Problem 26 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

Which of the following atoms has the highest first ionization energy? (a) \(\mathrm{K}\) (b) \(\mathrm{Sc}\) (c) \(\mathrm{Rb}\) (d) \(\mathrm{Na}\)

Step-by-Step Solution

Answer

Sodium ( \(\mathrm{Na}\)) has the highest first ionization energy.

Step 1: Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. Generally, ionization energy increases across a period and decreases down a group in the periodic table.

Step 2: Analyze the Options Based on the Periodic Table

We have the elements potassium (\(\mathrm{K}\), atomic number 19), scandium (\(\mathrm{Sc}\), atomic number 21), rubidium (\(\mathrm{Rb}\), atomic number 37), and sodium (\(\mathrm{Na}\), atomic number 11). All these elements are in groups 1 and 3, but different periods.

Step 3: Consider Periodic Trends

Since ionization energy increases across a period, elements on the far right have a higher ionization energy compared to elements on the left within the same period. Potassium (\(\mathrm{K}\), Period 4) and rubidium (\(\mathrm{Rb}\), Period 5) are in the same group, so rubidium will have a lower ionization energy than potassium because it is further down a group. Similarly, sodium (\(\mathrm{Na}\), Period 3) will have higher ionization energy than rubidium (Period 5).

Step 4: Identify the Highest Ionization Energy

Among the options, sodium (\(\mathrm{Na}\)) is in the third period, which is higher up in the periodic table compared to potassium (\(\mathrm{K}\)) and rubidium (\(\mathrm{Rb}\)), indicating it has a higher ionization energy as a group 1 element. Although scandium (\(\mathrm{Sc}\)) is from group 3, sodium, being a period 3 element, has relatively higher ionization energy due to its position among the given elements.

Key Concepts

Periodic Table Trends

The periodic table is a powerful tool that helps predict various properties of elements, such as ionization energy. Elements are arranged in order of increasing atomic number and are structured into periods (horizontal rows) and groups (vertical columns). As you move across a period from left to right, ionization energy generally increases. This is because the increased nuclear charge pulls the electrons closer, requiring more energy to remove an electron.
However, as you move down a group, ionization energy tends to decrease. This happens because the outer electrons are further away from the nucleus and are more shielded by inner electrons, making it easier to remove them. This predictable variation in properties as one moves across a period or down a group is known as a periodic trend.
Understanding these trends helps in predicting the chemical behavior of elements, such as reactivity and the type of bonds they form.

First Ionization Energy

First ionization energy is the minimum energy required to remove the outermost electron from a neutral atom in its gaseous state. When considering first ionization energy, we focus on how tightly an electron is held within an atom. Several factors impact it:

• Atomic Size: Smaller atoms generally have higher ionization energies because their electrons are closer to the nucleus.
• Nuclear Charge: Atoms with a higher positive charge in the nucleus will attract electrons more strongly, increasing ionization energy.
• Electron Shielding: Inner electrons can shield outer electrons from the full effect of the positive nuclear charge, which can decrease ionization energy.

As you analyze elements from the periodic table, keep in mind that the first ionization energy increases across a period due to increasing nuclear charge and decreases down a group due to increased electron shielding and distance from the nucleus.

Periodic Groups and Periods

The periodic table's structure is defined by groups and periods, providing a roadmap for understanding elemental properties. Groups are the columns, numbered 1 through 18, and each group shares similar chemical properties. For instance, all elements in group 1, also known as alkali metals, have one electron in their outer shell.
Periods, on the other hand, are the horizontal rows. Each period corresponds to a principal energy level filled with electrons. As you move across a period, elements gain electrons and protons, leading to varying chemical behaviors.
In our original exercise's context, sodium (\(\mathrm{Na}\)) stands out because it's located in period 3 and group 1. Despite this, it has higher ionization energy compared to potassium (\(\mathrm{K}\)) and rubidium (\(\mathrm{Rb}\)) because it's further up in the periodic table. Each group's and period's unique characteristics help us predict their chemical and physical properties effectively.