Problem 66 - A Complete Resource Book in Chemistry for JEE Main

Question

The configuration of inert gas with smallest size and highest IE is (a) \(1 \mathrm{~s}^{2}\) (b) \(1 s^{2} 2 s^{2} 2 p^{6}\) (c) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{5}\) (d) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6} 3 \mathrm{~s}^{2} 3 \mathrm{p}^{6}\)

Step-by-Step Solution

Answer

The configuration with the smallest size and highest ionization energy is (a) \(1s^2\), helium.

Step 1: Identify Inert Gases

Inert gases, also known as noble gases, are elements in Group 18 of the periodic table. They all have full valence electron shells, making them chemically stable. Their configurations typically end with a filled 'p' subshell.

Step 2: Examine Given Configurations

The provided electron configurations are:(a) \(1s^2\), belongs to helium; (b) \(1s^2 2s^2 2p^6\), belongs to neon; (c) \(1s^2 2s^2 2p^5\), belongs to fluorine (not an inert gas); (d) \(1s^2 2s^2 2p^6 3s^2 3p^6\), belongs to argon.

Step 3: Eliminate Non-Inert Option

Among the options, only (c) \(1s^2 2s^2 2p^5\) does not represent an inert gas because it is one electron short of the neon configuration.

Step 4: Identify Smallest Size and Highest Ionization Energy

Ionization energy tends to increase with decreasing atomic size and increasing nuclear charge, especially across a period where electrons are held more tightly. Among the various inert gas configurations identified:- Helium \((1s^2)\) has the smallest size and the highest ionization energy as it has only one shell.

Key Concepts

Electron Configuration

Electron configuration describes how electrons are distributed in an atom's electron shells. Each shell can hold a set number of electrons, and these are added in a specific sequence.
Understanding electron configuration helps predict an element's chemical properties, reactivity, and the nature of its bonds. For example:

• Helium: Its configuration is written as \( 1s^2 \), where both electrons fill the first energy level (the 's' orbital).
• Neon: With a configuration of \( 1s^2 2s^2 2p^6 \), it fills the first and second energy levels, illustrating its stable and noble gas characteristics.
• Argon: It has a configuration of \( 1s^2 2s^2 2p^6 3s^2 3p^6 \), with filled first, second, and third energy levels.

Atoms strive for full electron shells, often achieving an inert, noble gas configuration. This stability explains why noble gases are chemically non-reactive and are found at the far right of the periodic table.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is a key factor influencing an element's reactivity. Simply put, the higher the ionization energy, the more difficult it is to remove an electron. This typically means:

• Atoms with high ionization energy hold onto their electrons more tightly.
• Elements with a filled valence shell, like noble gases, have extremely high ionization levels.

For instance, helium has the highest ionization energy among all elements because of its small atomic size and effective nuclear charge. As you move across a period in the periodic table, ionization energy generally increases because electrons are added to the same shell, increasing the nuclear charge without increasing the shell size.

Periodic Table Groups

The periodic table is structured in a way that groups elements with similar properties in vertical columns. These are known as groups or families. Each group shares a common number of valence electrons, determining the elements' chemical behavior.
Noble gases, found in Group 18, are renowned for being unreactive due to their complete valence shell. This complete electron configuration means:

• They are stable and usually do not form chemical bonds.
• They have significant applications, like in lighting and inert atmospheres for chemical reactions.

While most groups are numbered 1 through 18, the properties within these groups can vary as one descends the table. However, noble gases consistently display similar attributes due to their full outer electron shells, emphasizing their stability across the group.