What Is a Hydrogen Bond
Nitrogen, oxygen and fluorine are the most electronegative elements in everyday chemistry. When one of them is joined to a hydrogen atom by a covalent bond, the shared electron pair is pulled strongly towards the more electronegative atom. The hydrogen, stripped of much of its electron density, is left with a sizeable fractional positive charge. That partially positive hydrogen is then attracted to the electronegative atom of a neighbouring molecule. The resulting link is the hydrogen bond, and it is weaker than the covalent bond that holds the same hydrogen in place.
NCERT defines it precisely: a hydrogen bond is the attractive force which binds the hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule. The hydrogen acts as a bridge between two atoms — held on one side by a true covalent bond and on the other by the weaker electrostatic hydrogen bond. By convention the covalent bond is drawn as a solid line and the hydrogen bond as a dotted line.
The hydrogen fluoride chain is the textbook illustration. Each $\ce{HF}$ molecule is internally covalent, yet the molecules string together through hydrogen bonds:
$\ce{\overset{\delta+}{H}-\overset{\delta-}{F} \bond{...} \overset{\delta+}{H}-\overset{\delta-}{F} \bond{...} \overset{\delta+}{H}-\overset{\delta-}{F}}$
Cause of Formation
The origin of the hydrogen bond is purely a matter of charge separation. When hydrogen is bonded to a strongly electronegative element $\ce{X}$ (where $\ce{X}$ is $\ce{F}$, $\ce{O}$ or $\ce{N}$), the bonding electron pair shifts far away from the hydrogen. As a result the hydrogen becomes highly electropositive relative to $\ce{X}$, acquiring a fractional positive charge $\delta+$, while $\ce{X}$ takes on a fractional negative charge $\delta-$. The bond $\ce{H-X}$ is therefore strongly polar.
This polarity sets up an electrostatic attraction between the exposed $\delta+$ hydrogen and the $\delta-$ atom of a neighbour, generalised by NCERT as:
$\ce{\overset{\delta+}{H}-\overset{\delta-}{X} \bond{...} \overset{\delta+}{H}-\overset{\delta-}{X} \bond{...} \overset{\delta+}{H}-\overset{\delta-}{X}}$
The magnitude of hydrogen bonding depends on the physical state of the compound: it is maximum in the solid state (molecules locked closest together) and minimum in the gaseous state (molecules far apart). Because of this dependence, hydrogen bonds exert a powerful influence on the structure and properties of the compound across all three states.
Conditions and Strength
Two requirements must be met for a hydrogen bond to form. First, the hydrogen must be covalently attached to a highly electronegative atom so that it carries a large $\delta+$. Second, the acceptor atom must itself be small and electronegative and must possess a lone pair to receive the bond. This is why only $\ce{F}$, $\ce{O}$ and $\ce{N}$ qualify — they are electronegative enough and compact enough to concentrate the negative charge effectively. Larger atoms such as chlorine, despite reasonable electronegativity, are too diffuse and do not form appreciable hydrogen bonds.
| Bond type | Approx. energy to break | Drawn as | Relative strength |
|---|---|---|---|
| Covalent bond | A few hundred kJ mol−1 | Solid line | Strongest |
| Hydrogen bond | About 4–25 kJ mol−1 | Dotted line | Weak, but directional |
| van der Waals forces | Very low | — | Weakest |
As the NIOS text records, the strength of a hydrogen bond is only about 4 to 25 kJ mol−1 — small compared with a covalent bond that needs hundreds of kJ mol−1 to break, yet noticeably stronger than ordinary van der Waals attraction. This intermediate position is the key to its importance: weak enough to break and re-form constantly, strong enough to hold molecules together and dictate boiling points, densities and the shapes of large biomolecules.
