Explain hydrogen bonding class 11 Chapter Chemical bonding

1. Basic Idea

A hydrogen bond forms when:

1. One molecule (or part of the same molecule) has a hydrogen directly attached to a highly electronegative atom—mainly $N$, $O$, or $F$.
2. Another molecule or another part of the same molecule has a lone-pair of electrons on $N$, $O$, or $F$.
3. The partly positive hydrogen ($\delta^+$) is attracted to the lone-pair bearing atom ($\delta^-$).

2. Notation

A hydrogen bond is written with a dotted line:
$X-H \;\cdots\; Y$
Here, $X$ is $N/O/F$ bearing positive $\delta^+$ and $Y$ is $N/O/F$ with lone pair.

3. Types

• Inter-molecular hydrogen bonding: Between two different molecules – e.g.

  • Water ($H_2O$): each molecule can form up to four $H$-bonds.
  • Hydrogen fluoride ($HF$) chain formation.

• Intra-molecular hydrogen bonding: Within the same molecule – e.g.

  • o-nitrophenol forms a five-member chelate ring, causing low boiling point because molecules do not link with others.

4. Strength Order

Covalent bond $>$ Ionic interaction $>$ Hydrogen bond $>$ Dipole-dipole $>$ van der Waals.
Typical hydrogen bond energy: $10$–$40\, \text{kJ mol}^{-1}$.

5. Consequences

• High boiling/melting points – water, HF, alcohols.
• Density anomaly of water – ice is less dense due to open $H$-bonded lattice.
• Viscosity & surface tension – glycerol, water.
• DNA double helix – A–T and G–C base pairing via $H$-bonds.

Logical chain to identify hydrogen bonding in problems

1. Check if $H$ is attached straight to $N$, $O$, or $F$.
2. Look for lone-pair on another $N/O/F$.
3. Decide whether bonding is intra- or inter-molecular.
4. Predict physical properties (bp, mp, solubility) accordingly.

What is Hydrogen Bonding?

Imagine water molecules like small magnets. Each magnet has a tiny positive side and a tiny negative side.

• The positive side is a hydrogen (H) atom that is a little bit short of electrons (so it feels positive).
• The negative side is an electronegative atom like oxygen (O), nitrogen (N) or fluorine (F) that is full of electrons (so it feels negative).

Because opposites attract, the positive hydrogen of one molecule sticks (weakly) to the negative oxygen/nitrogen/fluorine of a neighbour molecule.
This weak sticking force is called hydrogen bond.
It is stronger than ordinary van der Waals forces but weaker than a normal covalent bond.

So, hydrogen bonding is just a friendly handshake between molecules, helping them stay close to each other.

Worked Example: Why is the boiling point of water higher than that of hydrogen sulphide ($H_2S$)?

Step 1 – Identify electronegativity:

• $O$ has higher electronegativity ($3.5$) than $S$ ($2.5$).

Step 2 – Check for $H$ attached to $N/O/F$:

• $H_2O$ has $O–H$ bonds → eligible for hydrogen bonding.
• $H_2S$ has $S–H$ bonds → $S$ is not electronegative enough → no hydrogen bonding.

Step 3 – Predict interactions:

• Water molecules form extensive network of $H$-bonds (up to 4 per molecule).
• $H_2S$ molecules interact only by weak van der Waals forces.

Step 4 – Relate to boiling point:

• Extra energy (heat) is needed to break $H$-bonds in water.
• Less energy needed for $H_2S$.

Therefore:
$\text{bp}(H_2O) \gg \text{bp}(H_2S)$
Boiling point of water $= 100\,^\circ C$, hydrogen sulphide $\approx -60\,^\circ C$.

Answer: Water has higher boiling point due to intermolecular hydrogen bonding; $H_2S$ lacks such bonding.