Problem 42 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

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

Step-by-Step Solution

Answer

The correct order is (a) Na < Mg > Al < Si.

Step 1: Understand Ionisation Potential

Ionisation potential (or ionisation energy) is the energy required to remove an electron from a gaseous atom or ion. It generally increases across a period in the periodic table from left to right due to increasing nuclear charge.

Step 2: Identify the Elements

The elements in the problem are Na (Sodium), Mg (Magnesium), Al (Aluminum), and Si (Silicon). These elements are consecutive members of period 3 in the periodic table.

Step 3: Predict Trends Across a Period

Across a period, ionisation energy generally increases due to the addition of protons, which increases the effective nuclear charge and pulls the electrons closer to the nucleus. This makes them harder to remove.

Step 4: Consider Exceptions to the Trend

Magnesium (Mg) has a filled s-orbital, and Aluminum (Al) starts filling the p-orbital. Despite the general trend, Mg has a slightly higher ionisation energy than Al due to the stable filled configuration in Mg.

Step 5: Order the Ionisation Potentials

Considering the general trend and the exception, the order of ionisation potentials should be: 1. Na (lowest, as it starts the period) 2. Al (slightly lower than Mg due to p-orbital) 3. Mg (slightly higher than Al because of filled s-orbital) 4. Si (highest, as it follows the general trend). Hence, the correct order is: Na < Mg > Al < Si.

Key Concepts

Periodic Trends

Periodic trends refer to the predictable patterns seen across periods and groups in the periodic table. As you move from left to right across a period, ionization energy tends to increase. Each successive element has more protons, thus a greater positive charge pulling electrons closer to the nucleus. This means more energy is needed to remove an electron from the outer shell. For instance, in period 3, this trend is seen when moving from Sodium (Na) to Silicon (Si). However, there are notable exceptions due to electron configurations that can alter these trends slightly.

Generally, understanding periodic trends helps in predicting and explaining the chemical behavior of elements in the periodic table, such as ionization energies, atomic radius, and electronegativity.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by valence electrons. It is essentially the charge that the outermost electrons feel after accounting for electron-electron repulsions that reduce the actual nuclear charge. Mathematically, it can be understood as the atomic number minus the shielding or screening effect provided by inner electrons.

• As we move across a period, the effective nuclear charge increases.
• This is because each adjacent element has more protons, pulling electrons in more tightly.
• Consequently, electrons are held more tightly and the ionization energy increases.

This concept is particularly useful in predicting how the electron configurations affect an atom's chemical properties and reactivity.

s-orbital Stability

The stability of s-orbitals is an important factor when considering ionization energies, especially when comparing elements within the same period or group. A filled s-orbital, such as in magnesium (Mg) with its 3s² configuration, is energetically favorable.

This stability makes Mg slightly more resistant to losing an electron compared to Al, which starts filling the less stable p-orbital. Thus, Mg often has a higher ionization energy than expected based solely on periodic trends. This phenomenon is an example of how electron orbital configurations can lead to exceptions in general periodic patterns.

p-orbital Electron Configuration

P-orbitals begin filling after the s-orbitals in each energy level, starting from the third period of the periodic table. Following the more stable s-orbitals, p-orbitals accommodate up to six electrons. The shift from having a stable s-orbital to beginning to fill a p-orbital can affect the ionization energy trend.

In our example, after magnesium (Mg), aluminum (Al) has the electron configuration of 3s² 3p¹. The single electron in the 3p-orbital is less tightly held than the paired 3s electrons due to the 3p's higher energy and less stable nature. Thus, this makes it easier to remove, resulting in a deviation from the expected trend where Al has a slightly lower ionization energy than Mg. Understanding p-orbital configuration is crucial for explaining the deviations in ionization energy trends across periods.