Structures of molecules of two compounds are given below : (a) Which of the two compounds will have intermolec

NCERT Exemplar · Q38

Q. Structures of molecules of two compounds are given below :

Two benzene-ring structures labelled (I) and (II): in (I) the NO2 and OH groups sit on adjacent ring carbons, in (II) they sit on opposite ring carbons — Figure

(a) Which of the two compounds will have intermolecular hydrogen bonding and which compound is expected to show intramolecular hydrogen bonding.

(b) The melting point of a compound depends on, among other things, the extent of hydrogen bonding. On this basis explain which of the above two compounds will show higher melting point.

(c) Solubility of compounds in water depends on power to form hydrogen bonds with water. Which of the above compounds will form hydrogen bond with water easily and be more soluble in it.

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Concept understanding — Hydrogen Bonding

Hydrogen Bonding: From Intuition to Precision

Imagine you're holding two magnets. If you bring the north pole of one close to the south pole of another, they snap together. Now imagine a much weaker version of that — a tiny tug, not a full lock. That's the spirit of hydrogen bonding.

In chemistry, atoms in a molecule share electrons through covalent bonds. But electrons aren't shared equally in all cases. Some atoms are greedy — they pull the shared electrons closer to themselves. Oxygen, nitrogen, and fluorine are the biggest electron-hoarders. When one of these atoms bonds with hydrogen, the hydrogen ends up with a slight positive charge (because its electron has been pulled away), and the other atom gets a slight negative charge.

Now here's the key: that slightly positive hydrogen is attracted to any nearby slightly negative atom (like oxygen, nitrogen, or fluorine) on another molecule. This attraction is a hydrogen bond.

Note

A hydrogen bond is not a true chemical bond like a covalent or ionic bond. It's an intermolecular force — a strong dipole-dipole attraction — but weaker than covalent bonds (about 1/10th to 1/20th the strength).

The Precise Definition

A hydrogen bond is an attractive interaction between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and another electronegative atom (N, O, or F) that has a lone pair of electrons.

We can write it as:

\text{X—H} \cdots \text{Y}

where X and Y are N, O, or F. The dotted line ( $\cdots$ ) represents the hydrogen bond. X—H is the donor (the molecule that provides the hydrogen), and Y is the acceptor (the molecule that provides the lone pair).

Why Only N, O, and F?

Three things make these three elements special:

High electronegativity — They pull electrons hard, creating a large partial positive charge on hydrogen.
Small size — The lone pair on Y is compact, allowing the hydrogen to get very close. Closer distance means stronger attraction.
Lone pairs — They have unshared electron pairs that can act as the acceptor.

Chlorine is electronegative, but it's too large — the hydrogen can't get close enough for a strong bond. Carbon is not electronegative enough.

What Makes Hydrogen Bonding Special?

Unlike other dipole-dipole interactions, hydrogen bonds are directional and stronger. They're about 5–30 kJ/mol, compared to 0.5–2 kJ/mol for ordinary van der Waals forces. This strength has dramatic consequences.

Real-World Consequences

Water's high boiling point — Water ( $H_2O$ ) boils at $100^\circ C$ , while hydrogen sulfide ( $H_2S$ ) boils at $-60^\circ C$ . Both are similar molecules, but water forms hydrogen bonds; $H_2S$ does not (sulfur is not electronegative enough). Those bonds must be broken to boil water, requiring much more energy.

Ice floats — In liquid water, molecules jostle and form temporary hydrogen bonds. When water freezes, the molecules arrange into a hexagonal lattice held open by hydrogen bonds. This structure is less dense than liquid water — hence ice floats. Without hydrogen bonding, ice would sink, and lakes would freeze from the bottom up, killing aquatic life.

DNA double helix — The two strands of DNA are held together by hydrogen bonds between base pairs (adenine-thymine and guanine-cytosine). These bonds are strong enough to keep the strands together, but weak enough to be unzipped during replication. …

Why this formula?

Hydrogen Bonding: Why It Happens — The Reasoning, Not Just the Rule

Hydrogen bonding is not a full covalent bond — it's a special type of intermolecular attraction. To understand why it occurs, we must look at the electronic structure of the atoms involved.

1. The Core Requirement: A "Naked" Proton

A hydrogen bond forms when a hydrogen atom is covalently bonded to a highly electronegative atom (like F, O, or N). Why?

Electronegativity difference pulls the bonding electron pair strongly toward the electronegative atom.
The hydrogen atom is left with almost no electron cloud — it becomes a partially positive proton ( $\delta^+$ ).

Key idea: The hydrogen is now a small, dense positive charge — it can get very close to a lone pair on another electronegative atom.

2. The Electrostatic Attraction (The "Why")

The partially positive hydrogen ( $\delta^+$ ) is attracted to a lone pair of electrons on another electronegative atom (the acceptor).

This is electrostatic — Coulomb's law governs it:

F = \frac{1}{4\pi\varepsilon_0} \cdot \frac{q_1 q_2}{r^2}

$q_1$ = partial positive charge on H
$q_2$ = partial negative charge on lone pair
$r$ = distance between them

Because the hydrogen is so small, $r$ is very small → force is strong (stronger than van der Waals, weaker than covalent).

