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 → Chemical Bonding and Molecular Structure ↓


Bond parameters


In chemistry, it is important to understand how atoms join together to form molecules. This bonding is determined by several factors, known as bond parameters, which help us understand the nature of chemical bonds. These parameters include bond length, bond angle, bond energy, bond order, and polarity. Gaining insight into these parameters helps students predict the molecular structure and properties of substances.

Bond length

Bond length is the distance between the nuclei of two bonded atoms. It is an important indicator of the strength and stability of a bond. Bond length can be affected by the size of the atoms and the bond order. Generally, as the size of the atoms increases, so does the bond length. Conversely, as the bond order increases, the bond length decreases. This is because double bonds are shorter than single bonds, and triple bonds are shorter than double bonds.

Example: The bond length of a single bond CC is approximately 154 pm. The bond length of a double bond C=C is approximately 134 pm. The bond length of a triple bond C≡C is approximately 120 pm.

Looking at bond length

Single Bonds Double Bond Triple Bond

Bond angle

The bond angle is the angle formed between three atoms in at least two bonds. Like bond length, bond angles are also important because they determine the shape of the molecule. VSEPR (valence shell electron pair repulsion) theory helps predict bond angles. According to this theory, electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion forces.

Example: In a water molecule (H2O), the bond angle between HOH is approximately 104.5°. In carbon dioxide (CO2), the bond angle between OCO is 180°, making it linear.

Visualization of bond angles

H H O

Bond energy

Bond energy refers to the amount of energy needed to break one mole of bonds in gaseous molecules. It is a direct measure of bond strength: the higher the bond energy, the stronger the bond. Different bonds have different bond energies affected by factors such as bond order and bond length. Stronger bonds with higher bond orders generally have higher bond energies.

Example: The bond energy of a CH bond is about 413 kJ/mol. For a C=O double bond, it is approximately 743 kJ/mol.

Visualization of bond energy

Single Bonds Double Bond Triple Bond

Bond order

Bond order is an indication of the number of chemical bonds between a pair of atoms. Higher bond orders indicate stronger, shorter bonds. It is calculated by dividing the difference between the number of bonding and antibonding electrons by two. In simple diatomic molecules or entities, it can be equal to the number of shared electron pairs.

Example: For a single bond (HH) in H2, the bond order is 1. For a double bond (O=O) in O2, the bond order is 2. For a triple bond (N≡N) in N2, the bond order is 3.

Looking at bond order

Order 1 Order 2 Order 3

Bond polarity

Bond polarity occurs when there is a difference in electronegativities between bonded atoms. In polar covalent bonds, electrons are not shared equally, resulting in a partial charge. Nonpolar bonds involve equal sharing of electrons. The degree of polarity affects the physical properties of compounds, such as their state, boiling point, and solubility.

Example: In an H-Cl bond, the chlorine atom is more electronegative, making it partially negative, while hydrogen is partially positive. In an NH bond in ammonia (NH3), nitrogen is more electronegative than hydrogen, creating a dipole.

Visualization of bond polarity

δ+ δ-

Conclusion

Understanding bond parameters is the key to understanding how and why substances behave the way they do. By analyzing bond lengths, bond angles, bond energies, bond orders, and polarity, we can predict and explain the structures and properties of various chemical compounds. This insight is fundamental for students as they advance in the study of chemistry and move on to deal with more complex molecular interactions and reactions.


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Bond parameters

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