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

When we talk about noble gases, we're referring to elements like Helium (He), Neon (Ne), and Argon (Ar), among others. These are special because they have a complete set of electrons in their outermost shell, which makes them uniquely stable among the elements. This electron configuration is often a baseline for understanding other elements since noble gases don't tend to react with other substances.

For students diving into chemistry, it's crucial to recognize that a noble gas electron configuration means a full s and p orbital at the last energy level for a given period. In the exercise provided, options (a), (b), and (d) represent such complete configurations, whereas option (c) is one electron short, indicating an incomplete valence shell and hence not a noble gas configuration.

Understanding the Basics

Elements strive to achieve this stable arrangement through chemical reactions, which often involves gaining, losing, or sharing electrons. The noble gas configuration is a goal for other elements to reach either a full or empty valence shell, generally leading to increased stability.

Ionization energy is the amount of energy required to remove the outermost electron from an atom in its gaseous state. It's a valuable indicator of an element's reactivity and is particularly high for noble gases due to their full valence shells — an electron is not easily removed since it would disrupt the stable configuration.

As students explore the periodic table, they'll notice that ionization energy generally increases from left to right across a period. This is due to the increased positive charge in the nucleus, which pulls electrons in closer and makes them harder to remove. In the exercise, we're asked to identify an inert gas with the highest ionization energy. Option (a), represented by Helium (He), tops this category because it not only has a full valence shell but is also located in the first period, where the effective nuclear charge is greatest on the electron due to fewer electron shields.

Reactivity and Trends

In practical terms, elements with high ionization energy are less likely to lose electrons, and thus, less reactive. Knowing ionization energy trends can help predict and understand chemical behavior and reactions.

The atomic size, or atomic radius, refers to the distance from the nucleus to the boundary of the surrounding cloud of electrons. It is a key concept that impacts other properties, such as ionization energy and electronegativity. For example, as you move down a group on the periodic table, atomic size increases because each successive element adds a new electron shell, making the atom larger.

However, it's important to note that within a period, atomic size decreases from left to right. This happens because, although all elements within a period have the same number of electron shells, the increasing nuclear charge across the period pulls the electrons in closer. In the exercise we've been studying, option (a), which is Helium, will display the smallest atomic size since it's the topmost element in the noble gas group with only one electron shell and a high nuclear charge pulling that shell tight.

Significance in Chemistry

The concept of atomic size is fundamental for understanding patterns in the periodic table and underpins principles of molecular structure and bonding. It's a factor in determining the distances between atoms within a molecule and the size of molecules and ions.