Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Classification of elements and periodicity in properties → Periodic trends in properties ↓


Atomic and Ionic Radius


Introduction

In the fascinating world of chemistry, it is important to understand the size of atoms and ions. The terms "atomic radius" and "ionic radius" refer to the size of an atom and an ion, respectively. These concepts are essential for predicting and explaining the behavior of elements in chemical reactions and bonding. Let us learn about these properties in detail and understand their periodic trends.

Atomic radius

The atomic radius of an element is essentially the distance from the center of the nucleus to the outermost shell of the electron cloud in the atom. This measurement helps chemists understand the size of the atom and how it changes across the periodic table. Because electron clouds are not clearly defined, atomic radius is determined in a number of ways, usually related to how atoms bond in different situations.

Typically, atomic radii are measured in picometers (pm) or angstroms (Å), with 1 Å equal to 100 pm.

Covalent radius

It is half the distance between two identical atoms bonded together. It is generally used for non-metals where the atoms form covalent bonds.

d_A = 2 * r_c :r_c = Covalent radius :d_A = Distance between two nuclei of identical atoms

Metal radius

It is useful for metals, where the atoms are very close to each other. It is half the distance between two adjacent atoms in the metal lattice.

Periodic trends in atomic radius

Atomic radius varies as you move across a period and down a group in the periodic table. Understanding these trends is fundamental in predicting how the elements interact.

Trends over a period

As we move from left to right in a period, the atomic radius decreases. This happens because the nuclear charge increases due to the addition of protons to the atomic nucleus. The increased positive charge pulls the electrons closer to the nucleus, which reduces the size of the atomic radius.

H He Took Happen B C N Hey F by

Downward trend in the group

As one goes down a group in the periodic table, the atomic radius increases. This trend is due to the addition of a new electron shell with each successive element. Although the effective nuclear charge increases, its effect is reduced by the increase in the average distance of the electrons from the nucleus.

Ionic radius

Ionic radius refers to the radius of an atom's ion and measures the size of the atom after losing or gaining one or more electrons. Ions can be classified into two types: cations and anions.

Cations

Cations are positively charged ions formed by the loss of electrons. As electrons are removed, the electron-electron repulsion in the resulting ion decreases, allowing the remaining electrons to be drawn closer to the nucleus. As a result, the ionic radius of cations is smaller than that of their parent atoms.

For example, consider sodium:

Na → Na⁺ + e⁻

The radius of Na + ion is smaller than that of the neutral sodium atom, since it has one less electron.

Anion

Anions are negatively charged ions formed by the gain of electrons. The addition of extra electrons increases electron-electron repulsion, causing the electron cloud to expand. Therefore, the ionic radius of anions is larger than that of their parent atoms.

For example, consider chlorine:

Cl + e⁻ → Cl⁻

In this case, the radius of Cl - ion is larger than that of the neutral chlorine atom.

Periodic trends in ionic radii

The trends in ionic radii are very closely related to those in atomic radii, but they also involve the charges on the ions.

Trends over a period

As we move across a period from left to right, first the cations and then the anions show decreasing ionic radius. This trend exists due to the increase in nuclear charge, which makes the ionic radius pack more tightly. Between the cation and the anion, there is usually a significant jump in size, as the anions start with a larger volume due to the extra electrons. For a better understanding, analyze the size of the ions in the second period:

Lee + be 2+ B 3+ C 4+ N3- O2- F -

Downward trend in the group

Going down a group in the periodic table, the ionic radius increases. This pattern reflects the trend of atomic radius increasing due to the addition of a new electron shell with each element. As a result, even as ions, the radius of these atoms will be much wider than their predecessors.

Consider the alkali metal ions, such as Li +, Na +, and K +, all of which show increasing ionic radius as we go down the group.

Visualization of periodic trends

Understanding periodic trends helps to predict and logically explain the behavior of an element. Here is a simple model using the atomic radius trend through a representative block of elements.

This diagram symbolically shows the trend in atomic radius along periods (horizontal) and groups (vertical). Moving to the right (across a period), the boxes shrink, indicating a decrease in atomic radius. Moving down the group, they increase in size.

Conclusion

In conclusion, understanding atomic and ionic radius trends is fundamental to the study of chemistry. These trends provide information about the properties of elements and help predict their chemical behavior. As you explore these concepts, remember that atomic and ionic radii go far beyond these generalizations, but serve as invaluable guides into the larger world of chemistry. Understanding and visualizing these trends gives us a deeper understanding of the nature of the elements in the periodic table.


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Atomic and Ionic Radius

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