Anand Classes explains that in every atom, electrons are held in place by the strong electrostatic attraction of the positively charged nucleus. This attraction is due to the opposite charges between the negatively charged electrons and the positively charged protons in the nucleus. The outermost or most loosely bound electrons experience a weaker pull compared to the inner electrons because they are farther from the nucleus and shielded by inner shells. To remove such an electron, energy must be supplied to overcome this attraction. The amount of energy needed for this process is called Ionization Enthalpy or Energy (IE) or Ionization Potential (IP). It serves as a quantitative measure in kJ mol⁻¹ of how easily an atom can lose an electron, providing important insight into the atom’s reactivity and chemical behavior.

Anand Classes presents a comprehensive collection of important questions and answers on atomic and ionic radii, specially designed for students preparing for JEE, NEET, and CBSE exams. This compilation focuses on understanding the size variations in atoms and ions, helping learners grasp key concepts such as nuclear charge, electron configuration, and periodic trends. Each question is paired with a detailed explanation, making complex topics easier to remember and apply during exams. Whether you are revising or practicing, these carefully crafted Q&As will strengthen your foundation in chemistry and boost your confidence.

Anand Classes Isoelectronic ions are ions of different elements that contain the same number of electrons but differ in their nuclear charges. A group of such ions is called an isoelectronic series. In any isoelectronic series, the electron count remains constant, but as the nuclear charge increases, the force of attraction between the nucleus and the electrons also increases. This stronger pull draws the electron cloud closer to the nucleus, resulting in a smaller ionic radius. Consequently, in an isoelectronic series, the size of the ions decreases as the nuclear charge increases.

Anand Classes explains the size of a negative ion (anion) is always larger than that of its corresponding neutral atom. This is because an anion is formed when a neutral atom gains one or more electrons, increasing the total number of electrons while the nuclear charge remains the same. The added electrons increase electron–electron repulsion and reduce the effective nuclear pull on each electron. As a result, the electron cloud expands, making the anion significantly larger than the parent atom. For example, chlorine (Cl) has an atomic radius of 99 pm, whereas the chloride ion (Cl⁻) swells to 181 pm.

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.

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 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 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 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 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 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 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 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 explain in the periodic table, elements are arranged in periods (horizontal rows) based on the way their electrons fill up different energy levels (shells). Each period starts with the filling of a new energy shell, and the number of elements in a period depends on how many orbitals are available in that shell to hold electrons. Let’s explore the electronic configurations of elements in the first three periods and understand why each period has a specific number of elements.

Anand Classes explain how To find the position of an element in the periodic table, we can use its electronic configuration. This helps us easily determine the period, group, and block to which the element belongs. In this post, we’ll go through simple solved examples to understand how to apply these rules step by step. This is very helpful for students preparing for JEE, NEET, and CBSE Class 11 exams.

Anand Classs Notes explain In Chemistry we can easily find the period, group, and block of an element in the periodic table by looking at its electronic configuration. This means, by knowing how the electrons are arranged in an atom, we can tell where the element is placed in the periodic table. It helps us understand the element’s properties and how it reacts with other elements.

Anand Classes explains the number of elements in each period of the periodic table follows the pattern 2, 8, 8, 18, 18, 32, 32 due to the way electrons fill atomic orbitals according to the Aufbau principle and the (n + l) rule. Each period corresponds to the filling of a new principal energy level, and the number of elements in that period depends on the total electron capacity of the subshells being filled. Understanding this explains why the 1st period has only 2 elements, the 2nd and 3rd periods have 8 elements each, the 4th and 5th periods have 18, and the 6th and 7th periods have 32 elements — an important concept for Class 11 Chemistry, JEE, and NEET.

The period and group number of an element in the periodic table can be determined from its electronic configuration using specific rules for s-, p-, d-, and f-block elements. These rules depend on the principal quantum number (n), the distribution of electrons in the outermost orbitals, and the block to which the element belongs. Knowing these rules is essential for solving periodic table questions in Class 11 Chemistry, JEE, and NEET.

Anand Classes Notes of Metals, non-metals, and metalloids are distributed in the periodic table according to their electronic configuration and chemical properties. Metals occupy the left and central portions of the table, non-metals are located on the right side, and metalloids lie along the zig-zag line separating the two. Understanding their position helps explain periodic trends such as metallic and non-metallic character, reactivity, and the types of oxides they form — an important topic in Class 11 Chemistry, JEE, and NEET.

The periodic table is divided into four main blocks — s-block, p-block, d-block, and f-block elements — based on the type of orbital that receives the last electron. Each block has unique position, electronic configuration, properties, and examples, making their comparison important for Class 11 Chemistry, JEE, and NEET preparation. This comparison covers their position in the periodic table, oxidation states, reactivity, metallic character, and special properties in a clear tabular format along with FAQs for quick revision.