Chemistry Bridge Course - Lecture 12 | CHEMCA JEE & NEET

CHEMCA

JEE & NEET Academy by Abhishek Sengar

Lecture 12 Molecular Bonding Target: Class 10 to 11 Transition (JEE/NEET)

Chemical Bonding Basics

Welcome to Lecture 12 of the CHEMCA Bridge Course! Chemical Bonding is the absolute core of both organic and inorganic chemistry. In this session, Abhishek Sengar Sir answers fundamental questions on why bonds are formed, how bond formation lowers energy to bring thermodynamic stability, and explores the three central bonding classes: Ionic, Covalent, and Coordinate bonds.

Video Lecture Broadcast

Instructor: Abhishek Sengar Sir Published: April 26, 2026 Subject: Chemical Bonding

Interactive Lecture Timestamps

Click any topic to skip the video directly to that specific concept explanation.

In-Depth Lecture Notes & Summary

What is a Chemical Bond?

A Chemical Bond is simply an attractive force that holds constituent particles (atoms, molecules, or ions) together in various chemical species.

Why is it formed? Atoms undergo bonding to achieve thermodynamic stability. When two atoms approach each other and form a bond, the net potential energy of the system decreases. Energy and stability are inversely related: $$\text{Stability} \propto \frac{1}{\text{Energy}}$$

The Three Main Classes of Bonds

While there are many minor attractions, main chemical bonds are categorized into three classes:

1. Ionic / Electrovalent Bond (Complete Electron Exchange)

Formed via complete transfer of one or more electrons from one atom (which becomes a cation) to another (which becomes an anion). The resulting electrostatic force of attraction binds them.

Example: Sodium Chloride ($NaCl$)

• $Na \, ([Ne]\,3s^1) \xrightarrow{-e^-} Na^+ \, (2, 8)$

• $Cl \, ([Ne]\,3s^2 3p^5) \xrightarrow{+e^-} Cl^- \, (2, 8, 8)$

2. Covalent Bond (Mutual Sharing)

Formed when participating atoms mutually share electrons to complete their respective valence shells (duplet for Hydrogen, octet for others).

Examples:

Hydrogen: $H-H$ (Single)

Oxygen: $O=O$ (Double)

Nitrogen: $N\equiv N$ (Triple)

3. Coordinate / Dative Bond (One-Sided Donation, Shared by Both)

A special covalent bond where the shared electron pair is contributed entirely by only one of the bonded atoms (known as the donor) but shared by both atoms (including the acceptor). It is represented by an arrow ($\to$) pointing from the donor to the acceptor.

Examples:

• Ammonium: $NH_3 + H^+ \to NH_4^+$

• Ozone: $O_2 + O \to O_3$

Bond Pairs (BP) vs. Lone Pairs (LP)

When analyzing Lewis dot structures, valence electrons are classified into:

Bond Pairs (BP): Electron pairs involved in the formation of chemical bonds.
Lone Pairs (LP): Pairs of valence shell electrons that are not involved in bonding.

Calculating Formal Charges ($FC$)

In polyatomic ions or complex molecules, the net charge belongs to the entire species. However, we can calculate a Formal Charge for each individual atom to see how electron densities are distributed:

$$FC = V - L - \frac{1}{2}B$$

$V$: Total valence electrons in the free, unbonded atom.
$L$: Total number of non-bonding lone-pair electrons.
$B$: Total number of bonding electrons.

Key Rule: The sum of the formal charges on all atoms in a molecule/ion must equal the total net charge of that species.

Formal Charge Visualizer

Select a molecule or polyatomic ion. Click on any atom in the structural model to see its live formal charge calculation!

Select Molecule

Atom Inspector: Select an atom

Click any atom in the diagram above to display its formal charge formula ($V - L - \frac{1}{2}B$).

Lecture 12 Concept Test

Validate your understanding of electron shares, coordinate dative bonds, and formal charges.

Question 1 of 5

Score: 0/0

Stuck on Formal Charges?

If you are struggling to compute lone-pair or bonding electron counts in complex molecular systems, drop a query to Abhishek Sengar Sir!

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Chemistry Bridge Course Lecture 12