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In earlier chapters, you were introduced to the concepts of the physical world and matter, including their properties, various atomic theories, quantum theories of the atom, and the principles of electronic configuration. In this chapter, you will learn how atoms combine or bond, the different bonding theories, their types, and the underlying principles governing chemical compound formation.
Contents
- 1 Nature and Development of Chemical Bonding Theories
- 2 Kössel-Lewis Approach to Chemical Bonding
- 3 Octet Rule
- 4 Lewis Dot Structures
- 5 Covalent Bonding
- 6 Drawing Lewis Structures
- 7 Formal Charge
- 8 Ionic Bond Formation
- 9 Factors Favoring Ionic Bond Formation
- 10 Lattice Enthalpy
- 11 Bond Parameters
- 12 Covalent Bonds: Polar and Nonpolar
- 13 Dipole Moment
- 14 Valence Shell Electron Pair Repulsion (VSEPR) Theory and Molecular Shapes
- 15 Introduction to Valence Bond Theory
- 16 Types of Covalent Bonds Based on Overlap
- 17 Important Formulas in NCERT Notes Class 11 Chemistry (Part-I) Chapter 4: Chemical Bonding & Molecular Structure
- 18 FAQs
Explore Notes of Class 11 Chemistry
Nature and Development of Chemical Bonding Theories
Matter is composed of one or more types of elements. Under normal conditions, elements, except noble gases, do not exist as independent atoms in nature. Atoms usually occur in groups called molecules, which possess characteristic properties. This suggests the existence of a force that holds these atoms together, known as a chemical bond.
The concept of chemical bonding helps in understanding:
- Why do atoms combine?
- Why are only certain combinations possible?
- Why do some atoms react while others do not?
- Why do molecules possess definite shapes?
Kössel-Lewis Approach (1916)
Atoms combine to achieve noble gas configuration (octet or duplet). Bond formation occurs via:
- Electron transfer, forming ionic bonds (e.g., NaCl).
- Electron sharing, forming covalent bonds (e.g., Cl₂, H₂).
- Lewis symbols represent valence electrons as dots around the element’s symbol.
- The Octet Rule states that atoms tend to have eight electrons in their valence shell for stability.
- Limitations of this theory:
- Incomplete and expanded octets.
- Molecules with odd numbers of electrons.
- Inability to explain molecular shapes and relative stability.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Developed by Sidgwick and Powell; refined by Nyholm and Gillespie. It predicts molecular geometry based on the repulsion between valence electron pairs (bonded or lone). Electron pairs arrange themselves in space to minimize repulsion, determining the molecular shape.
Valence Bond (VB) Theory
Proposed by Heitler and London; extended by Pauling. Based on quantum mechanics and atomic orbital overlap.
- Covalent bonds are formed by the overlapping orbitals containing unpaired electrons with opposite spins.
- Explains directional properties of bonds and concepts like hybridization.
- Accounts for bond strength and geometry in polyatomic molecules (e.g., CH₄, NH₃).
Molecular Orbital (MO) Theory
Atomic orbitals combine to form molecular orbitals that extend over the entire molecule. Differentiates between bonding and antibonding orbitals. Explains bond order, magnetic behaviour, and delocalization of electrons in molecules.
The development of these theories has been closely associated with advances in understanding the structure of atoms, their electronic configurations, and the periodic classification of elements.
Objective of Bonding
Bond formation allows atoms to attain a more stable state by lowering the system’s energy, reflecting nature’s inherent tendency toward stability.
Kössel-Lewis Approach to Chemical Bonding
In 1916, Kössel and G.N. Lewis independently proposed a theory of chemical bonding based on the electronic structure of atoms. Their approach was among the earliest to explain valency in terms of electrons, particularly the inertness of noble gases.
Lewis Concept of Atom and Valence
Lewis envisioned an atom as composed of:
- A positively charged kernel (nucleus + inner electrons).
- An outer shell accommodating up to eight electrons, which are positioned at the corners of a cube.
- Atoms tend to achieve a stable octet (like noble gases) via bonding.
- In bonding:
- Sodium (Na) donates an electron → forms Na⁺.
- Chlorine (Cl) gains an electron → forms Cl⁻.
- The resulting NaCl molecule is held by electrostatic attraction (ionic bond).
- In molecules like Cl₂, H₂, etc., atoms share electrons to complete the octet.
Lewis Symbols
Only valence electrons participate in chemical bonding. Inner shell electrons are generally not involved. Lewis symbols: Simple notations where valence electrons are shown as dots around the element’s symbol.
Examples (Second Period Elements):
- Li: ·Li
- Be: ··Be
- B: ···B
- C: ····C
- N: ·····N
- O: ······O
- F: ·······F
- Ne: ········Ne
Significance:
- Number of dots = number of valence electrons.
