Problem 24 - A Complete Resource Book in Chemistry for JEE Main

Question

The order of first ionization energies of the elements \(\mathrm{Li}, \mathrm{Be}, \mathrm{B}, \mathrm{Na}\) is (a) \(\mathrm{Be}>\mathrm{Li}>\mathrm{B}>\mathrm{Na}\) (b) \(\mathrm{B}>\mathrm{Be}>\mathrm{Li}>\mathrm{Na}\) (c) \(\mathrm{Na}>\mathrm{Li}>\mathrm{B}>\mathrm{Be}\) (d) \(\mathrm{Be}>\mathrm{B}>\mathrm{Li}>\mathrm{Na}\)

Step-by-Step Solution

Answer

The correct order is (d) \( \mathrm{Be} > \mathrm{B} > \mathrm{Li} > \mathrm{Na} \).

Step 1: Understanding Ionization Energy

First ionization energy is the energy required to remove the outermost electron from a neutral atom in the gaseous state. Generally, ionization energy increases across a period due to increasing effective nuclear charge and decreases down a group due to increasing atomic size.

Step 2: Identify Position in the Periodic Table

Locate the elements in the periodic table: Lithium (Li) and Sodium (Na) are in Group 1, Beryllium (Be) and Boron (B) are in Group 2 and 13 respectively. Be and B are in the same period, while Na is in the next period compared to Li.

Step 3: Compare Elements in the Same Period

In the same period (like Be and B), ionization energy generally increases from left to right. Thus, Be has a higher ionization energy than B.

Step 4: Compare Elements in the Same Group

Within a group, ionization energy decreases as the atomic number increases. Therefore, Li has a higher ionization energy than Na since they are both in Group 1 and Li is above Na.

Step 5: Integrating Periodic Trends

Combine the trends: Be > B within the period, and Li > Na within the group. Now, we must decide between elements in the same row or across groups. Be > Li generally because it is to the right in the same period, even with exceptions in Boron's placement.

Step 6: Final Ordering

According to the above comparisons, the order is: Be > B > Li > Na.

Key Concepts

Periodic Trends

Understanding periodic trends, especially when it comes to ionization energy, is key when analyzing elemental behaviors across the periodic table. Ionization energy, the energy needed to remove an electron from an atom, generally follows predictable patterns. Across a period (left to right), ionization energy tends to increase. This is because atoms have more protons and an increasing effective nuclear charge, pulling the electrons closer and making them harder to remove.

Conversely, moving down a group (top to bottom), ionization energy decreases. This is due to the additional electron shells that increase the distance between the outer electrons and the nucleus, making it easier to remove these electrons. Understanding these trends helps us to correctly predict the order of ionization energies for elements, as demonstrated in the original exercise.

Periodic Table

The periodic table is a chart that organizes all known elements according to increasing atomic number. It also reveals patterns in chemical properties and behaviors. The table is arranged into rows called periods and columns called groups (or families). The elements in these rows and columns show periodicity in their properties, meaning they show repeating patterns.

Elements that fall in the same group often display similar chemical behaviors. For example, Group 1 elements (alkali metals like lithium and sodium) are known for their high reactivity, while Group 2 (alkaline earth metals like beryllium) are slightly less reactive. Understanding how elements are positioned helps to predict their properties, including their ionization energies.

Group 1 and Group 2 Elements

Group 1 and Group 2 elements have distinct properties and positions on the periodic table. Group 1 elements, or alkali metals, include lithium (Li) and sodium (Na). These metals are highly reactive and have relatively low ionization energies compared to elements across the table. As you move down Group 1, from Li to Na, ionization energy decreases due to increased atomic size and electron shielding.

Group 2 elements, known as alkaline earth metals, include beryllium (Be). These elements are less reactive than alkali metals and have slightly higher ionization energies. Within Group 2, beryllium has a notably high ionization energy, especially compared to boron (B) from period 2 due to its smaller atomic size and higher effective nuclear charge. Grasping these group characteristics aids in understanding the trends in ionization energies and the reasoning behind the order determined in the exercise.

Atomic Structure

Atomic structure is crucial in determining an element's properties, including its ionization energy. Atoms consist of a nucleus, containing protons and neutrons, surrounded by electrons in various shells or energy levels. The number of protons (atomic number) determines the element's identity and properties.

Electrons occupy energy levels with those in the outermost shell being the valence electrons. These electrons are important in chemical reactions and ionization processes. The effective nuclear charge, which is the net positive charge experienced by an electron in a multi-electron atom, affects how strongly an outer electron is held. The higher the effective nuclear charge, the more tightly an electron is held, and thus, the higher the ionization energy. Understanding atomic structure helps explain trends such as why beryllium has a higher ionization energy than boron.