Chemistry Form 5: Chemical Bonding and Molecular Structure Notes

Chemical Bonding and Molecular Structure

What is a Chemical Bond?

A chemical bond is the attractive force that holds atoms together in a molecule. It’s the “glue” that keeps everything stable.

Main Types of Chemical Bonds

Ionic (or Electrovalent) bond
Covalent bond
Coordinate bond
Polar covalent bond
Non-polar covalent bond
Hydrogen bonding

1. Ionic (Electrovalent) Bond

An ionic bond forms when one or more electrons are completely transferred from a metal atom to a non-metal atom.

What happens during this process?

The metal atom loses electron(s) and becomes a positively charged ion (cation).
The non-metal atom gains those electron(s) and becomes a negatively charged ion (anion).
Both atoms achieve stable electron configurations like the noble gases.
The opposite charges of the cation and anion attract each other strongly (Coulombic force), forming the ionic bond.

Example: Sodium + Chlorine → Sodium chloride Na + Cl → Na⁺Cl⁻

This bond can also be described as the electrostatic force of attraction between oppositely charged ions. Energy is released when the bond forms.

Electrovalency

Electrovalency is the number of electrons an atom loses or gains to form an ion.

Metals usually show positive electrovalency (lose electrons).
Non-metals show negative electrovalency (gain electrons).

Note: Some elements like iron show variable electrovalency — Fe²⁺ in ferrous compounds and Fe³⁺ in ferric compounds.

When electrons are lost, the atom is oxidized. When electrons are gained, the atom is reduced.

Properties of Ionic Bonds

Formed by strong electrostatic attraction between oppositely charged ions.
Non-directional (the force depends on distance, not direction).
Break easily when the compound is dissolved in water (a polar solvent) or melted.
Typical example: NaCl (table salt).

Factors Influencing Ionic Bond Formation

Three main energy steps are involved:

Ionization energy — Energy absorbed to remove electron(s) from the metal.
Electron affinity — Energy released when the non-metal gains electron(s).
Lattice energy — Large amount of energy released when cations and anions come together to form the solid crystal.

2. Covalent Bond

A covalent bond is formed when two atoms (same or different) share electrons mutually so both can achieve stable noble gas configurations.

Definition: It is the force of attraction that results from the sharing of one or more pairs of electrons between two atoms.

Covalency

Covalency is the number of electrons an atom shares in forming the bond. Example: In H₂, each hydrogen atom has a covalency of 1.

Characteristics of Covalent Bonds

Formed by mutual sharing of electrons.
Directional in nature — the shared electrons stay localized in a specific region between the two nuclei, giving the bond a definite direction.

Types of Covalent Bonds

Single Covalent Bond One pair of electrons is shared. Represented by a single line (–). Example: H–H (hydrogen molecule)

Multiple Covalent Bonds

Double Bond: Two pairs of electrons are shared (represented by =). Examples: O=O (oxygen), C=O (carbon dioxide)
Triple Bond: Three pairs of electrons are shared (represented by ≡). Examples: N≡N (nitrogen), H–C≡C–H (acetylene)

Examples of Molecule Formation

Oxygen (O₂) Molecule Each oxygen atom has 6 valence electrons and needs 2 more for stability. They share two pairs of electrons, forming a double bond (O=O).

Carbon Dioxide (CO₂) Carbon has 4 valence electrons and needs 4 more. Each oxygen needs 2 more. They form O=C=O, with two double bonds.

Nitrogen (N₂) Molecule Each nitrogen atom has 5 valence electrons and needs 3 more. They share three pairs of electrons, forming a triple bond (N≡N).

Comparison Between Single, Double, and Triple Covalent Bonds

Bond Length: Triple bond < Double bond < Single bond (Triple bonds are the shortest, single bonds are the longest)
Bond Strength / Bond Energy: Triple bond > Double bond > Single bond

Because shorter bonds are stronger, more energy is required to break a triple bond than a double bond, and more for a double bond than a single bond.