- Helps determine:
- Group valence: Equal to the number of dots or 8 − number of dots.
Kössel’s Ionic Bond Formation
Kössel focused on the formation of ions and ionic compounds.
- Halogens (highly electronegative) and alkali metals (highly electropositive) are on opposite ends of the periodic table, separated by noble gases.
- Ionic bond forms by transfer of electrons:
- Halogens gain electrons → become anions.
- Alkali metals lose electrons → become cations.
- Resulting ions attain noble gas configuration (ns² np⁶).
- The electrostatic attraction between oppositely charged ions leads to bond formation.
Examples:
(a) Formation of NaCl:
Na → Na⁺ + e⁻ [Ne]3s¹ → [Ne]
Cl + e⁻ → Cl⁻ [Ne]3s² 3p⁵ → [Ne]3s² 3p⁶
Na⁺ + Cl⁻ → NaCl
(b) Formation of CaF₂:
Ca → Ca²⁺ + 2e⁻ [Ar]4s² → [Ar]
F + e⁻ → F⁻ [He]2s²2p⁵ → [Ne]
Ca²⁺ + 2F⁻ → CaF₂
- The electrovalent bond is the bond formed due to electrostatic attraction between cations and anions.
- Electrovalence = number of unit charges on the ion.
Octet Rule
Kössel and Lewis proposed the Octet Rule: Atoms tend to gain, lose, or share electrons in order to have eight electrons in their outermost shell (octet), resembling noble gases. Bond formation mechanisms under the octet rule:
- Transfer of electrons → Ionic bonds.
- Sharing of electrons → Covalent bonds.
Lewis Dot Structures
Represent bonding and lone pairs in molecules using dots for electrons. Main assumptions:
- Bonds form by the sharing of electron pairs.
- Each atom contributes at least one electron to the shared pair.
- Resulting atoms attain noble gas configurations.
Examples:
- H₂O: Two single covalent bonds; oxygen has two lone pairs.
- CCl₄: Four single bonds between C and Cl atoms.
- CO₂: Two double bonds between carbon and oxygen.
- N₂: Triple bond between nitrogen atoms.
Covalent Bonding
Introduced by Langmuir (1919) to explain electron sharing. A covalent bond is formed when atoms share one or more pairs of electrons. Each shared electron pair constitutes one covalent bond.
Drawing Lewis Structures
Steps for Writing Lewis Structures:
- Calculate the total valence electrons of all atoms in the molecule.
- Determine the central atom (least electronegative, usually written first).
- Form single bonds between the central atom and outer atoms.
- Complete octets of outer atoms (except H).
- Place remaining electrons on the central atom.
- Use multiple bonds if the central atom lacks an octet after the above steps.
Example: CH₄ (Methane)
- Carbon (4e⁻) + 4 Hydrogens (1e⁻ each) = 8 e⁻ total
- C is central; it forms 4 single bonds with H.
- All atoms attain a duplet/octet.
Formal Charge
The formal charge (FC) on an atom in a Lewis structure is the charge it would have if all atoms shared electrons equally. Helps identify the most stable and realistic Lewis structure when multiple are possible.
Example: O₃ (Ozone)
- Two resonance structures.
- Formal charges help determine which resonance forms are significant.
Exceptions to the Octet Rule
Despite its wide applicability, the octet rule has several exceptions:
(i) Incomplete Octet
(ii) Odd-Electron Molecules
(iii) Expanded Octet
Ionic Bond Formation
An ionic bond (also called an electrovalent bond) is formed by the complete transfer of electrons from one atom (usually a metal) to another (usually a non-metal). The atom losing electrons becomes a cation, and the one gaining electrons becomes an anion. The resulting electrostatic attraction between oppositely charged ions forms a stable ionic compound. The formation of an ionic compound depends on several energy factors:
Ionization Enthalpy (ΔiH)
The energy required to remove an electron from a neutral atom in the gaseous state. High ionization enthalpy makes cation formation difficult. Metals like alkali metals have low ionization enthalpies, thus form cations easily.
Electron Gain Enthalpy (ΔegH)
The energy is released when an atom gains an electron in the gaseous state. A highly negative value implies ease of anion formation. Non-metals like halogens have highly negative electron gain enthalpies.
Lattice Enthalpy (ΔlatticeH)
The energy is released when gaseous ions combine to form an ionic solid. A measure of the strength of the ionic bond
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Factors Favoring Ionic Bond Formation
The following are the factors favouring the formation of ionic bond.
- Low ionization enthalpy of metal.
- High electron gain enthalpy (more negative) of a non-metal.
- High lattice enthalpy of the resulting ionic compound.
