing elements in the same period, such as chlorine (Cl) and sulfur (S), chlorine has a higher nuclear charge because it has more protons. This increased nuclear charge attracts the outer electrons more strongly, making it more difficult to remove them. Hence, chlorine has a higher ionization energy than sulfur. In summary:

• Ionization energy increases across a period due to increasing nuclear charge.
• Ionization energy decreases down a group due to electron shielding.
• Chlorine has lower ionization energy than fluorine but higher than sulfur due to these trends.