Which of the following compound cannot act as a Lewis base?
(a) NF3
(b) PCl5
(c) SF4
(d) ClF3

Q: Why is PCl5 a Lewis acid if its octet is already expanded?

Even though phosphorus has 10 electrons in its valence shell in $PCl_5$, it still has empty 3d orbitals. It can accept another pair of electrons (e.g., from $Cl^-$) to form the $PCl_6^-$ complex ion, thus acting as a Lewis acid.

Q: Can a molecule be both a Lewis acid and a Lewis base?

Yes, some molecules can act as both depending on the reaction. For example, water ($H_2O$) has lone pairs on oxygen (Lewis base) but can also accept electrons in certain complex reactions, though it's primarily a Lewis base. However, $PCl_5$ strictly lacks lone pairs on the central atom.

Q: How do I quickly find the number of valence electrons?

The number of valence electrons corresponds to the group number (for groups 1-2 and 13-18, it's the last digit of the group number). For example, N and P are in Group 15, so they have 5 valence electrons.

Q: Does the electronegativity of the surrounding atoms affect Lewis basicity?

Yes! Highly electronegative atoms like fluorine pull electron density away from the central atom, making the lone pair less available for donation. For instance, $NF_3$ is a much weaker Lewis base than $NH_3$.

Q: Why doesn't NF3 act as a Lewis acid?

Nitrogen is a second-period element and does not have empty d-orbitals in its valence shell. Therefore, it cannot expand its octet to accept more electron pairs, meaning it cannot act as a Lewis acid.

Question

Which of the following compound cannot act as a Lewis base?

Options:

(A) $	ext{NF}_3$

(B) $	ext{PCl}_5$

(C) $	ext{SF}_4$

(D) $	ext{ClF}_3$

Sigma Percentile · Accepted Answer

### The Essence of Lewis Theory

Imagine you are at a grand chemical marketplace. In this bustling bazaar, molecules are constantly trading, sharing, and negotiating for electrons.

At the heart of this marketplace are two distinct types of traders: the **Lewis acids** and the **Lewis bases**.

A **Lewis base** is like a wealthy merchant who has a surplus of goods—specifically, a non-bonding pair of electrons, commonly known as a **lone pair**. Because they have this extra pair, they are more than happy to donate it to someone in need.

On the other hand, a **Lewis acid** is a trader with empty warehouse space. They have vacant orbitals and are desperately looking to accept an electron pair to achieve stability.

Our mission in this problem is to inspect four specific molecules and determine which one is absolutely incapable of being a Lewis base. To do this, we must play the role of a chemical detective and search for the presence of lone pairs on the central atom of each molecule.

### Decoding the Candidates

Let's bring our first suspect to the stand: Nitrogen trifluoride, or $NF_3$.

Nitrogen sits comfortably in Group 15 of the periodic table, which means it brings 5 valence electrons to the bonding table. In $NF_3$, it forms three single bonds with three fluorine atoms, using up 3 of its electrons.

What remains? Exactly 2 electrons, which pair up to form **1 lone pair**. Because it has this lone pair, $NF_3$ can definitely act as a Lewis base.

Next up is Sulphur tetrafluoride, $SF_4$.

Sulphur is a member of Group 16, boasting 6 valence electrons. It uses 4 of these electrons to bond with four fluorine atoms.

Simple math tells us that $6 - 4 = 2$ electrons left over. That's **1 lone pair** sitting on the sulphur atom. Thus, $SF_4$ is also fully capable of acting as a Lewis base.

Our third candidate is Chlorine trifluoride, $ClF_3$.

Chlorine is a halogen from Group 17, armed with 7 valence electrons. It shares 3 of these electrons to form bonds with three fluorine atoms.

Subtracting 3 from 7 leaves us with 4 non-bonding electrons. These arrange themselves into **2 lone pairs**. With not just one, but two lone pairs, $ClF_3$ is undeniably a Lewis base.

### The Case of Phosphorus Pentachloride

Finally, we arrive at Phosphorus pentachloride, $PCl_5$.

Phosphorus, like nitrogen, belongs to Group 15 and has 5 valence electrons. However, in $PCl_5$, phosphorus goes all out. It forms five single bonds with five chlorine atoms, utilizing every single one of its 5 valence electrons.

How many electrons are left? **Zero**.

There are absolutely no lone pairs on the central phosphorus atom. Without a lone pair, $PCl_5$ has nothing to donate. Therefore, it **cannot act as a Lewis base**.

But the story of $PCl_5$ doesn't end there. Phosphorus is a Period 3 element, which means it has a secret weapon: **empty 3d-orbitals**.

Because it has these vacant orbitals, $PCl_5$ can actually accept an incoming pair of electrons (for example, from a chloride ion to form the $PCl_6^-$ complex). This ability to accept electrons makes $PCl_5$ a classic **Lewis acid**.

### The Final Verdict

By systematically analyzing the valence electrons and bonding patterns of each molecule, we've uncovered the truth.

While $NF_3$, $SF_4$, and $ClF_3$ all possess at least one lone pair and can act as Lewis bases, $PCl_5$ has exhausted its valence electrons in bonding.

Lacking a lone pair and possessing empty d-orbitals, $PCl_5$ is strictly a Lewis acid.

**Therefore, the correct answer is $PCl_5$.**

Animated Solution for Chemistry - s and p-Block Elements: Which of the following compound cannot act as a Lewis base?

Select Answer:

Visualized Solution

$Lewis Base Concept$

$Checking for Lone Pairs$

$Analysis of NF_{3}$

$Analysis of SF_{4}$

$Analysis of ClF_{3}$

$Analysis of PCl_{5}$

$Conclusion$

$Further Thinking$

The Sigma Insight: Group 15 Elements

The Essence of Lewis Theory

Decoding the Candidates

The Case of Phosphorus Pentachloride

The Final Verdict

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