Problem 106 - A Complete Resource Book in Chemistry for JEE Main

Question

Noble gases do not react with other elements because (a) completely paired up electrons lead to stable electronic configuration (b) the size of their atoms is very small (c) they are found in abundance (d) they are monoatomic

Step-by-Step Solution

Answer

Noble gases do not react because they have a stable electronic configuration (option a).

Step 1: Understanding the Noble Gases

Noble gases are a group of chemical elements with very similar properties. These gases include helium, neon, argon, krypton, xenon, and radon. They are found in the far right column of the periodic table (Group 18).

Step 2: Identifying Electron Configuration

Noble gases have a full valence shell of electrons, meaning their outer electron shell is complete. Specifically, helium has two electrons filling its sole energy level, while other noble gases have eight electrons filling their outermost level, achieving a stable octet.

Step 3: Stability Through Full Valence Shells

A complete valence shell of electrons makes an atom energetically stable. Noble gases have all their electrons paired within their electron shells, reaching a stable configuration that does not require reactions with other elements.

Step 4: Choosing the Correct Answer

Given the understanding that noble gases do not react due to their stable electronic configuration, option (a) 'completely paired up electrons lead to stable electronic configuration' correctly explains why noble gases are chemically inert.

Key Concepts

Stable Electron Configuration

All atoms strive to reach a stable electron configuration. But what does this mean? Simply put, atoms seek to have a fully filled outermost electron shell, which gives them stability. For the noble gases, this is their natural state, which explains their lack of reactivity. Many elements on the periodic table will react with another to either gain or lose electrons, completing their outer shell like noble gases. Noble gases naturally have this configuration without needing to bond with other elements. - **Helium** has only 2 electrons in its outer shell but that's all it needs, since it's the only electron shell it has. - **Other noble gases** such as neon, argon, krypton, xenon, and radon, all have 8 electrons in their outer shell. This filling of their valence shell results in what's called a 'stable electronic configuration.' Stable means no tendency to change, which in chemistry translates to no need for reaction.

Chemical Inertness

Noble gases are famously known for their chemical inertness. This means they don’t typically react with other elements. But why is this the case? Their inertness is attributed to their stable electron configuration. With a fully filled valence shell, noble gases have no tendency to gain, lose, or share electrons. - Since their electron shells are full, there's no 'need' for them to interact with other elements to achieve stability. This inertness is why you rarely see noble gases forming compounds under normal conditions. In most chemical reactions, elements either share or exchange electrons to fill their outer shells. Noble gases, being content with their configuration, simply don't join the party.

Valence Shell Electrons

The concept of valence shell electrons is crucial to understanding chemical reactions. The valence shell is the outermost shell of an atom. The electrons in this shell are called 'valence electrons.' These electrons play a significant role in chemical reactivity. For most elements, having a complete set of valence electrons is key to achieving a stable state. - **Noble gases** already have a complete valence shell: - Helium has 2 electrons, while others like neon and argon have 8. With all their positions filled, there’s just no room for more electrons without creating instability. This complete set renders them disinterested in forming bonds with other elements. Understanding the importance of valence shell electrons helps explain the unique behaviors of different elements on the periodic table.