Question

The \(\mathrm{BCl}_{3}\) is a planar molecule whereas \(\mathrm{NCl}_{3}\) is pyramidal, becanse (a) B-Cl bond is more polar than \(\mathbf{N}-\) Cl bond (b) N-Cl bond is more covalent than \(\mathbb{B}\)-Cl bood (c) nitrogen atom is smaller than boron atoms (d) \(\mathrm{BCl}_{3}\) has no lone pair but \(\mathrm{NCl}_{3}\) has a lone pair of electrons.

Short answer

\textbf{NCl}\(_3\) has a pyramidal shape due to the presence of a lone pair of electrons on the nitrogen atom, while \textbf{BCl}\(_3\) is planar as boron has no lone pairs. The correct answer is (d).

Step-by-step solution

Understanding Molecular Geometry

The shape of a molecule is determined by the arrangement of electron pairs around the central atom, based on the VSEPR (Valence Shell Electron Pair Repulsion) theory. According to this theory, electron pairs will arrange themselves to minimize repulsion, and the presence of lone pairs can affect the molecular geometry.

Analyzing \textbf{BCl}\(_3\) Structure

Boron in \textbf{BCl}\(_3\) has three valence electrons and forms three covalent bonds with chlorine atoms, using all its valence electrons. There are no lone pairs of electrons on the boron atom, resulting in a trigonal planar geometry.

Analyzing \textbf{\textbf{NCl}\(_3\) Structure

Nitrogen in \textbf{NCl}\(_3\) has five valence electrons. It forms three covalent bonds with chlorine atoms and has one lone pair of electrons remaining. The presence of this lone pair means \textbf{NCl}\(_3\) has a pyramidal geometry, according to the VSEPR theory.

Identifying the Correct Reason for the Difference in Shape

The key to the different shapes lies in the presence or absence of lone pairs of electrons. \textbf{BCl}\(_3\) is planar because it has no lone pairs, whereas \textbf{NCl}\(_3\) is pyramidal because it has a lone pair. The other options do not directly explain the difference in molecular geometry.

Key concepts

VSEPR theory

Visualizing how molecules form their unique shapes isn't always straightforward, but the Valence Shell Electron Pair Repulsion (VSEPR) theory offers us a roadmap. Simply put, VSEPR theory proposes that the three-dimensional arrangement of atoms within a molecule is determined by the repulsions between electron pairs in the valence shell of the central atom. Each electron pair, be it in a bond or on its own as a lone pair, behaves like a negatively charged cloud, trying to stay as far away from each other as possible due to their like charges.

Imagine a room where people are holding balloons; naturally, they'd spread out to give each other space. Similarly, electron pairs spread out around the central atom, and the variety of possible spatial arrangements leads to the different molecular geometries observed. It's a bit like a dance of balance and symmetry, with the electron pairs choreographing the molecular structure.

Lone pair of electrons

A lone pair of electrons is like a reserved guest at a party – it takes up space but doesn't directly engage with other guests. In chemical terms, a lone pair refers to a pair of valence electrons that are not shared with another atom and therefore not involved in forming a covalent bond.

The presence of a lone pair on the central atom has a significant influence on the shape of a molecule. Due to the electronic repulsions, lone pairs tend to require more space than bonded pairs, thus pushing the bonded atoms closer together and leading to deviations from the idealized geometries. They are the behind-the-scenes players in the game of molecular shape, subtly but powerfully altering the arrangement of atoms.

Covalent bonds

Covalent bonds are the glue that holds molecules together in a sturdy embrace. These bonds form when atoms share pairs of electrons, and they are the result of each atom’s desire to achieve a more energetically favorable electron configuration, often akin to that of noble gases.

The dance of these shared electrons is delicate – each covalent bond represents a balance between attraction and repulsion. The electrons are attracted to the positively charged nuclei of the bonded atoms, while at the same time, the nuclei repel each other. When this bond forms, each atom in the pair brings its own electrons to the relationship, creating a shared electron cloud that orbits both nuclei – a truly collaborative effort in the microscopic world.

Electron pair repulsion

Picture yourself at a crowded party, trying to keep a decent amount of personal space. That's what electron pair repulsion is like at a molecular level. Electron pair repulsion is the idea that electron pairs – whether they're involved in bonds or are lone pairs – repel each other. This is because electrons are negatively charged, and like charges repel.

These repulsive forces are what dictate the VSEPR theory's rules and ultimately shape the molecule. For instance, in a molecule with only bonded pairs, the repulsions are distributed evenly, and the molecule adopts a shape that minimizes this repulsion, often resulting in symmetrical shapes. Introduce a lone pair, and the geometry changes as the repulsions become uneven. This dance of push and pull between electron pairs is fundamental to understanding the three-dimensional structure of molecules.