Neeraj Anand, Param Anand

Anand Classes Notes explain the arrangement of elements in the periodic table follows a specific order based on the filling of electrons in atomic orbitals. From the fourth period onward, the pattern becomes more complex due to the involvement of d- and f-block orbitals. These periods include not only s- and p-block elements but also transition and inner transition elements. Understanding how these orbitals are filled explains why some periods have 18 or even 32 elements, unlike the shorter first three periods.

Anand Classes explains the periodic table is not just a list of elements—it’s a powerful tool that reveals the underlying pattern in chemical behavior. This repeating pattern, known as periodicity, occurs because elements with similar valence shell electron configurations exhibit similar chemical and physical properties. As we move across periods or down groups, these patterns repeat at regular intervals due to the systematic arrangement of electrons, especially in the outermost shell. Understanding this cause of periodicity helps explain why elements in the same group behave alike in reactions and form similar compounds.

Anand Classes Notes explains Atomic radius refers to the distance from an atom’s nucleus to its outermost electron shell, measured in different ways depending on context—covalent radius, metallic radius, or van der Waals radius. Understanding these radius types and their trends is vital for solving Class 11, NEET, and JEE chemistry problems on periodic properties and atomic structure.

Anand Classes Notes explains in atoms with more than one electron, the outer electrons are repelled by the inner electrons due to electron–electron repulsion. This repulsion reduces the full attractive force of the positively charged nucleus on the valence electrons. This phenomenon is known as the screening effect or shielding effect. The greater the number of inner electrons, the larger the shielding effect, which decreases the effective nuclear charge (Zeff) experienced by the outermost electrons. Slater, a scientist, formulated rules to calculate this shielding effect quantitatively, which are essential for understanding atomic structure and periodic trends.

Anand Classes Notes explains across any period in the periodic table, atomic radius generally decreases from left to right. This trend occurs because as the atomic number increases, the nuclear charge grows while electrons are added to the same principal energy level. Without significant shielding from electrons in the same shell, the increased effective nuclear pull draws electrons closer to the nucleus, resulting in a smaller atomic size.

Anand Classes explains the atomic radius of elements shows a clear trend when we move down a group in the periodic table. For both alkali metals (Group 1) and halogens (Group 17), the atomic radius increases as we go from top to bottom. This happens because each successive element down the group has an additional electron shell in its electronic configuration. The added shell increases the distance between the outermost electrons and the nucleus, while the shielding effect of inner electrons reduces the effective nuclear pull. As a result, the atoms become larger despite the increasing nuclear charge.

Anand Classes explains the Van der Waals radius is a measure of the size of an atom when it is not bonded to another atom but held together by weak Van der Waals forces in the solid state. Unlike covalent radius, which is measured for bonded atoms, Van der Waals radius is determined for non-bonded atoms in neighbouring molecules. This concept is especially important for noble gases, which do not form covalent bonds and therefore have their atomic radii expressed in terms of Van der Waals radii.

Anand Classes explains Ionic radius is the measure of the size of an ion, defined as the distance from the nucleus of the ion to the outermost electron shell. When an atom loses electrons to form a cation, the number of protons exceeds the number of electrons, pulling the remaining electrons closer to the nucleus and reducing the ionic radius. In contrast, when an atom gains electrons to form an anion, the increased electron–electron repulsion expands the outer shell, increasing the ionic radius. Understanding ionic radius trends across a period and down a group is important for predicting chemical reactivity, bond strength, and physical properties, making it a key topic for JEE, NEET, and CBSE Board examinations.

Anand Classes explains the radius of a cation is always smaller than that of its parent atom. This is because a cation is formed when one or more electrons are lost from a gaseous atom, often resulting in the complete removal of the outermost electron shell. For example, when a sodium atom (Na) loses its single 3s electron to form a Na+ ion, the entire 3s shell disappears, leading to a significant reduction in size—from 186 pm in Na to just 95 pm in Na+. This size decrease occurs because the nuclear charge remains the same while the number of electrons decreases, causing the remaining electrons to be pulled closer to the nucleus due to the increased effective nuclear charge.