Problem 155 - A Complete Resource Book in Chemistry for JEE Main

Question

The increasing order of the first ionization enthalpies of the element \(\mathrm{B}, \mathrm{P}, \mathrm{S}\) and \(\mathrm{F}\) (lowest first) is (a) \(\mathrm{F}<\mathrm{S}<\mathrm{P}<\mathrm{B}\) (b) \(\mathrm{P}<\mathrm{S}<\mathrm{B}<\mathrm{F}\) (c) \(\mathrm{B}<\mathrm{P}<\mathrm{S}<\mathrm{F}\) (d) \(\mathrm{B}<\mathrm{S}<\mathrm{P}<\mathrm{F}\)

Step-by-Step Solution

Answer

The correct order is (c) \( \mathrm{B}<\mathrm{P}<\mathrm{S}<\mathrm{F} \).

Step 1: Understand Ionization Enthalpy

Ionization enthalpy refers to the energy required to remove an electron from an atom in its gaseous state. Higher ionization enthalpy indicates that it is more difficult to remove an electron.

Step 2: Review Periodic Trends

In the periodic table, ionization enthalpy generally increases across a period from left to right and decreases down a group. This is because atoms become smaller and have a stronger attraction for electrons across a period while they get larger and electron shielding increases down a group.

Step 3: Analyze the Elements

The elements given are B (Boron), P (Phosphorus), S (Sulfur), and F (Fluorine). Boron, being in group 13, will have a lower ionization enthalpy than the others. Phosphorus and Sulfur are closer in ionization energy, but generally, Phosphorus (group 15) has a lower ionization enthalpy than Sulfur (group 16). Fluorine, in group 17, typically has the highest ionization enthalpy among these.

Step 4: Arrange Elements by Ionization Enthalpy

Based on the periodic trends and group positions: B < P < S < F. Boron, being the leftmost and low group element among these, has the lowest ionization enthalpy, while Fluorine, being the rightmost and in a higher group, has the highest.

Key Concepts

Periodic Table Trends

Understanding periodic table trends is key to mastering chemistry concepts like ionization enthalpy. As you move across the periodic table from left to right, the ionization enthalpy generally increases. This is due to the increase in nuclear charge, which means that electrons are held more tightly by the positively charged nucleus, making them harder to remove. Thus, it takes more energy to ionize elements that sit more to the right of a period.
From top to bottom in a group, the ionization enthalpy tends to decrease. As atoms get larger, the outer electrons are farther from the nucleus and are less tightly held due to increased electron shielding. This means that it is easier to remove these outer electrons, resulting in lower ionization enthalpy.
Overall, understanding these trends helps predict the behavior of elements, such as in determining the order of ionization enthalpy for B, P, S, and F.

Periodic Groups

In the context of the periodic table, periodic groups are the columns, labeling elements with similar outer electron configurations and, often, similar chemical properties. Grouping elements make it easier to predict their behavior in reactions and help explain trends observed across different elements.
For example, in Group 13, Boron (B) is known to have one primary outer electron, making it easier to remove electrons compared to elements in groups with fuller outer shells. Therefore, it has a comparatively lower ionization enthalpy.
On the other hand, elements like Fluorine in Group 17 have nearly filled outer shells, making their electrons harder to remove and resulting in higher ionization enthalpies. So, understanding which elements belong to which group can aid in arranging them in order of ionization enthalpy, like B < P < S < F as noted in the exercise.

Electron Shielding

Electron shielding is a core concept when discussing why ionization enthalpy decreases down a group on the periodic table. As atoms have more electron shells, the inner electrons effectively block the pull of the nucleus on the outer electrons. This "shielding effect" reduces the force exerted by the nucleus on outer electrons, allowing them to be removed with less energy.
This effect is why larger atoms found further down a group have lower ionization enthalpies. For instance, as you compare elements down a group, the increase in inner electron shells means the outermost electron is further away from the nucleus, leading to a reduced effective nuclear charge.
Understanding electron shielding clarifies why an element like Boron, even though it is part of the same period as Fluorine, has a lower ionization enthalpy due to its open electron shells and greater distance over which the nuclear charge acts effectively.