Problem 63 - A Complete Resource Book in Chemistry for JEE Main

Question

In which of the following molecules the central atom does not follow the octet rule? (a) \(\mathrm{CO}_{2}\) (b) \(\mathrm{H}_{2} \mathrm{~S}\) (c) \(\mathrm{BF}_{3}\) (d) \(\mathrm{PCl}_{3}\)

Step-by-Step Solution

Answer

The molecule BF3 does not follow the octet rule.

Step 1: Evaluate CO2

In carbon dioxide, CO2, carbon is the central atom. Carbon forms two double bonds with oxygen. Each double bond consists of two pairs of electrons, for a total of 8 electrons around carbon: 4 from bonds with the first oxygen and 4 from bonds with the second oxygen. Thus, carbon obeys the octet rule in CO2.

Step 2: Evaluate H2S

In hydrogen sulfide, H2S, sulfur is the central atom. Sulfur forms two single bonds with hydrogen atoms. Additionally, sulfur has 2 lone pairs, for a total of 8 electrons surrounding it (4 from lone pairs and 4 from bonds with hydrogens). Thus, sulfur follows the octet rule.

Step 3: Evaluate BF3

In boron trifluoride, BF3, boron is the central atom. Boron forms three single bonds with fluorine atoms. This gives boron a total of 6 electrons (2 from each bond), which is less than 8. Boron often forms compounds that do not complete the octet rule, thus BF3 does not follow the octet rule.

Step 4: Evaluate PCl3

In phosphorus trichloride, PCl3, phosphorus is the central atom. Phosphorus forms three single bonds with chlorine atoms and has one lone pair. This gives phosphorus a total of 8 electrons (6 from the bonds with chlorine and 2 from a lone pair), allowing it to follow the octet rule.

Key Concepts

Electron Configuration

Electron configuration describes how electrons are distributed within an atom's orbitals, determining its chemical properties. Each atom follows a specific pattern based on its atomic number, which corresponds to the number of electrons and protons the atom contains. The arrangement of electrons is crucial because:

• It predicts how an atom will bond with other atoms.
• It provides insight into the chemical reactivity and properties of the element.
• It helps in understanding the concept of the octet rule, which is essential for forming stable molecules.

For instance, consider carbon, which has an electron configuration of 1s² 2s² 2p². This configuration shows that carbon has four valence electrons in its outermost shell, allowing it to form up to four bonds with other atoms to achieve a stable octet configuration. Conversely, elements like boron have fewer valence electrons and may not achieve a full octet, as seen in boron trifluoride (BF₃). Understanding electron configuration can thus be seen as foundational for grasping why certain atoms do or do not follow the octet rule.

Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds, and it plays a fundamental role in determining the structure and properties of substances. There are several types of chemical bonds:

• Ionic Bonds: Formed through the transfer of electrons from one atom to another, typically between metals and non-metals.
• Covalent Bonds: Involves the sharing of electrons between atoms. For example, atoms in \\( \mathrm{CO}_{2} \) are bonded covalently.
• Metallic Bonds: Present in metal atoms where electrons form a cloud around a lattice of positive ions.

In molecular compounds, like \\( \mathrm{PCl}_{3} \), phosphorus and chlorine atoms are covalently bonded, ensuring the octet rule is fulfilled for each atom involved. However, some molecules like \\( \mathrm{BF}_{3} \) don't comply with the octet rule due to unique characteristics of boron that allows it to form stable compounds with underfilled electron shells. Understanding chemical bonding is crucial to predict how molecules form, behave, and interact.

Lewis Structures

Lewis structures are diagrams that help visualize the electron configuration and bonding between atoms in a molecule. They are essential tools in chemistry for depicting electrons in the valence shell. Here's what they show:

• The layout of bonds between atoms, highlighting single, double, or triple covalent bonds.
• Any lone pairs of electrons that are not involved in bonding.

Using Lewis structures, we can understand how molecules like \\( \mathrm{CO}_{2} \), \\( \mathrm{H}_{2}\mathrm{S} \), and \\( \mathrm{BF}_{3} \) map out their electron distribution. For instance, in \\( \mathrm{CO}_{2} \), the Lewis structure shows carbon forming double bonds with each oxygen atom, completing carbon's octet. In contrast, \\( \mathrm{BF}_{3} \) displays three single bonds to fluorine, illustrating how boron doesn't fulfill the octet rule. Thus, Lewis structures allow us to examine exceptions and patterns in chemical bonding, making them a valuable visual tool in understanding complex chemical concepts.