Problem 113 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following statements is/are correct? (a) Successive ionization energies of an atom always increase. (b) The electron affinity for sulphur is more exothermic than that for oxygen. (c) First ionization enegy of As is greater than that of Se. (d) Chlorine has larger atomic size as well as electron affinity than that of fluorine.

Step-by-Step Solution

Answer

Correct statements are (a) and (b). Statements (c) and (d) are incorrect.

Step 1: Understanding Ionization Energies

Successive ionization energies always increase because each electron removal decreases electron-electron repulsion but increases the effective nuclear charge on the remaining electrons, making them harder to remove. Thus, statement (a) is correct.

Step 2: Comparing Electron Affinities

Sulfur and oxygen are both in group 16 of the periodic table, but sulfur is lower down than oxygen. Electron affinities become more exothermic going down the group due to larger atomic size and better accommodation of additional electrons. Therefore, the electron affinity for sulfur is more exothermic than for oxygen, making statement (b) correct.

Step 3: First Ionization Energy Comparison

The first ionization energy of As (Arsenic) is less than that of Se (Selenium) due to Se having a smaller atomic radius and slightly higher effective nuclear charge, which means it holds onto its valence electrons more tightly. Consequently, statement (c) is incorrect.

Step 4: Analyzing Atomic Size and Electron Affinity

Chlorine is larger in size compared to fluorine as it is located lower in the periodic table. Despite its larger size, chlorine has a higher (more exothermic) electron affinity than fluorine, contrary to initial assumptions based on size. Therefore, statement (d) is only partially correct in terms of electron affinity and incorrect in terms of atomic size comparison.

Key Concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or ion in the gaseous state. This energy depends on several factors, including the atomic size and effective nuclear charge. Successive ionization energies refer to the energy needed to remove each subsequent electron after the first has been removed. With every electron removed, the remaining electrons experience a stronger attractive force from the nucleus as there are fewer electron-electron repulsions.

• As each electron is removed, the effective nuclear charge increases, making it harder to remove the next electron.
• Therefore, successive ionization energies always increase, and each step requires more energy than the last.

This is important because it explains the stability of electron configurations and why elements like the noble gases have such high ionization energies. Understanding this concept helps predict how and why atoms interact to form compounds.

Electron Affinity

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom in the gaseous state, forming an anion. Generally, elements with greater electron affinity release more energy when gaining an electron, making the process more exothermic.

• In the periodic table, electron affinity values tend to become more negative across a period and less negative going down a group.
• This is because, across a period, atoms having smaller atomic sizes and greater effective nuclear charges attract additional electrons more strongly.

For elements like sulfur and oxygen, sulfur's larger atomic radius allows for a better accommodation of additional electrons compared to oxygen, making its electron affinity more exothermic. A better understanding of electron affinity aids in predicting the chemical reactivity of elements in forming ionic bonds.

Atomic Size

Atomic size, or atomic radius, is the distance from the center of the nucleus of an atom to the outermost shell of electrons. This size tends to decrease across a period and increase down a group in the periodic table.

• As you move across a period from left to right, the effective nuclear charge increases, pulling the outer electrons closer to the nucleus and resulting in a smaller atomic size.
• When moving down a group, additional electron shells are added, increasing the atomic size despite a higher effective nuclear charge.

For example, chlorine is larger than fluorine because it has more electron shells. Understanding atomic size helps in predicting the physical and chemical properties of an element, such as its reactivity and the nature of its bonds in compounds.

Effective Nuclear Charge

Effective nuclear charge (Z_{eff}) is the net positive charge experienced by an electron in a multi-electron atom. It is the result of the shielding effect of other electrons, where inner electrons partially block the outer electrons from the full attractive force of the nucleus.

• Effective nuclear charge increases across a period because the number of protons in the nucleus increases while shielding remains relatively constant.
• Down a group, the effective nuclear charge does increase but not as significantly as it does across a period, due to the increase in shielding from additional electron shells.

Having a higher effective nuclear charge means electrons are held more tightly to the nucleus, influencing properties such as ionization energy and electronic structure. Understanding how Z_{eff} functions is crucial for predicting atomic behavior and interactions.