Types of Hydrogen Bonds
NCERT classifies hydrogen bonds into two types depending on whether the partner electronegative atom lies in a different molecule or in the same molecule.
| Feature | Intermolecular H-bond | Intramolecular H-bond |
|---|---|---|
| Formed between | Two different molecules (same or different compounds) | Two electronegative atoms within the same molecule |
| Effect on molecules | Links molecules into associated chains/networks | Forms a closed ring inside one molecule; no association |
| Effect on boiling point | Raises boiling point (extra energy needed to separate molecules) | Lowers boiling point (molecules stay independent and volatile) |
| Examples | $\ce{HF}$, water, alcohols | o-nitrophenol, salicylaldehyde |
Intermolecular hydrogen bond
This bond is formed between two different molecules of the same or different compounds. The hydrogen of one molecule reaches across to the electronegative atom of a neighbour, knitting many molecules together. The hydrogen fluoride chain of Figure 1, the association of alcohol molecules, and the network in liquid water are all intermolecular. Because energy is required to pull these associated molecules apart, intermolecular hydrogen bonding always raises melting and boiling points.
Intramolecular hydrogen bond
This bond forms when the hydrogen atom lies between two highly electronegative ($\ce{F}$, $\ce{O}$ or $\ce{N}$) atoms present within the same molecule. NCERT's standard example is o-nitrophenol, where the hydrogen sits between the two oxygen atoms of the adjacent hydroxyl ($\ce{-OH}$) and nitro ($\ce{-NO2}$) groups, closing a small ring inside the molecule. Salicylaldehyde behaves the same way. Because the hydrogen's bonding need is satisfied internally, the molecule does not link strongly to its neighbours.
A hydrogen bond is a polar covalent $\ce{H-X}$ bond plus an electrostatic pull. Revise how electronegativity sets up bond polarity in Bond Parameters.
o- vs p-Nitrophenol
The two isomers of nitrophenol are the most-tested illustration of how the type of hydrogen bond changes physical behaviour. In ortho-nitrophenol the nitro and hydroxyl groups sit on adjacent carbons, close enough for the $\ce{O-H}$ hydrogen to reach the oxygen of the nitro group and form an intramolecular hydrogen bond. The molecule essentially satisfies itself, so there is little association between separate molecules.
In para-nitrophenol the two groups sit on opposite sides of the ring, far too distant for an internal bond. Instead, the $\ce{O-H}$ of one molecule hydrogen-bonds to the nitro oxygen of a neighbouring molecule — an intermolecular bond. Many such bonds chain the molecules together.
Why is o-nitrophenol MORE volatile than p-nitrophenol?
Students often reason "ortho has hydrogen bonding, so it should boil higher" — and pick the wrong isomer. The trap is forgetting which type of bond is at work. In o-nitrophenol the bonding is intramolecular: it is locked inside one molecule and cannot tie molecules to one another, so the molecules separate easily and the compound is volatile (lower boiling point). In p-nitrophenol the bonding is intermolecular: it links many molecules into associated chains, so far more energy is needed to vaporise it (higher boiling point, lower volatility).
Intramolecular H-bond → no association → MORE volatile. Intermolecular H-bond → association → LESS volatile. Hence o-nitrophenol is more volatile than p-nitrophenol.
Water and the Ice Structure
Water is the supreme example of intermolecular hydrogen bonding. Each oxygen carries two lone pairs and each molecule has two $\ce{O-H}$ bonds, so a single water molecule can both donate two hydrogen bonds and accept two — up to four bonds in all. NIOS notes that it is "due to hydrogen bonding only that water exists as a liquid"; without it, a molecule as light as $\ce{H2O}$ would be a gas at room temperature.
When water freezes, these four bonds per molecule arrange themselves into a fixed, open, cage-like tetrahedral network. The molecules in this rigid lattice are held farther apart than in the jostling liquid, so ice occupies more volume than the same mass of water. The consequence is the famous low density of ice — it floats — which NIOS explicitly attributes to hydrogen bonding.
