Why does ethanol have a higher boiling point than dimethyl ether?

Ethanol has an O-H group forming strong intermolecular hydrogen bonds. Dimethyl ether has all H atoms bonded to carbon and cannot hydrogen-bond. So ethanol needs much more energy to vaporise, giving a much higher boiling point.

Which of CH3COOH, NH3, HF, and CH4 shows maximum hydrogen bonding?

HF shows the strongest hydrogen bonding because fluorine has the highest electronegativity combined with a small size, giving maximum charge separation.

Hydrogen Bonding

Q: Why doesn't chlorine form hydrogen bonds despite being electronegative?

Chlorine has electronegativity comparable to nitrogen but its atomic size is much larger. The negative charge is spread over a bigger volume with low charge density, making the interaction with a delta-plus hydrogen too weak to count as a proper hydrogen bond.

Q: Which is the strongest hydrogen bond among F-H, O-H, and N-H?

F-H-F is strongest, followed by O-H-O, then N-H-N, following the order of electronegativity F greater than O greater than N.

Q: Why is ice less dense than liquid water?

In ice, each water molecule forms four hydrogen bonds in a tetrahedral cage-like open structure with lots of empty space. When ice melts, some H-bonds break and molecules pack closer, so liquid water is denser than ice.

Q: What is the role of hydrogen bonding in DNA?

Hydrogen bonds hold the two DNA strands together through base pairing. Adenine pairs with thymine via two H-bonds and guanine pairs with cytosine via three H-bonds, providing stability while still allowing strand separation for replication.

Q: How do van der Waals forces differ from hydrogen bonds?

Van der Waals forces are weaker (0.05 to 25 kJ/mol), non-directional attractions from temporary or induced dipoles present between all molecules. Hydrogen bonds are stronger (8 to 42 kJ/mol), directional, and require H bonded to F, O, or N.

ChemistryChemical Bonding And Molecular StructureFor JEE aspirants

A hydrogen bond is a weak attractive force between a hydrogen atom already covalently bonded to a highly electronegative atom (F, O, or N) and another electronegative atom nearby that carries a lone pair. Written as $X - H \dots Y$ , the hydrogen bond is stronger than ordinary van der Waals forces but much weaker than a covalent bond, with an energy of roughly $8 - 42 kJ mo l^{- 1}$ compared to $200 - 400 kJ mo l^{- 1}$ for covalent bonds. Hydrogen bonding explains why water is a liquid at room temperature while $H_{2} S$ is a gas, why ice floats, why ethanol boils higher than dimethyl ether of the same molecular mass, and why DNA holds its double helix together. This concept covers the conditions for H-bonding, intermolecular and intramolecular types, effects on physical properties, and a brief look at the wider family of intermolecular forces.

Key Ideas - Quick Reference

Hydrogen bond form: $X - H \dots Y$ , where X and Y are highly electronegative (F, O, N).
Bond energy: $8 - 42 kJ mo l^{- 1}$ (much weaker than covalent, stronger than van der Waals).
Order of H-bond strength: $F - H \dots F > O - H \dots O > N - H \dots N$ .
Two types: intermolecular (between separate molecules) and intramolecular (within one molecule).
Intermolecular H-bonding raises boiling point, melting point, viscosity, and surface tension.
Intramolecular H-bonding usually lowers boiling point (no molecular association) and increases volatility.

1. What is a Hydrogen Bond?

A hydrogen bond is the attractive force that binds a hydrogen atom of one molecule (covalently attached to a strongly electronegative atom X) to an electronegative atom Y of the same or a different molecule.

$X - H \dots Y$ where X, Y $=$ F, O, or N

A hydrogen bond is not a true covalent bond. It is essentially a strong dipole-dipole electrostatic attraction, with a small contribution from partial orbital overlap. In notation, the covalent X-H bond is shown as a solid line and the hydrogen bond as a dotted line ( $X - H \dots Y$ ).

Typical hydrogen bond strengths:

$F - H \dots F$ : the strongest known, $\sim 40 kJ mo l^{- 1}$ .
$O - H \dots O$ : $\sim 20 - 25 kJ mo l^{- 1}$ (as in water, alcohols, carboxylic acids).
$N - H \dots N$ : $\sim 8 - 15 kJ mo l^{- 1}$ (as in ammonia, amines, DNA base pairing).

2. Cause of Hydrogen Bond Formation

When hydrogen is covalently bonded to a strongly electronegative atom like F, O, or N, the shared electron pair is pulled strongly towards the electronegative atom. Since hydrogen has no inner electron shell, its nucleus becomes very exposed and acquires a substantial fractional positive charge ( $δ^{+}$ ). The electronegative atom acquires a matching fractional negative charge ( $δ^{-}$ ).

