Hydrogen Bonding & Intermolecular Forces

The Unsung Hero of Chemistry

Water boils at $100\,^{\circ}C$. But $H_2S$, $H_2Se$ and $H_2Te$ — heavier cousins from the same group — all boil below room temperature. By the trend of increasing molar mass, water 'should' boil around $-90\,^{\circ}C$. Something is holding water molecules together far more strongly than expected. That something is the hydrogen bond.

Definition: A hydrogen bond is an attractive force between a hydrogen atom covalently bonded to a small, highly electronegative atom (F, O, or N) and the lone pair of another electronegative atom (F, O, or N) nearby.

It is represented by a dotted line: $X-H \cdots Y$ (where X, Y = F, O, N).

Why it forms

When H is bonded to F, O, or N, the bond is highly polar — H carries a large $\delta^+$. This partially-bare proton is strongly attracted to the lone pair ($\delta^-$) on a neighbouring electronegative atom. It is essentially a strong dipole-dipole attraction.

Key Point: A hydrogen bond is much weaker than a covalent bond (about 5–40 kJ/mol vs 200–800 kJ/mol) but much stronger than ordinary van der Waals forces.

[NEET Important] Only F, O, and N are electronegative and small enough to form effective hydrogen bonds. Cl is electronegative but too large, so HCl shows only weak H-bonding.

Conditions & Types of Hydrogen Bonding

Conditions for hydrogen bonding

1. Hydrogen must be bonded to a highly electronegative atom (F, O, or N).
2. That electronegative atom should be small (high charge density).
3. A nearby atom must have a lone pair to accept the H.

Two types

Intermolecular hydrogen bonding — between different molecules. It links many molecules together, raising boiling point, viscosity and surface tension.

• Examples: water, $HF$, alcohols ($R-OH$), ammonia, carboxylic acids (which even dimerise).

Intramolecular hydrogen bonding — within the same molecule, when the donor and acceptor are positioned to form a ring (usually 5- or 6-membered).

• Examples: o-nitrophenol, salicylaldehyde, o-hydroxybenzaldehyde.
• This 'locks up' the H internally, so such molecules have lower boiling points than their isomers that form intermolecular bonds.

The o- vs p-nitrophenol classic

ortho-nitrophenol forms an intramolecular H-bond (lower b.p., more volatile, steam-volatile), while para-nitrophenol forms intermolecular H-bonds (higher b.p., less volatile). A perennial exam question.

[JEE Tip] Intramolecular H-bonding lowers boiling point (no inter-molecular network); intermolecular H-bonding raises it.

The Anomalous Properties of Water

Hydrogen bonding explains water's life-supporting peculiarities.

Ice floats on water (density anomaly)

In ice, each water molecule is hydrogen-bonded to four neighbours in an open, cage-like tetrahedral structure with large hollow spaces. This makes ice less dense than liquid water, so ice floats. When ice melts, the cage partly collapses and density increases, reaching a maximum at $4\,^{\circ}C$. This is why lakes freeze top-down, letting aquatic life survive winter.

High boiling point & high specific heat

Breaking the extensive H-bond network needs lots of energy → high boiling point and high heat capacity (water moderates climate).

High surface tension & viscosity

H-bonds pull surface molecules inward, giving water its high surface tension (insects walk on water) and notable viscosity.

Solvent power

Water's polarity and H-bonding let it dissolve ionic and polar substances exceptionally well.

Key Point: The open tetrahedral H-bonded lattice of ice is the direct reason ice is less dense than water — a fact that quite literally allows aquatic life to survive frozen winters.

[NEET Important] Maximum density of water is at $4\,^{\circ}C$; ice is less dense than liquid water due to its open H-bonded structure.

Hydrogen Bonding in Biology & A Note on van der Waals Forces

In living systems

Hydrogen bonds are everywhere in biology:

• DNA: the two strands of the double helix are held together by H-bonds between base pairs (A=T with two H-bonds, G≡C with three). These bonds are weak enough to 'unzip' for replication, yet strong enough to store information faithfully.
• Proteins: H-bonds stabilise the α-helix and β-pleated sheet secondary structures.

The broader family: van der Waals forces

Hydrogen bonds are the strongest of the intermolecular forces. The weaker van der Waals forces include:

• Dipole-dipole forces — between polar molecules (e.g. HCl).
• Dipole-induced dipole forces — a polar molecule distorts a non-polar one.
• London dispersion forces — instantaneous-induced-dipole attractions present in all molecules; they increase with molecular size/mass (why $I_2$ is a solid but $F_2$ a gas).

Key Point — strength order of intermolecular forces: ionic/covalent (intramolecular, strongest) ≫ hydrogen bond > dipole-dipole > London dispersion (weakest).

[JEE Tip] Boiling-point trends within a series are governed by H-bonding first, then molecular mass (dispersion forces).

Solved Examples

Example 1: Why does HF have a higher boiling point than HCl?

Explain using hydrogen bonding.

Solution:

1. F is small and highly electronegative, so $HF$ forms strong intermolecular H-bonds ($F-H \cdots F$).
2. Cl is larger and less electronegative, so $HCl$ forms only weak dipole-dipole forces.
3. Result: more energy is needed to separate $HF$ molecules → higher boiling point despite $HF$'s lower molar mass.

Takeaway: Hydrogen bonding in $HF$ overrides the molar-mass trend, giving it an anomalously high boiling point.