Effects on Physical Properties
The single most exam-relevant consequence of hydrogen bonding is its effect on boiling points. Molecules that hydrogen-bond intermolecularly need extra energy to be separated into the gas phase, so their boiling points are abnormally high for their molecular mass. This is precisely why $\ce{HF}$ boils higher than $\ce{HCl}$, and why $\ce{H2O}$ boils far above $\ce{H2S}$, despite being lighter — an outcome that contradicts the usual "heavier means higher boiling point" expectation.
Arrange $\ce{HF}$, $\ce{HCl}$, $\ce{HBr}$ and $\ce{HI}$ in order of boiling point and explain the position of $\ce{HF}$.
Among the heavier halides, boiling point rises with molecular mass through stronger van der Waals forces, giving $\ce{HCl} < \ce{HBr} < \ce{HI}$. But $\ce{HF}$ breaks the pattern: fluorine is the most electronegative and smallest halogen, so $\ce{HF}$ molecules associate through strong intermolecular hydrogen bonds. Breaking these bonds at the boiling point demands extra energy, lifting $\ce{HF}$ above $\ce{HCl}$ and even above $\ce{HBr}$. The observed order is therefore $\ce{HCl} < \ce{HBr} < \ce{HF} < \ce{HI}$, with $\ce{HF}$ standing out as the hydrogen-bonded anomaly. The same logic explains why $\ce{H2O}$ (hydrogen-bonded) boils much higher than $\ce{H2S}$ (not appreciably hydrogen-bonded).
Beyond boiling point, hydrogen bonding also raises melting points, increases viscosity and surface tension, and governs solubility — alcohols and ammonia dissolve readily in water because they can hydrogen-bond with it. Since the bond is maximum in the solid state and minimum in the gaseous state, its grip on properties is strongest in solids and liquids and essentially absent in vapour.
Hydrogen bond vs hydrogen bonding to chlorine
A common mistake is to assume any electronegative atom forms a hydrogen bond. NCERT restricts effective hydrogen bonding to $\ce{F}$, $\ce{O}$ and $\ce{N}$. Chlorine has a comparable electronegativity to nitrogen but is much larger, so its charge is too spread out to hold a hydrogen bridge. This is why $\ce{HCl}$ does not show the boiling-point anomaly that $\ce{HF}$ does.
Only F, O, N form significant hydrogen bonds — small size plus high electronegativity. Cl is electronegative but too large.
Biological Relevance
Hydrogen bonding is not merely an inorganic curiosity. NIOS notes that it "plays an important role in the structure and function of many biomolecules like proteins and nucleic acids." In proteins, hydrogen bonds between backbone groups hold helices and sheets in shape; in nucleic acids they pair the two strands of the double helix. Because each individual bond is weak (4–25 kJ mol−1), these structures can be assembled, read and taken apart by the cell without breaking strong covalent bonds — the very weakness of the hydrogen bond is what makes life's machinery reversible.
For NEET this connects Unit 4 chemistry directly to the biology and biomolecules syllabus, where the same hydrogen-bonded base pairing and protein folding reappear. Recognising the chemical basis here makes those later topics far easier to reason about.
Hydrogen Bonding in one screen
- A hydrogen bond binds the $\delta+$ hydrogen of one molecule to a highly electronegative atom — only $\ce{F}$, $\ce{O}$ or $\ce{N}$ — of another; it is weaker than a covalent bond and drawn dotted.
- Cause: $\ce{H-X}$ polarity leaves H electropositive; bonding is maximum in the solid state, minimum in the gas state.
- Strength is about 4–25 kJ mol−1 — stronger than van der Waals, far weaker than a covalent bond.
- Intermolecular H-bonds (e.g. $\ce{HF}$, water, alcohols) link molecules and raise boiling point; intramolecular H-bonds (e.g. o-nitrophenol) close an internal ring and lower it.
- o-Nitrophenol (intramolecular) is more volatile than p-nitrophenol (intermolecular).
- Ice has an open four-bond tetrahedral network, making it less dense than water so it floats; H-bonding also explains the high boiling points of $\ce{HF}$ and $\ce{H2O}$ and is central to proteins and nucleic acids.