This highly polar $X - H$ bond means the tiny, exposed, positively-charged H nucleus can interact strongly with a lone pair on a nearby electronegative atom, forming the hydrogen bond.

$H^{δ +} - X^{δ -} \dots H^{δ +} - X^{δ -} \dots H^{δ +} - X^{δ -}$

Conditions for hydrogen bonding

Hydrogen must be covalently bonded to a highly electronegative element (F, O, or N only).
The electronegative element must have a small atomic size (so the negative charge is concentrated).
The electronegative element must have at least one lone pair to accept the interaction.

Chlorine, though about as electronegative as nitrogen, does not usually form hydrogen bonds because its atomic size is too large - the negative charge is diffuse, and the electrostatic interaction is weakened.

3. Types of Hydrogen Bonding

Intermolecular hydrogen bonding

Formed between two different molecules of the same or different substances. This is what gives water, alcohols, ammonia, carboxylic acids, and HF their characteristic properties. In liquid water, each molecule is bonded to about four neighbours on average; in crystalline HF, the molecules form long zig-zag chains.

Figure 1: Intermolecular hydrogen bonding in water shown as a 3D ball-and-stick model. Each water molecule (orange oxygen with two hydrogens) is linked to its neighbours by hydrogen bonds (dashed orange lines) — the

δ^{+}

H of one molecule is attracted to a lone pair on the

δ^{-}

O of another. Each O can donate two H-bonds and accept two, giving a nearly tetrahedral local network.

Figure 2: Solid hydrogen fluoride forms a zig-zag polymer chain

(HF)_{n}

. Solid lines show covalent H-F bonds; dashed lines show hydrogen bonds.

Intramolecular hydrogen bonding

Formed within a single molecule when two functional groups are positioned such that a hydrogen and an electronegative atom on the same molecule can come close enough. This creates a closed ring structure, called chelation.

Figure 3: Intramolecular hydrogen bonding, shown as a dashed orange line. Left: o-nitrophenol — the hydroxyl H forms a hydrogen bond to an oxygen of the ortho nitro group, closing a six-membered chelate ring. Right: salicylaldehyde — the hydroxyl H forms a hydrogen bond to the carbonyl O of the ortho aldehyde group. This locking-in of the H raises melting/boiling points less than intermolecular H-bonding would, so the ortho isomers are more volatile than the meta/para counterparts.

Key contrast: Intermolecular hydrogen bonding leads to molecular association (many molecules linked together) and raises the boiling point. Intramolecular hydrogen bonding traps the H atom inside one molecule so it can no longer link to others, which prevents association and lowers the boiling point.

4. Effects of Hydrogen Bonding

Boiling point and physical state

Solved Example 1

Water is a liquid at room temperature while

H_{2} S

is a gas, even though both belong to Group 16. Explain.

Solution:

Oxygen has both high electronegativity and small size, so $H_{2} O$ forms strong intermolecular hydrogen bonds. These bonds link many water molecules together, effectively increasing the mass that must be evaporated as a unit. This raises the boiling point of water to 100°C. Sulphur is much larger and less electronegative than oxygen, so $H_{2} S$ cannot form hydrogen bonds. Its molecules interact only through weak van der Waals forces, so it boils at $- 60$ °C and exists as a gas at room temperature.

Solved Example 2

Ethanol (

C_{2} H_{5} OH

) has a much higher boiling point (78°C) than dimethyl ether (

C H_{3} - O - C H_{3}

- 24

°C), even though both have the same molecular formula (

C_{2} H_{6} O

). Why?

Solution:

In ethanol, the hydrogen atom of the O-H group is bonded to O (highly electronegative and small) and can form strong intermolecular hydrogen bonds with the oxygen of another ethanol molecule. In dimethyl ether, all hydrogens are bonded to carbon, which is not electronegative enough to make the H atom $δ^{+}$ . So dimethyl ether cannot form hydrogen bonds; its molecules interact only through weak van der Waals forces. Ethanol therefore needs much more energy to evaporate.

Solubility

Substances that can form hydrogen bonds with water dissolve readily in it. Alcohols, ammonia, sugars, and low-molecular-mass carboxylic acids are all water-soluble because their O-H, N-H, or C=O groups can hydrogen-bond with water molecules. Hydrocarbons cannot form H-bonds and are insoluble.

Density anomaly of ice

Ice is less dense than liquid water - a rare property. In ice, each water molecule forms four hydrogen bonds arranged tetrahedrally, creating an open cage-like structure. When ice melts, some hydrogen bonds break, the cage partially collapses, and molecules pack closer together, so liquid water is denser than ice. This is why ice floats and why lakes freeze from the top down (protecting aquatic life).