3. Why Only F, O, N?

Not all electronegative atoms work. The atom must have:

Property	Why it matters
High electronegativity	Pulls electron density away from H, creating $\delta^+$
Small atomic size	Allows close approach of the H to the lone pair
At least one lone pair	Provides the negative site for attraction

F, O, and N satisfy all three. Cl is electronegative but too large — the H cannot get close enough for a strong bond.

4. The "Formula" for Hydrogen Bond Strength

There is no single formula for hydrogen bond energy, but the strength depends on:

E_{\text{H-bond}} \propto \frac{\delta^+ \cdot \delta^-}{r^2}

Where:

$\delta^+$ = partial charge on H (depends on electronegativity of donor atom)
$\delta^-$ = partial charge on acceptor lone pair
$r$ = distance between H and acceptor atom

Typical strengths (for context):

Covalent bond: ~400 kJ/mol
Hydrogen bond: 10–40 kJ/mol
van der Waals: ~1–5 kJ/mol

5. Directionality — The "Linear" Preference

Hydrogen bonds are directional: the strongest interaction occurs when the donor H–X bond and the acceptor lone pair are collinear (180° angle).

Why? Because:

The positive charge on H is concentrated along the bond axis …

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Ortho-nitrophenol exhibits intramolecular hydrogen bonding, while para-nitrophenol forms intermolecular hydrogen bonds. Due to stronger intermolecular forces, para-nitrophenol has a higher melting point and is more soluble in water than ortho-nitrophenol.

Understanding the nature of hydrogen bonding is key to explaining the physical properties of these compounds. Hydrogen bonding occurs when a hydrogen atom bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine) is attracted to another highly electronegative atom in the same or a different molecule. The position of the nitro group relative to the hydroxyl group significantly influences the type of hydrogen bonding possible.

Understanding Hydrogen Bonding in Nitrophenols
- Compound (I): Ortho-nitrophenol (2-nitrophenol)
  
  In ortho-nitrophenol, the hydroxyl ( $-\text{OH}$ ) group and the nitro ( $-\text{NO}_2$ ) group are on adjacent carbon atoms of the benzene ring. This close proximity allows the hydrogen atom of the hydroxyl group to form a hydrogen bond with an oxygen atom of the nitro group within the same molecule. This type of hydrogen bonding is called intramolecular hydrogen bonding.
  
  The structure can be visualized as:

\text{C}_6\text{H}_4(\text{OH})(\text{NO}_2) \quad \text{(ortho-isomer)}

    where the H of -OH forms a bond with an O of -NO$_2$ on the adjacent carbon.

*   **Compound (II): Para-nitrophenol (4-nitrophenol)**
    In para-nitrophenol, the hydroxyl ($-\text{OH}$) group and the nitro ($-\text{NO}_2$) group are on opposite carbon atoms of the benzene ring. The distance between these groups is too large for hydrogen bonding to occur within the same molecule. Instead, the hydrogen atom of the hydroxyl group of one molecule forms a hydrogen bond with an oxygen atom of the nitro group (or hydroxyl group) of an *adjacent molecule*. This type of hydrogen bonding is called **intermolecular hydrogen bonding**.

    The structure can be visualized as:

\text{C}_6\text{H}_4(\text{OH})(\text{NO}_2) \quad \text{(para-isomer)}

    where the H of -OH on one molecule forms a bond with an O of -NO$_2$ on another molecule.

> [!IMPORTANT]
> Intramolecular hydrogen bonding occurs *within* a single molecule, forming a ring-like structure. Intermolecular hydrogen bonding occurs *between* different molecules, leading to molecular association.

2. Explaining Melting Point Differences

The melting point of a compound is directly related to the energy required to overcome the intermolecular forces holding the molecules together in the solid state.

*   **Ortho-nitrophenol (Intramolecular H-bonding)**
    Since ortho-nitrophenol forms intramolecular hydrogen bonds, its hydrogen bonding capacity is satisfied *within* each molecule. This reduces the availability of the hydroxyl hydrogen and nitro oxygen atoms to form strong hydrogen bonds with *other* ortho-nitrophenol molecules. Consequently, the overall intermolecular forces (van der Waals forces, dipole-dipole interactions, and weaker intermolecular H-bonds) between ortho-nitrophenol molecules are relatively weaker. Less energy is required to separate these molecules, leading to a lower melting point.

*   **Para-nitrophenol (Intermolecular H-bonding)**
    Para-nitrophenol forms strong intermolecular hydrogen bonds. This means that individual para-nitrophenol molecules associate with each other, forming larger, more stable aggregates or chains in the solid state. These extensive intermolecular hydrogen bonds create a strong network of attractive forces between molecules. A significant amount of energy is needed to break these strong intermolecular attractions to transition from solid to liquid, resulting in a higher melting point.

Therefore, para-nitrophenol, with its stronger intermolecular hydrogen bonding, will have a higher melting point than ortho-nitrophenol. …

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