Example 2: Why ice floats on water

Explain why solid water (ice) is less dense than liquid water.

Solution:

1. In ice, each $H_2O$ is H-bonded to four others in an open, tetrahedral cage with empty spaces.
2. This open structure spreads molecules out → lower density.
3. On melting, the cage partially collapses, packing molecules closer → liquid water is denser.
4. Conclusion: ice (less dense) floats on water.

Takeaway: The open H-bonded lattice of ice makes it less dense than liquid water.

Example 3: o- vs p-nitrophenol

Which has the higher boiling point, ortho- or para-nitrophenol, and why?

Solution:

1. ortho-nitrophenol forms an intramolecular H-bond (a 6-membered ring), tying up its own H.
2. para-nitrophenol cannot do this; instead it forms intermolecular H-bonds linking many molecules.
3. Result: para has a stronger intermolecular network → higher boiling point.

Takeaway: Intramolecular H-bonding (ortho) lowers boiling point; intermolecular (para) raises it.

Example 4: Which molecule cannot form hydrogen bonds?

Among $H_2O$, $NH_3$, $CH_4$ and $HF$, which cannot form hydrogen bonds?

Solution:

1. H must be bonded to F, O, or N for H-bonding.
2. $CH_4$: hydrogen is bonded to carbon (not sufficiently electronegative) → no H-bonding.
3. The others have H on O, N, or F → they H-bond.

Takeaway: No H-bonding without H attached to F, O, or N — methane is the odd one out.

Example 5: DNA base pairing

How many hydrogen bonds hold the G≡C and A=T base pairs in DNA?

Solution:

1. Guanine-Cytosine (G–C): three hydrogen bonds.
2. Adenine-Thymine (A–T): two hydrogen bonds.
3. Consequence: G–C-rich DNA is more thermally stable (needs more energy to separate).

Takeaway: G≡C (3 H-bonds) is stronger than A=T (2 H-bonds) — H-bonding underpins DNA stability.

Example 6: Boiling-point order of hydrides of group 16

Why does $H_2O$ have a much higher boiling point than $H_2S$?

Solution:

1. $H_2O$ forms strong intermolecular hydrogen bonds (O is small and very electronegative).
2. $H_2S$ cannot form effective H-bonds (S is larger, less electronegative) — only weak van der Waals forces.
3. Result: $H_2O$ boils at $100\,^{\circ}C$, far above $H_2S$ ($-60\,^{\circ}C$).

Takeaway: Hydrogen bonding makes water's boiling point anomalously high within group 16.

Example 7: Carboxylic acid dimer

Why does acetic acid have a higher-than-expected molar mass in benzene?

Solution:

1. Carboxylic acids form cyclic dimers through two intermolecular H-bonds between the –COOH groups.
2. In a non-polar solvent (benzene), the dimer is stable, so the effective particle is $(CH_3COOH)_2$.
3. Result: the apparent molar mass is roughly double (about 120 instead of 60).

Takeaway: Hydrogen-bonded dimerisation doubles the apparent molar mass of carboxylic acids in non-polar solvents.

Example 8: Comparing strengths of forces

Arrange in increasing strength: London dispersion forces, hydrogen bonding, covalent bonding.

Solution:

1. London dispersion: weakest intermolecular force.
2. Hydrogen bond: stronger intermolecular force (5–40 kJ/mol).
3. Covalent bond: an intramolecular bond, far stronger (200–800 kJ/mol).
4. Order: London dispersion < hydrogen bond < covalent bond.

Takeaway: H-bonds sit between weak van der Waals forces and strong covalent bonds.

Example 9: Solubility of alcohols in water

Why is ethanol completely miscible with water but ethane is not?

Solution:

1. Ethanol ($C_2H_5OH$) has an –OH group that forms hydrogen bonds with water.
2. These H-bonds between ethanol and water make mixing favourable → miscible.
3. Ethane is non-polar, cannot H-bond, and is essentially insoluble.

Takeaway: H-bonding between solute and water explains the solubility of alcohols, sugars, and amines.

Example 10: Why does water have maximum density at 4 °C?

Explain the density behaviour of water near freezing.

Solution:

1. Below 4 °C, H-bonds begin building the open, low-density ice-like structure → density decreases.
2. At 4 °C, the molecules are most closely packed → maximum density.
3. Above 4 °C, normal thermal expansion lowers density.

Takeaway: Competition between H-bond cage formation and thermal contraction gives water its density maximum at 4 °C.

Example 11: Surface tension of water

How does hydrogen bonding give water its high surface tension?

Solution:

1. Surface molecules are pulled inward and sideways by H-bonds with neighbours, but have no molecules above.
2. The net inward pull minimises surface area, creating a 'skin'-like tension.
3. Result: high surface tension (e.g., water striders can walk on water).

Takeaway: Extensive H-bonding pulls surface molecules together, giving water its unusually high surface tension.

Example 12: London forces and halogen states

Why is $F_2$ a gas, $Br_2$ a liquid, and $I_2$ a solid at room temperature?

Solution:

1. None of these are polar, so the only intermolecular force is London dispersion.
2. Dispersion forces increase with molecular size/electron count: $F_2 < Cl_2 < Br_2 < I_2$.
3. Result: larger, more polarisable molecules have stronger forces → $I_2$ is a solid, $Br_2$ a liquid, $F_2$ a gas.

Takeaway: For non-polar molecules, physical state is set by London dispersion forces, which grow with molecular size.