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 ↓


Hybridization and its types


The fascinating concept of hybridisation plays a vital role in understanding chemical bonding and molecular structure. Hybridisation is a theoretical model that describes the rearrangement of atomic orbitals to form new hybrid orbitals that are suitable for the pairing of electrons to form chemical bonds, usually covalent bonds. Let's delve deeper into the concept of hybridisation, explore its types and understand how it helps in predicting the geometry of molecules.

What is hybridization?

Hybridization is the process of combining atomic orbitals into new orbitals called hybrid orbitals. This concept was introduced to explain the structure of molecules such as methane (CH4), which could not be explained satisfactorily using existing notions of atomic structure. Hybrid orbitals are degenerate, which means they have the same energy level, and they share characteristics of the constituent atomic orbitals.

Why is hybridization important?

Hybridisation helps to explain the following:

  • Molecular geometry of compounds.
  • Bond angles in molecules.
  • Stability and strength of chemical bonds.

Basic principles of hybridization

There are several key principles about hybridisation that need to be understood:

  1. Hybrid orbitals are formed by the mixing of atomic orbitals located on the same atom.
  2. The number of hybrid orbitals formed is equal to the number of mixed atomic orbitals.
  3. The hybrid orbitals are oriented in such a way that they maintain maximum distance from each other, and minimize electron pair repulsion.

Types of hybridization

Hybridization can result in different geometries and types of bonds. The type of hybridization depends on the number and type of orbitals involved in the mixing. Here are the main types:

1. sp hybridisation

In sp hybridization, an s orbital and a p orbital combine to form two equivalent sp hybrid orbitals. Each orbital is of 50% s and 50% p character. In this type the orbitals are oriented at 180 degrees relative to each other, forming a linear geometry.

Example: Acetylene (C2H2)

In acetylene, the carbon atoms are sp hybridised. Each carbon atom forms two sp hybridised orbitals. One of these forms a sigma bond with hydrogen and the other forms a sigma bond with the other carbon atom. The unhybridised p orbitals overlap each other to form pi bonds, resulting in a triple bond between the carbon atoms.

C - H : sp ≡ H - C : sp
H H SP SP

2. sp2 hybridisation

In sp2 hybridisation, one s orbital mixes with two p orbitals, resulting in the formation of three equivalent sp2 hybrid orbitals. These orbitals are arranged in a planar triangular configuration with an angle of 120° between them.

Example: Ethylene (C2H4)

In ethylene, each carbon atom is sp2 hybridized, and sp2 hybridized orbitals form sigma bonds with hydrogen and the other carbon. The remaining unhybridized p orbitals overlap to form a pi bond, forming a double bond between the carbon atoms.

H - C = C - H sp2
H H sp2 sp2

3. sp3 hybridisation

In sp3 hybridization, one s orbital mixes with three p orbitals, forming four equivalent sp3 hybrid orbitals. These orbitals adopt a tetrahedral shape with a bond angle of 109.5°.

Example: methane (CH4)

In methane, carbon is sp3 hybridised, with four sp3 hybrid orbitals forming sigma bonds with hydrogen.

H - C - H sp3
H H H H

4. sp3d hybridisation

In sp3d hybridization, one s, three p, and one d orbital combine to form five sp3d hybrid orbitals. These form a trigonal bipyramidal arrangement.

Example: Phosphorus pentachloride (PCl5)

In phosphorus pentachloride, phosphorus is sp3d hybridized, and it forms five equivalent bonds with chlorine, forming a trigonal bipyramidal shape.

P / |  Cl Cl Cl | | Cl Cl
Chlorine Chlorine Chlorine Chlorine Chlorine Chlorine

5. sp3d2 hybridisation

In sp3d2 hybridization, one s, three p, and two d orbitals combine to form six equivalent hybrid orbitals. This results in an octahedral arrangement.

Example: Sulfur hexafluoride (SF6)

In sulfur hexafluoride, sulfur is sp3d2 hybridized. It forms six bonds with fluorine in an octahedral shape, with 90° bond angles.

F | FSF | F  | / F
F F F F F F

How to determine hybridization?

Step 1: Count the valence electrons

Start with the number of valence electrons on the central atom. Consider the effect of surrounding atoms and any charge on the molecule or ion.

Step 2: Determine the molecular geometry

Use VSEPR (valence shell electron pair repulsion) theory to predict the geometry of the molecule. This helps to predict how many hybrid orbitals are needed.

Step 3: Use the hybridization shortcut

The regions of electron density around the central atom can help predict hybridization. Here are some guidelines:

  • 2 area: sp
  • 3 region: sp2
  • 4 region: sp3
  • 5 region: sp3d
  • 6 region: sp3d2

Hybridisation in different compounds

Water (H2O)

The structure of water is bent and its bond angle is 104.5°. Oxygen in water is sp3 hybridized, with two lone pairs and two single bonds with hydrogen.

H - O - H sp3

Ammonia (NH3)

The nitrogen atom in ammonia is sp3 hybridized. It forms three sigma bonds with hydrogen and retains a lone pair, resulting in a trigonal pyramidal shape.

H - N - H | H

Boron trifluoride (BF3)

Boron trifluoride shows sp2 hybridization, in which boron forms three sigma bonds with fluorine. This results in a planar triangular shape with a 120° bond angle.

F | BF | F

Importance of hybridization in chemistry

Hybridization is an important concept in understanding how molecules adopt certain shapes and bond types, which affect the physical and chemical properties of chemical reactions. It offered a coherent model that correlated with experimental data, helping to predict molecular structures and reactivity.

While hybridization is a valuable theoretical tool, it is important to understand that it oversimplifies the true quantum mechanical nature of electrons in atoms. Nevertheless, it remains a cornerstone in chemical education for explaining molecular geometry and bonding characteristics.

Conclusion

Hybridization laid the foundation for understanding how atoms combine to form molecules with specific geometries. By mixing atomic orbitals, hybridization explains the shapes, bond angles, and bond patterns of molecules, ultimately shaping the world of chemistry.


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Hybridization and its types

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