Problem 42 - A Complete Resource Book in Chemistry for JEE Main

Question

The strongest base is (a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{PH}_{3}\) (c) \(\mathrm{AsH}_{3}\) (d) \(\mathrm{SbH}_{3}\)

Step-by-Step Solution

Answer

The strongest base is \(\mathrm{NH}_3\).

Step 1: Understand the problem

We need to identify which of the given compounds is the strongest base: \(\mathrm{NH}_3\), \(\mathrm{PH}_3\), \(\mathrm{AsH}_3\), and \(\mathrm{SbH}_3\). The basicity generally depends on the ability of a molecule to donate a lone pair of electrons.

Step 2: Analyze the periodic trends

Elements \(\mathrm{N}\), \(\mathrm{P}\), \(\mathrm{As}\), and \(\mathrm{Sb}\) all belong to the same group (group 15) in the periodic table. As we move down the group, the atomic size increases and the electronegativity decreases, affecting the ability to donate lone pairs.

Step 3: Compare each compound's basicity

\(\mathrm{NH}_3\) (ammonia) has more electron density around the nitrogen atom due to its small size and high electronegativity, making it a better lone pair donor. As a result, \(\mathrm{NH}_3\) is a stronger base compared to the other hydrides in the group.

Key Concepts

Periodic Trends in Basicity

In chemistry, understanding periodic trends helps us predict the properties of elements and their compounds. When it comes to basicity, these trends are quite revealing. Basicity in group 15 elements is influenced by two main factors: atomic size and electronegativity.

• Atomic Size: As we descend the group, moving from nitrogen to antimony, the atomic radius increases. This increase makes it more difficult for the atom to tightly hold onto the lone pair electrons, and thus lessens its ability to donate those electrons to act as a base.
• Electronegativity: Elements higher up in the group, like nitrogen, are more electronegative. This means they have a stronger tendency to attract electrons towards themselves, enhancing their ability to hold and donate lone pair electrons. As we move down the group, electronegativity decreases, making these elements less basic.

By combining these factors, it becomes clear why ammonia (\(\mathrm{NH}_3\)) shows stronger basicity compared to its heavier counterparts. Ammonia's small size coupled with higher electronegativity allows it to donate the lone pair of electrons more readily.

Lone Pair Electron Donation

The ability to donate lone pair electrons is what defines a compound's basic character. Within group 15, each element’s hydride, such as \(\mathrm{NH}_3\), has a lone pair of electrons. However, the ease with which these electrons are donated varies significantly.

Influence of Atomic Size and Shielding

Atomic size plays a crucial role in lone pair donation. Smaller atoms like nitrogen have electron clouds that are closer and more concentrated. This makes their lone pairs more available for donation compared to the loosely held electrons in larger atoms like antimony.

Electronegativity and Its Effects

High electronegativity causes atoms to draw electrons towards themselves, stabilizing the electron pair. Nitrogen's higher electronegativity relative to phosphorus, arsenic, and antimony means its electrons are held more closely, ready to be donated, making ammonia a stronger base.

Group 15 Hydrides Comparison

The hydrides of group 15—ammonia \(\mathrm{NH}_3\), phosphine \(\mathrm{PH}_3\), arsine \(\mathrm{AsH}_3\), and stibine \(\mathrm{SbH}_3\)—show a clear trend in basicity. Ammonia tops the list with phosphine trailing behind, followed by arsine and stibine. Understanding why ammonia is the most basic can be simplified by comparing these compounds' structures and properties.

• Ammonia (\(\mathrm{NH}_3\): Contains a tetrahedral shape with a lone pair of electrons on nitrogen. Its smaller size and higher electronegativity mean the lone pairs are readily available for donation.
• Phosphine (\(\mathrm{PH}_3\): Larger phosphorus atom results in weaker hold on the electron pair, reducing basicity compared to ammonia.
• Arsine (\(\mathrm{AsH}_3\) and Stibine (\(\mathrm{SbH}_3\): With increasing atomic size and decreasing electronegativity, these hydrides become progressively weaker as bases.

Thus, the combination of smaller atomic radius and higher electronegativity makes \(\mathrm{NH}_3\) the strongest base within these group 15 hydrides.