Lattice Enthalpy
Lattice enthalpy (ΔlatticeH) is defined as the energy required to separate one mole of an ionic solid into its gaseous ions.
Example:
MX(s) → M⁺(g) + X⁻(g) ΔlatticeH > 0
Example:
M⁺(g) + X⁻(g) → MX(s) ΔlatticeH < 0
The numerical value is the same in both cases, but the sign differs based on the direction of the process.
Role in Ionic Bond Formation
Lattice enthalpy measures the electrostatic force between oppositely charged ions in an ionic crystal. It is a major contributor to the overall energy change during the formation of an ionic compound.
Bond Parameters
Bond parameters are measurable quantities that describe the nature and strength of a chemical bond. These include:
Bond Length
The average distance between the nuclei of two bonded atoms in a molecule. Determined experimentally using X-ray diffraction or spectroscopic methods.
Unit: Picometre (pm)
1 pm = 10⁻¹² m.
Depend on:
- Sizes of atoms involved
- Type of bond (single, double, triple)
Bond Angle
The angle between the two bonds formed by the central atom in a molecule.
Unit: Degrees (°).
Determined by:
- Number of lone pairs on the central atom
- Electronegativity of bonded atoms
- Hybridization of the central atom
Bond Enthalpy
The amount of energy required to break one mole of bonds of a particular type in gaseous molecules. Also called as bond dissociation enthalpy
Unit: kJ mol⁻¹
Bond Order
The number of covalent bonds between two atoms in a molecule.
Calculation:
Bond Order = (Number of bonding electron pairs) − (Number of antibonding electron pairs) / 2
Resonance
When a single Lewis structure cannot represent a molecule accurately, two or more structures (resonance structures) are used. The actual structure is a resonance hybrid, an average of all valid resonance structures.
Rules:
- All resonance structures must have the same arrangement of atoms.
- Differ only in the distribution of electrons.
- The actual molecule is more stable than any individual resonance structure.
Covalent Bonds: Polar and Nonpolar
Below, we have discussed the types of covalent bonds.
Nonpolar Covalent Bond
- A nonpolar bond is formed between two atoms of identical or similar electronegativity.
- The shared electron pair is equally distributed between the two atoms.
Example:
- H₂ molecule → Both atoms have the same electronegativity.
- The electron cloud is symmetrical.
- No separation of charge → bond is nonpolar.
Polar Covalent Bond
- A polar bond arises when the two atoms involved in a covalent bond have different electronegativities.
- The more electronegative atom attracts the shared electron pair more strongly.
- This leads to partial charges:
- δ⁻ on the more electronegative atom
- δ⁺ on the less electronegative atom
Example:
- HF molecule:
- F is more electronegative than H.
- F develops a partial negative charge (δ⁻), and H becomes partially positive (δ⁺).
- Bond becomes polar covalent.
Dipole Moment
The product of the magnitude of the charge (Q) and the distance (r) between the two charges.
Formula:
Dipole moment (μ) = Q × r
Unit: Debye (D)
1 D = 3.33564 × 10⁻³⁰ C·m
Vector Quantity:
Has both magnitude and direction (from δ⁺ to δ⁻).
Significance
Helps in:
- Distinguishing between polar and nonpolar molecules.
- Determining the symmetry and shape of molecules.
- Predicting solubility, boiling points, and chemical reactivity.
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Valence Shell Electron Pair Repulsion (VSEPR) Theory and Molecular Shapes
Developed by Sidgwick and Powell (1940), refined by Nyholm and Gillespie.
Purpose: To predict the shape of molecules based on repulsions between electron pairs in the valence shell of atoms.
Postulates of VSEPR Theory
The postulates of VSEPR theory are:
- The shape of a molecule depends on the number of valence shell electron pairs (bonded + lone pairs) around the central atom.
- Electron pairs (bond pairs and lone pairs) repel each other.
→ They orient themselves as far apart as possible to minimize repulsion. - Lone pair–lone pair repulsion > lone pair–bond pair repulsion > bond pair–bond pair repulsion.
- Multiple bonds (double/triple) are treated as one electron pair for geometry prediction.
- The geometry is determined by the total number of electron pairs, but the molecular shape is determined by the positions of the atoms only.
