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 ↓


VSEPR Theory


Valence shell electron pair repulsion (VSEPR) theory is a simple but effective method used in chemistry to predict the shape or geometry of molecules. This theory, proposed by Ronald Gillespie and Ronald Nyholm in the early 1960s, is essential for understanding molecular structures determined by electron pair interactions around the central atom in the molecule.

Fundamentals of vsepr theory

VSEPR theory is based on several fundamental principles that guide the prediction of molecular shapes:

  1. Electron pairs, both bonded and unbonded (lone pairs), repel each other and therefore arrange themselves as far apart as possible around the central atom.
  2. The geometry of a molecule is determined by the number and type of electron pairs present around the central atom.
  3. Lone pairs exert greater repulsion than bonding pairs, causing the bond angles to deviate from their ideal values.

Types of electron pairs

Before we explore molecular shapes, it is necessary to understand the two types of electron pairs considered in the VSEPR theory:

  • Bond pair: These are shared between two atoms to form a covalent bond.
  • Non-bonding pairs (lone pairs): These are not shared between atoms but belong to the same atom, which usually affect the shape more due to their repulsion force.

General molecular geometry

Below are common molecular geometries according to VSEPR theory along with visual examples:

Linear geometry

AB

Molecules of the type AB 2 have a linear geometry, where the two B atoms are 180 degrees away from the central A atom. An example of this is carbon dioxide, CO 2.

Trigonal planar geometry

ABB

Molecules like BF 3 have trigonal planar geometry. Here, the B atoms are spread out at 120 degrees in one plane.

Tetrahedral geometry

ABBB

Tetrahedral geometry, with a bond angle of about 109.5°, is seen in molecules such as methane, CH 4.

The effect of single joints

Lone pairs are important in determining the shape of the molecule. They occupy more space due to increased repulsion, which leads to variations in bond angles:

Bent geometry

Consider water, H 2 O The presence of two lone pairs on oxygen compresses the bond angle between the hydrogen atoms to about 104.5 degrees, whereas 109.5 degrees would be expected in a tetrahedral setup.

Triangular pyramid geometry

In ammonia, NH 3, a lone pair causes the molecule to adopt a trigonal pyramidal shape, reducing the ideal 109.5° angle to about 107°.

Steps to determine molecular geometry using VSEPR

  1. Write the Lewis structure: Determine the structure by arranging the electrons around the atoms to satisfy their octets.
  2. Identify electron pairs: Calculate the regions of electron density (bonding and nonbonding) around the central atom.
  3. Determine the electron pair geometry: Associate the total number of electron pair regions with its corresponding geometry.
  4. Account for lone pairs: Adjust the geometry to account for repulsion from lone pairs to determine the actual molecular shape.

Examples of molecular shape prediction

Example 1: Methane (CH 4)

Lewis Structure: Carbon is the central atom with four hydrogen atoms attached to it. No lone pair is found on carbon.

Classification: Four bond pairs, consistent with tetrahedral geometry.

Shape: Tetrahedral, bond angle of about 109.5°.

Example 2: Sulfur dioxide (SO 2)

Lewis Structure: Sulfur is the central atom with a lone pair and is bonded to two oxygen atoms.

Classification: Two bond pairs and one lone pair give the initial geometry of trigonal planar line.

Shape: Bent or V shaped due to lone pair repulsion, bond angle about 119 degrees.

Applications in chemistry

VSEPR theory is invaluable for predicting molecular geometry in covalent compounds. This insight is important for determining a compound's physical properties, reactivity, and designing new chemicals within a laboratory setting. Understanding molecular shapes aids in discovering interactions between different molecules, which is fundamental in fields such as drug design or the synthesis of new materials.


Grade 11 → 4.5


U
username
0%
completed in Grade 11

← Prev (4.4)
Bond parameters
Next (4.6) →
Valence bond theory

Comments 



VSEPR Theory

https://www.chemistryelite.com/g11/4.5?ceal=en