Problem 91 - A Complete Resource Book in Chemistry for JEE Main

Question

The first ionization potential of \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}\) and \(\mathrm{Si}\) are in the order: (a) \(\mathrm{Na}>\mathrm{Mg}>\mathrm{Al}<\mathrm{Si}\) (b) \(\mathrm{Na}>\mathrm{Mg}>\mathrm{Al}>\mathrm{Si}\) (c) \(\mathrm{Na}<\mathrm{Mg}<\mathrm{Al}<\mathrm{Si}\) (d) \(\mathrm{Na}<\mathrm{Mg}>\mathrm{Al}>\mathrm{Si}\)

Step-by-Step Solution

Answer

The correct order is option (d): \(\mathrm{Na}<\mathrm{Mg}>\mathrm{Al}<\mathrm{Si}\).

Step 1: Understanding Ionization Potential

The ionization potential (or ionization energy) is the energy required to remove an electron from a neutral atom in the gaseous phase. Generally, ionization energy increases across a period from left to right due to increasing nuclear charge, making it more difficult to remove an electron.

Step 2: Analyze Periodic Trends

Sodium (Na), Magnesium (Mg), Aluminum (Al), and Silicon (Si) all belong to the same period (Period 3) of the periodic table. According to periodic trends, we expect the ionization energy to increase from Na to Si. Therefore, typical order without exceptions would be Na < Mg < Al < Si.

Step 3: Consider Electron Configuration Exceptions

The electron configuration of each element can cause variations in expected trends. Mg's electron configuration is [Ne]3s², which is a full subshell and more stable than Al's [Ne]3s²3p¹, making Mg's ionization energy higher than Al's despite Al being further to the right.

Step 4: Examine Options

Based on stability due to filled sublevels (Mg) being higher than partially filled sublevels (Al), the order should reflect Mg > Al. The potential sequence considering ionic energies and the exceptions should be: Na < Mg > Al < Si. This aligns with option (d).

Key Concepts

Periodic Trends in Ionization Potential

Ionization potential, also known as ionization energy, is significantly influenced by periodic trends observed in the periodic table. This energy refers to the amount needed to remove an electron from an atom in its gaseous state. As a general trend, ionization energy increases as we move from left to right across a period. This is because the nuclear charge (or the positive charge of the nucleus) increases, leading to a stronger attraction between the nucleus and the electrons.

As this attraction strengthens, it becomes increasingly harder to remove an electron, thus elevating the ionization energy. For elements within the same period, such as sodium (Na), magnesium (Mg), aluminum (Al), and silicon (Si), this trend can be noted: ionization potential typically increases from Na to Si. However, exceptions due to electron configuration can occur, requiring a closer inspection to understand the nuances.

Electron Configuration and Its Impact

The arrangement of electrons within an atom, known as electron configuration, plays a crucial role in determining ionization potential. For instance, magnesium (Mg) has the electron configuration of [Ne]3s², which is a filled subshell. This configuration imparts stability, making it require more energy to remove an electron compared to an atom with a less stable configuration.

On the other hand, aluminum (Al) has an electron configuration of [Ne]3s²3p¹. The presence of an extra electron in the 3p sublevel makes Al less stable compared to Mg, resulting in a lower energy requirement to remove this electron. This is why Mg, despite being to the left of Al on the periodic table, has a higher ionization energy than Al. Understanding these configurations helps clarify why ionization trends might not strictly follow the expected trend across a period.

Navigating Periodic Table Trends

The periodic table is more than just a collection of elements; it offers crucial insights into elemental properties, such as ionization potential. Elements in the same period often show a gradual increase in ionization energy as one moves rightward due to increased nuclear charge. However, subtleties within this trend, like the stability associated with filled or half-filled subshells (as seen with Mg versus Al), can cause deviations.

Therefore, effectively navigating the periodic table involves recognizing these trends, but also understanding when and why exceptions occur. For elements like Na, Mg, Al, and Si, knowing their placement and considering electron configurations provides a full picture of why their ionization potentials are ordered as Na < Mg > Al < Si, rather than strictly increasing. This comprehensive approach allows students to grasp how periodic trends and electron configuration intricately interact.