Factors Favouring the Formation of Covalent Bonds

Covalent bonds are more likely to form under these conditions:

High Ionization Energy: The atom does not lose electrons easily.
Similar Electron Affinity: The two atoms have nearly equal tendency to gain electrons, so they prefer to share rather than transfer.
Similar Electronegativity: When electronegativities are almost the same, electron transfer doesn’t occur; sharing takes place instead.
High Nuclear Charge and Small Atomic Size: Small atoms with high nuclear charge hold their electrons tightly, making complete transfer difficult.

Coordinate (Dative) Bond

A coordinate bond (also called a dative bond) is a type of covalent bond in which both shared electrons come from only one of the combining atoms.

The atom that donates the pair of electrons is called the donor.
The atom that accepts the pair is called the acceptor.

It is represented by an arrow (→) pointing from the donor to the acceptor.

Examples:

Formation of Ammonium ion (NH₄⁺) Nitrogen (in NH₃) donates its lone pair to H⁺. Nitrogen is the donor, H⁺ is the acceptor.
Ammonia-Boron Trifluoride Complex (NH₃ → BF₃) Nitrogen donates its lone pair to the boron atom.

Polarity in Covalent Bonds

Covalent bonds can be of two types depending on the atoms involved:

1. Non-Polar Covalent Bond

Forms between two identical atoms.
Electrons are shared equally.
The molecule is electrically neutral/symmetrical.
Example: H₂, Cl₂, O₂, N₂.

2. Polar Covalent Bond

Forms between two different atoms.
The shared electrons are attracted more towards the atom with higher electronegativity.
This creates a slight negative charge (δ⁻) on the more electronegative atom and a slight positive charge (δ⁺) on the less electronegative atom.
Examples: HCl, H₂O, HF.

Cause of Polarity: Difference in electronegativity. Example: H–F (1.9) > H–Cl (0.9) > H–Br (0.7) > H–I (0.4)

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Proposed by R.J. Gillespie and R.S. Nyholm in 1957.

Main Idea: The shape of a molecule is determined by the repulsion between electron pairs (bond pairs and lone pairs) around the central atom. Electron pairs arrange themselves as far apart as possible to minimize repulsion.

Order of Repulsion Strength: Lone pair – Lone pair > Lone pair – Bond pair > Bond pair – Bond pair

Shapes According to VSEPR Theory

Number of Bond Pairs	Shape	Bond Angle	Example
2	Linear	180°	BeCl₂
3	Trigonal Planar	120°	BF₃, BCl₃
4	Tetrahedral	109.5°	CH₄, SiCl₄
5	Trigonal Bipyramidal	90° & 120°	PCl₅
6	Octahedral	90°	SF₆

Shapes with Lone Pairs

3 Bond Pairs + 1 Lone Pair → Trigonal Pyramidal (e.g., NH₃, PCl₃)
2 Bond Pairs + 2 Lone Pairs → Bent (V-shaped) (e.g., H₂O, H₂S)
4 Bond Pairs + 2 Lone Pairs → Square Planar

Exercise Answers (VSEPR):

NH₃ → Trigonal Pyramidal
BeCl₂ → Linear
H₂O → Bent (V-shaped)
SiCl₄ → Tetrahedral

Hybridization

Definition: Hybridization is the mixing of atomic orbitals to form new hybrid orbitals that have different shapes and energies.

Key Points:

The number of hybrid orbitals formed equals the number of atomic orbitals mixed.
All hybrid orbitals of the same type are identical in shape and energy.
Hybrid orbitals are symmetrically arranged in space.

Types of Hybridization

sp³ Hybridization (Tetrahedral)
- One s + three p orbitals → four sp³ hybrid orbitals
- Bond angle: 109.5°
- Example: CH₄ (Methane)
sp² Hybridization (Trigonal Planar)
- One s + two p orbitals → three sp² hybrid orbitals
- Bond angle: 120°
- One unhybridized p orbital remains for π bond.
- Example: C₂H₄ (Ethene)
sp Hybridization (Linear)
- One s + one p orbital → two sp hybrid orbitals
- Bond angle: 180°
- Example: C₂H₂ (Acetylene)

Hydrogen Bonding

When a hydrogen atom is covalently bonded to a highly electronegative atom (F, O, or N), it develops a partial positive charge (δ⁺). This hydrogen can then form a weak electrostatic attraction with another electronegative atom (F, O, N) of a nearby molecule. This attraction is called a hydrogen bond.

