Problem 45 - A Complete Resource Book in Chemistry for JEE Main

Question

Which corresponds to the most electropositive character? (a) \([\mathrm{Xe}] 6 \mathrm{~s}^{1}\) (b) \([\mathrm{Xe}] 6 \mathrm{~s}^{2}\) (c) \([\mathrm{He}] 2 \mathrm{~s}^{1}\) (d) \([\mathrm{He}] 2 \mathrm{~s}^{\mathrm{l}}\)

Step-by-Step Solution

Answer

Option (a) [Xe] 6s¹ corresponds to the most electropositive character.

Step 1: Understanding Electropositivity

The term 'electropositive' refers to an element's tendency to lose electrons and form positive ions. Electropositivity increases as you move down the groups and towards the left of the periodic table. Elements toward the bottom left are usually more electropositive.

Step 2: Analyze Element Configurations

Evaluate the given options: (a) [Xe] 6s¹: This is Francium, the most electropositive element. (b) [Xe] 6s²: This is Radium, which is less electropositive than Francium. (c) [He] 2s¹: This is Lithium, which is a highly reactive alkali metal. (d) [He] 2s¹ (likely a typo identical to option C).

Step 3: Compare Elements' Electropositive Character

Francium ([Xe] 6s¹) is the most electropositive element due to its position at the bottom of Group 1 in the periodic table. It is more likely to lose its single valence electron compared to other options.

Step 4: Conclusion

Given the periodic table trends, Francium with configuration [Xe] 6s¹ is the most electropositive element among the options provided.

Key Concepts

Periodic Table Trends

The periodic table is more than just a list of elements. It showcases significant patterns called trends. One of the crucial trends is how electronegativity and electropositivity vary across the table. As you move from left to right across a period, atoms generally become more electronegative. This means they are more inclined to attract electrons. In contrast, as you move down a group, atoms become more electropositive, meaning they are more willing to lose electrons.

Electronegativity and electropositivity are inversely related. As such, elements in the bottom left of the periodic table, like Francium, are highly electropositive. They readily lose electrons due to their larger atomic radii and lower ionization energies, making them prone to forming positive ions. Observing these trends can help you predict the chemical behavior of an element based on its position in the periodic table.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It's a measure of the energy needed to overcome the attraction between the negatively charged electron and the positively charged nucleus. This concept helps explain why certain elements are more reactive, especially the highly electropositive metals.

The weaker the attraction (lower ionization energy), the easier it is for an atom to lose an electron. As you go down a group in the periodic table, the ionization energy decreases. This happens because electrons are added to higher energy levels, which are further from the nucleus. For example, Francium has a low ionization energy, making it highly electropositive and likely to give up its lone valence electron. This makes understanding ionization energy essential when considering why certain elements are more reactive than others.

Valence Electrons

Valence electrons determine an atom's chemical properties and its ability to bond with other atoms. These are the outermost electrons of an atom and are crucial in forming chemical bonds. Each group of the periodic table has elements with the same number of valence electrons, explaining their similar chemical behavior.

In the exercise example, Francium ([Xe] 6s¹) has one valence electron in its outermost s-orbital. This single electron is relatively easy to remove due to its position in the periodic table. Elements with fewer valence electrons, especially those located to the left of the table, like alkali metals, are more likely to lose these electrons and display electropositive behavior. It's important to recognize the role of valence electrons in dictating how an element will interact chemically with others.