Why Fluorine (F) has a more negative electron gain enthalpy than Oxygen (O) ?
because: (a) It has a 2s²2p⁵ configuration, needing only one electron to complete its octet.
(b) Its small size and high nuclear charge attract the extra electron strongly.
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Anand Classes explains that successive electron gain enthalpies (EA₁, EA₂) represent the energy changes when one or more electrons are added to a gaseous atom. The first electron gain enthalpy (EA₁) is usually negative because a neutral atom attracts the incoming electron and energy is released. However, the second electron gain enthalpy (EA₂) is always positive, since the electron is being added to a negatively charged ion and experiences strong Coulombic repulsion, requiring external energy input. For example, in oxygen, EA₁ = –141 kJ mol⁻¹ (energy released), while EA₂ = +780 kJ mol⁻¹ (energy absorbed).
Anand Classes explains that the electron gain enthalpy of second-period elements (such as N, O, F) is generally less negative than their third-period counterparts (such as P, S, Cl). This trend arises because second-period elements are very small in size, and when an extra electron is added, it experiences strong electron–electron repulsions in the compact 2p orbitals. In contrast, third-period elements have larger atomic sizes and more diffused orbitals, so the incoming electron experiences less repulsion and greater attraction from the nucleus, resulting in more negative electron gain enthalpy values. This concept is important for JEE, NEET, and CBSE Class 11 Chemistry as it explains anomalies in periodic trends.
Anand Classes explains the important concept of electron gain enthalpy of halogens in a simple yet detailed way for JEE, NEET, and CBSE students. The trend of electron gain enthalpy in halogens (F, Cl, Br, I, At) is an essential topic of the periodic table, often asked in competitive exams. While generally electron gain enthalpy becomes less negative down the group, fluorine shows an anomaly where its value is less negative than chlorine. Understanding this exception, along with the overall trend, helps students strengthen their grasp on periodic properties and improve exam performance.
Anand Classes explains the concept of electron gain enthalpy of noble gases, beryllium (Be), magnesium (Mg), nitrogen (N), and phosphorus (P) in detail for JEE, NEET, and CBSE examinations. While most elements have negative electron gain enthalpy due to their tendency to attract electrons, noble gases exhibit positive electron gain enthalpy because of their stable octet configuration (ns2np6). On the other hand, Be and Mg with completely filled s-orbitals, and N and P with half-filled p-orbitals, show very low or nearly zero electron gain enthalpy values due to their extra stability. Understanding these exceptions is crucial for mastering trends in the periodic table and is a frequently asked concept in competitive exams.
Anand Classes brings you a detailed explanation of why halogens have the highest negative electron gain enthalpies in the periodic table. Halogens (Group 17 elements) like fluorine, chlorine, bromine, iodine, and astatine possess the electronic configuration ns²np⁵, which makes them just one electron short of achieving the stable noble gas configuration. This unique property gives halogens a very strong tendency to accept an additional electron, releasing a large amount of energy in the process. As a result, their electron gain enthalpies are highly negative, a key concept often asked in JEE, NEET, and CBSE Class 11 Chemistry exams.
Anand Classes provides the best study material for JEE, NEET, and CBSE students. In this article, we explain one of the most important periodic properties – Electron Gain Enthalpy. It refers to the energy change when an atom gains an extra electron to form a negative ion. Understanding the periodic trends of electron gain enthalpy (variation across a period and down a group) is crucial for competitive exams like JEE Main, JEE Advanced, NEET, and Board examinations, as it helps in predicting the reactivity and chemical behavior of elements.
Anand Classes brings you a detailed explanation of the important topic Factors Affecting Electron Gain Enthalpy for JEE, NEET, and CBSE Class 11 Chemistry. Electron gain enthalpy refers to the energy change when an isolated gaseous atom gains an extra electron to form a negative ion. Understanding how nuclear charge, atomic size, electronic configuration, and shielding effect influence electron gain enthalpy is crucial for mastering periodic trends and solving exam-based questions. This article provides clear theory, examples, FAQs, and comparisons to help students build strong conceptual clarity and score high in competitive as well as board examinations.
Anand Classes brings you a clear and detailed explanation of Electron Gain Enthalpy (Electron Affinity), an important concept in Chemistry for JEE, NEET, and CBSE Class 11 students. Electron gain enthalpy is defined as the energy change that occurs when an electron is added to an isolated gaseous atom to form a negatively charged ion. Understanding this concept helps students grasp the reactivity trends of elements, periodic variations, and the stability of ions, making it a crucial topic for competitive exams.
Anand Classes presents detailed study material on Ionization Enthalpy (Periodic Table) Class 11 Important Conceptual NCERT Question Answers for JEE, NEET, and CBSE Board students. Ionization enthalpy is a fundamental periodic property that explains the energy required to remove an electron from an isolated gaseous atom. Understanding this concept not only strengthens your basics of atomic structure but also helps in solving higher-order conceptual and numerical problems. In this article, you will find well-explained answers to NCERT questions, along with extra tips, solved examples, and exam-oriented insights.
Anand Classes provides a clear and detailed explanation of the variation of ionization enthalpy down a group in the periodic table, an important topic for JEE, NEET, and CBSE Class 11 Chemistry. As we move from the top to the bottom of a group, ionization enthalpy gradually decreases due to factors such as an increase in atomic size, an increase in the shielding effect, and a relatively smaller impact of increasing nuclear charge. Understanding this periodic trend helps students predict the chemical reactivity, metallic character, and bonding nature of elements, making it a crucial concept for competitive exams and board preparation.
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.
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 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.