Types of Electron Pair Arrangements
Below is the table highlighting the types of electron pairs arrangements.
| Total Electron Pairs | Arrangement | Example | Shape of Molecule |
| 2 | Linear | BeCl₂ | Linear |
| 3 | Trigonal planar | BF₃ | Trigonal planar |
| 4 | Tetrahedral | CH₄ | Tetrahedral |
| 5 | Trigonal bipyramidal | PCl₅ | Trigonal bipyramidal |
| 6 | Octahedral | SF₆ | Octahedral |
Effect of Lone Pairs on Shape
Below is the table showcasing the effect of lone pairs on shape.
| Total Electron Pairs | Bond Pairs | Lone Pairs | Example | Geometry (Electron Pair) | Shape (Molecular) |
| 4 | 3 | 1 | NH₃ | Tetrahedral | Trigonal pyramidal |
| 4 | 2 | 2 | H₂O | Tetrahedral | Bent/V-shaped |
| 5 | 4 | 1 | SF₄ | Trigonal bipyramidal | See-saw |
| 5 | 3 | 2 | ClF₃ | Trigonal bipyramidal | T-shaped |
| 5 | 2 | 3 | XeF₂ | Trigonal bipyramidal | Linear |
| 6 | 5 | 1 | BrF₅ | Octahedral | Square pyramidal |
| 6 | 4 | 2 | XeF₄ | Octahedral | Square planar |
Bond Angle Deviations and Shape Explanation
Below is the table showcasing the bond angle deviations and shape explanations.
| Molecule | Expected Geometry | Bond Angle (°) | Observed Deviation | Reason |
| CH₄ | Tetrahedral | 109.5° | None | No lone pairs |
| NH₃ | Tetrahedral | 109.5° | ~107° | 1 lone pair → more repulsion |
| H₂O | Tetrahedral | 109.5° | ~104.5° | 2 lone pairs → more repulsion |
Introduction to Valence Bond Theory
Developed by Heitler and London (1927), extended by Linus Pauling.
A quantum mechanical approach to explain the formation and properties of covalent bonds.
Basic Principles of VBT
The basic principles of VBT are provided below.
- A covalent bond is formed when two atomic orbitals (from different atoms) overlap, and the electrons in them are paired with opposite spins.
- The strength of a covalent bond depends on the extent of overlap:
- The greater the overlap, → stronger the bond.
- Each atom retains its own identity, and the bond is localized between the two overlapping orbitals.
Orbital Overlap Concept
Bond formation occurs due to the overlap of half-filled atomic orbitals. The extent of overlap determines bond strength and bond length.
Types of Orbital Overlaps:
- s–s overlap: e.g., H–H bond in H₂
- s–p overlap: e.g., H–F molecule
- p–p overlap: e.g., Cl–Cl in Cl₂
Directional Properties of Bonds
Covalent bonds are directional in nature (unlike ionic bonds). The shape of the molecule depends on the direction in which orbitals overlap.
Types of Covalent Bonds Based on Overlap
Covalent bonds are classified into sigma (σ) and pi (π) bonds based on the type and extent of orbital overlap.
Sigma (σ) Bond
Formed by end-to-end (axial) overlap of atomic orbitals. Overlapping orbitals may be:
- s–s, s–p, or p–p (along the internuclear axis).
- Stronger than pi bonds due to greater overlap.
- Free rotation is possible around sigma bonds.
Pi (π) Bond
Formed by the sideways (lateral) overlap of two parallel p orbitals. Electron density lies above and below the internuclear axis. Weaker than sigma bonds due to less effective overlap. No free rotation due to the electron cloud locking the atoms in position.
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Important Formulas in NCERT Notes Class 11 Chemistry (Part-I) Chapter 4: Chemical Bonding & Molecular Structure
Here are the important formulas and equations from NCERT Class 11 Chemistry Chapter 2: Structure of Atom:
- Formal Charge (FC)
Formal Charge = (Valence electrons in free atom) – (Non-bonding electrons) – ½ × (Bonding electrons) - Dipole Moment (μ) = q × r
- Bond Order = ½ × (Number of bonding electrons – Number of antibonding electrons)
- ΔH (formation) = Sublimation energy + Ionization energy + ½ × Bond dissociation energy + Electron gain enthalpy + Lattice enthalpy
- Molecular Orbital Energy Order (for atomic number ≤ 7) σ(1s) < σ*(1s) < σ(2s) < σ*(2s) < π(2p_x) = π(2p_y) < σ(2p_z) < π*(2p_x) = π*(2p_y) < σ*(2p_z)
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FAQs
Ans: A sigma (σ) bond is formed by head-on (end-to-end) overlap of orbitals along the internuclear axis, while a pi (π) bond is formed by sidewise overlap of parallel p-orbitals. Sigma bonds are stronger than pi bonds due to a greater extent of overlapping.
Ans: Lattice enthalpy is the energy required to completely separate one mole of an ionic solid into its gaseous ions. It is important because it provides a measure of the stability of an ionic compound; the higher the lattice enthalpy, the more stable the compound.
Ans: Lone pairs are localized on the central atom and occupy more space than bonding pairs, which are shared between two atoms. This leads to greater repulsion by lone pairs, causing distortions in molecular shape and bond angles.
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