Chapter 4, Problem 4.104QP

Interpretation Introduction

Interpretation: From the given set of option the option that perfectly matches with the relatively sizes of isoelectronic ions of the given elements should be determined.

Concept Introduction:

Periodic Table: The available chemical elements are arranged considering their atomic number, the electronic configuration and their properties. The elements placed on the left of the table are metals and non-metals are placed on right side of the table.

In periodic table the horizontal rows are called periods and the vertical column are called group. There are seven periods and 18 groups present in the table and some of those groups are given special name as follows,

$\begin{array}{l}\text{Group-1}\to \text{Alkali metal}\\ \text{Group-2}\to \text{Alkaline metals}\\ \text{Group-16}\to \text{Chalcogens}\\ \text{Group-17}\to \text{Halogens}\\ \text{Group-18}\to \text{Noble gases}\end{array}$

Atomic Number: Atomic number of the element is equal to the number of protons present in the nucleus of the element which is denoted by symbol Z. The superscript presents on the left side of the symbol of the element.

Atomic radius:

Atomic radius is the distance between the atomic nucleus and outermost electron of an atom.  From the atomic radius, the size of atoms can be visualized.  But there is no specific distance from nucleus to electron due to electron cloud around the atom does not have well-defined boundary.

Cation: Removal of electron from the atom results to form positively charged ion called cation.

Anion: Addition of electron to atom results to form negatively charged ion called anion.

The net charge present in the element denotes the presence or absence of electrons in the element.

In periodic table the horizontal rows are called periods and the vertical column are called group.

Electronic configuration: It is used to denote the distribution of electrons placed over orbitals that present in the atom.

The rules followed by the electrons are as follows,

The electrons gets distributed starting from lower energy orbital to higher energy orbital, pairing of electrons starts after all the orbitals are singly filled and finally, no two electrons that have same spin can occupy the same orbital.

Isoelectronic species: Two chemical species are considered as isoelectronic if they both contain equal number of electrons with them.