Neeraj Anand, Param Anand

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.

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 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 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 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 explanation of the Factors Governing Ionization Enthalpy — the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Ionization enthalpy is a key concept in chemistry that reflects how strongly an atom holds its electrons. Its magnitude is influenced by several factors, including the size of the atom, nuclear charge, screening effect, penetration power of electrons, and the stability of the atom’s electronic configuration. Understanding these factors is essential for mastering periodic trends and predicting the chemical reactivity of elements.

Anand Classes explains that ionization enthalpy provides an excellent example for understanding the periodicity of elements in the periodic table. It is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state, and is expressed in kilojoules per mole (kJ mol⁻¹). By studying the variation of ionization enthalpy across periods and groups, we can clearly see the repeating patterns in element properties. The maxima in each period occur at noble gases, which have highly stable ns² np⁶ configurations, while the minima are found in alkali metals, which have a single loosely held ns¹ valence electron.