Volatility of ortho vs para/meta isomers

Solved Example 3

Why is o-hydroxybenzaldehyde (salicylaldehyde) a liquid at room temperature while p-hydroxybenzaldehyde is a high-melting solid?

Solution:

In the ortho isomer, the OH and CHO groups are close enough on the ring to form an intramolecular hydrogen bond, locking the H inside a single molecule. No intermolecular association can happen, so the compound is a low-melting liquid.

In the para isomer, the OH and CHO groups are on opposite ends of the ring, so intramolecular H-bonding is geometrically impossible. Instead, extensive intermolecular H-bonding links many molecules together into a rigid network, making it a high-melting solid.

Biological importance

Hydrogen bonds are central to biology:

Proteins: The $N - H \dots O = C$ hydrogen bonds along the backbone stabilise the alpha-helix and beta-sheet secondary structures.
DNA: The two strands of the double helix are held together by hydrogen bonds between complementary base pairs - two H-bonds between adenine and thymine, three between guanine and cytosine. These bonds are strong enough to hold the helix together but weak enough to allow the strands to separate for replication.
Water in living systems: The high specific heat capacity, high heat of vaporisation, and high surface tension of water - all consequences of hydrogen bonding - make it uniquely suited as the medium of life.

5. Van der Waals and Other Intermolecular Forces

Hydrogen bonding is a special case of a broader family of intermolecular forces - the weak attractions between molecules that hold matter together in the liquid and solid states. These are collectively called van der Waals forces (or London dispersion forces after Fritz London who explained them using wave mechanics).

Types of intermolecular forces

Dipole-dipole interaction: Between molecules with permanent dipole moments, e.g., HCl, $H_{2} O$ . Positive end of one molecule attracts the negative end of another.
Ion-dipole interaction: Between an ion and a polar molecule, e.g., $N a^{+}$ in water (each $N a^{+}$ is surrounded by water molecules with their O atoms pointing towards it). Responsible for solvation.
Induced dipole interaction (Debye force): A permanent dipole (or ion) polarises a nearby non-polar molecule, inducing a dipole in it. Example: $HCl$ with $Ar$ .
Instantaneous dipole - induced dipole (London dispersion force): Even non-polar molecules have transient dipoles from random electron motion. These transient dipoles induce dipoles in neighbours, giving weak attractive forces. This is why noble gases can be liquefied and why $I_{2}$ is a solid at room temperature.
Hydrogen bonding: A special, particularly strong form of dipole-dipole interaction involving H bonded to F, O, or N.

Relative strengths

Interaction	Typical energy ( $kJ mo l^{- 1}$ )
Covalent bond	$200 - 400$
Ionic bond (lattice)	$700 - 4000$
Hydrogen bond	$8 - 42$
Dipole-dipole	$5 - 25$
London dispersion	$0.05 - 40$ (grows with molecular size)

Van der Waals forces are non-directional and increase with the size of the interacting units. This is why boiling points of similar compounds increase down a group (in the absence of hydrogen bonding): the larger, more polarisable molecules have stronger London forces.

Common Mistakes to Avoid

Watch out

Confusing hydrogen bond with covalent bond. A hydrogen bond is an electrostatic attraction, roughly 10 to 40 times weaker than a covalent bond.
Expecting hydrogen bonding whenever H is present. H-bonding requires H to be covalently attached to F, O, or N. C-H bonds do not give H-bonding because carbon is not electronegative enough.
Assuming intramolecular H-bonds raise boiling point. They do the opposite - by locking H inside one molecule, they prevent intermolecular association and typically lower the boiling point.
Forgetting that HCl does not form hydrogen bonds. Chlorine is about as electronegative as nitrogen but its atomic size is too large; the resulting interaction is too weak to count as a proper H-bond.
Treating all intermolecular forces as van der Waals only. Hydrogen bonds, dipole-dipole, ion-dipole, induced dipole, and London dispersion are all distinct interactions with different strengths and directional properties.
Ignoring the effect of anomalous properties. The high boiling point of water, the density anomaly of ice, and the high surface tension all trace back to H-bonding and should be explained through it in exam answers.

Frequently Asked Questions

Q1. What is a hydrogen bond and how does it form?

A hydrogen bond is an attractive force between an H atom covalently bonded to a highly electronegative atom (F, O, or N) and another electronegative atom carrying a lone pair. It forms because the electronegative atom pulls the electron pair away from H, exposing the H nucleus and giving it a strong partial positive charge. This exposed H then attracts a lone pair on a nearby electronegative atom.

Q2. Why doesn't chlorine form hydrogen bonds despite being electronegative?

Chlorine has electronegativity comparable to nitrogen, but its atomic size is much larger. The negative charge on Cl is spread over a bigger volume (low charge density), so the electrostatic attraction with a nearby $δ^{+}$ hydrogen is too weak to form a proper hydrogen bond. F, O, and N combine high electronegativity with small size, which is why only they participate in hydrogen bonding.