Conditions for Hydrogen Bonding:

Hydrogen must be attached to a small, highly electronegative atom (F, O, N).
The electronegative atom should be small in size.

Examples:

Hydrogen fluoride (HF)
Water (H₂O)
Ammonia (NH₃)

Types of Hydrogen Bonding

Intermolecular Hydrogen Bonding Occurs between different molecules. Examples: Water, HF, NH₃ (this is why water has high boiling point).
Intramolecular Hydrogen Bonding Occurs within the same molecule (between two groups in one molecule).

Intramolecular Hydrogen Bonding

This type of hydrogen bonding occurs within the same molecule — between the hydrogen atom of one functional group and an electronegative atom (O, N, or F) of another functional group in the same molecule.

Example:

o-nitrophenol shows intramolecular hydrogen bonding.
p-nitrophenol shows intermolecular hydrogen bonding.

Effects of Hydrogen Bonding

Hydrogen bonding significantly influences the physical properties of compounds. Major effects include:

Molecular Association Molecules form aggregates or clusters due to hydrogen bonding. Example: Water molecules associate into groups, which explains many of its unusual properties.
Increase in Melting and Boiling Points Compounds with hydrogen bonding have higher melting and boiling points than expected. Examples: H₂O, NH₃, and HF have unusually high boiling points compared to similar compounds like H₂S, PH₃, and HCl.
Influence on Physical State Hydrogen bonding can determine whether a substance is liquid or gas at room temperature. Example: Water (H₂O) is a liquid, while hydrogen sulfide (H₂S) is a gas at room temperature.
Solubility in Water Covalent compounds that can form hydrogen bonds with water dissolve readily in it. Examples: Ethanol, ammonia, lower aldehydes, and ketones are soluble in water due to hydrogen bonding.

Review Questions (Atomic Structure & Related Topics)

Here’s a clean and organized version of the review questions:

Question 31

Using the electronic configuration and the periodic table, identify the element and the number of valence electrons.

(i) 1s² 2s² 2p⁴ (ii) 1s² 2s² 2p⁶ 3s² 3p³

Solution: (a) (i) Oxygen – 6 valence electrons (2s and 2p subshells). (ii) Phosphorus – 5 valence electrons (3s and 3p subshells).

Write the electronic configuration and orbital diagram for:

Cobalt
Nickel
Zinc (Using Hund’s rule)

Question 32

Provide the number of orbitals in each of the following: 1s, 2p, 3d, 4f.
How many sub-shells are there in shells K, L, and N?
Is it possible for an electron to have these quantum numbers: n=3, ℓ=3, m=3, s=+1/2? Explain.
An electron is in a 4f orbital. What possible values can its quantum numbers have?
What subshells are found in the shell with n=4?
Write the possible values of the four quantum numbers for the outermost 2 electrons of Calcium.
Write the four quantum numbers for an electron added to a Cl atom when forming Cl⁻.
Write the four quantum numbers for the outermost 2 electrons in Na.

Question 33

Write the total number of electrons possible for an atom with n=3 (assuming all orbitals are completely filled).
Write the electronic configuration of the following: O, K, Ni (Z=28), Cu⁺ (Z=29), Mo (Z=42), Cl⁻ (Z=17).
What is the number of unpaired electrons in the ground state of: Fe³⁺ (Z=26), Cr³⁺ (Z=24), Ni²⁺ (Z=28)?
Why is the electronic configuration 1s² 2s² 2pₓ² 2pᵧ⁰ 2p_z⁰ not correct for the ground state of nitrogen?
Mention the law disobeyed in the given arrangements.

Question 1 – Bohr’s Atomic Model

State two major postulates and five shortcomings of Bohr’s atomic model.

An electromagnetic radiation was emitted in the Balmer series from an electron transition between n=2 and n=5. Calculate:

Energy of the radiation (kJ/mol)
Frequency (Hz)
Wavelength (metres)

(Additional calculations on Balmer series transitions, energy, wavelength, etc., are given in the original notes.)

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