Q3. Which is the strongest hydrogen bond among F-H, O-H, and N-H?

$F - H \dots F$ is the strongest ( $\sim 40 kJ mo l^{- 1}$ ), followed by $O - H \dots O$ ( $\sim 25 kJ mo l^{- 1}$ ), then $N - H \dots N$ ( $\sim 10 kJ mo l^{- 1}$ ). This follows the order of electronegativity: F $>$ O $>$ N.

Q4. How does intramolecular H-bonding differ from intermolecular H-bonding?

Intermolecular H-bonding occurs between two separate molecules (like water molecules linking to each other) and causes molecular association, raising melting and boiling points. Intramolecular H-bonding occurs within one molecule (like in o-nitrophenol), forming a closed ring called a chelate. Because the H is locked inside its own molecule, no association happens between molecules, so intramolecular H-bonding typically lowers melting and boiling points.

Q5. Why is ice less dense than liquid water?

In ice, each water molecule forms four hydrogen bonds arranged tetrahedrally, creating an open, cage-like crystal with lots of empty space. When ice melts, some H-bonds break and the structure partially collapses, allowing molecules to pack closer together. So liquid water has higher density than ice, which is why ice floats.

Q6. Why does ethanol have a much higher boiling point than dimethyl ether, though both are $C_{2} H_{6} O$ ?

Ethanol has an O-H group whose H can form strong intermolecular hydrogen bonds with oxygen of neighbouring ethanol molecules. Dimethyl ether has all its H atoms bonded to carbon (which does not make H a $δ^{+}$ ), so it cannot hydrogen-bond. Ethanol therefore needs much more energy to vaporise (78°C boiling point vs $- 24$ °C for dimethyl ether).

Q7. Among $C H_{3} COOH$ , $N H_{3}$ , HF, and $C H_{4}$ , which shows maximum hydrogen bonding?

HF shows the strongest hydrogen bonding because fluorine has the highest electronegativity of all elements combined with a small size, giving maximum $δ^{+}$ on H and maximum $δ^{-}$ on F. Acetic acid also forms very strong H-bonds (dimers), $N H_{3}$ moderate, and $C H_{4}$ none.

Q8. What is the role of hydrogen bonding in DNA?

In DNA, hydrogen bonds hold the two strands of the double helix together through base pairing. Adenine pairs with thymine (two H-bonds) and guanine pairs with cytosine (three H-bonds). These bonds are strong enough to keep the double helix stable but weak enough to be broken during replication, allowing the strands to separate.

Q9. How do van der Waals forces differ from hydrogen bonds?

Van der Waals forces are weak, non-directional attractions between all molecules (including non-polar ones) arising from temporary or induced dipoles. They typically range from 0.05 to 25 kJ/mol. Hydrogen bonds are stronger (8 to 42 kJ/mol), directional, and require a specific arrangement of H covalently bonded to F, O, or N interacting with another electronegative atom.

Previous year questions on Hydrogen Bonding

3 questions from past papers, each with a step-by-step solution.

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1. What is a Hydrogen Bond?

2. Cause of Hydrogen Bond Formation

Conditions for hydrogen bonding

3. Types of Hydrogen Bonding

Intermolecular hydrogen bonding

Intramolecular hydrogen bonding

4. Effects of Hydrogen Bonding

Boiling point and physical state

Solubility

Density anomaly of ice

Volatility of ortho vs para/meta isomers

Biological importance

5. Van der Waals and Other Intermolecular Forces

Types of intermolecular forces

Relative strengths

Common Mistakes to Avoid

Frequently Asked Questions

Q1. What is a hydrogen bond and how does it form?

Q2. Why doesn't chlorine form hydrogen bonds despite being electronegative?

Q3. Which is the strongest hydrogen bond among F-H, O-H, and N-H?

Q4. How does intramolecular H-bonding differ from intermolecular H-bonding?

Q5. Why is ice less dense than liquid water?

Q6. Why does ethanol have a much higher boiling point than dimethyl ether, though both are C2​H6​O?

Q7. Among CH3​COOH, NH3​, HF, and CH4​, which shows maximum hydrogen bonding?

Q8. What is the role of hydrogen bonding in DNA?

Q9. How do van der Waals forces differ from hydrogen bonds?

Previous year questions on Hydrogen Bonding

Ready to master Chemical Bonding And Molecular Structure?

Q6. Why does ethanol have a much higher boiling point than dimethyl ether, though both are $C_{2} H_{6} O$ ?

Q7. Among $C H_{3} COOH$ , $N H_{3}$ , HF, and $C H_{4}$ , which shows maximum hydrogen bonding?