Q.Give reason for the higher boiling point of ethanol in comparison to methoxymethane.
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Boiling Point Trends (Organic Compounds)
A substance's boiling point is set by how much energy is needed to overcome the attractive forces HOLDING its molecules together in the liquid — the stronger those intermolecular forces, the higher the boiling point.
The Forces, Weakest to Strongest
- Van der Waals (London dispersion) forces — present in every molecule, and they grow stronger as the molecule gets bigger (more electrons, larger surface area of contact between neighbouring molecules) and more polarisable.
- Dipole–dipole forces — present in polar molecules, add an extra attraction on top of dispersion forces.
- Hydrogen bonding — present when H is bonded directly to N, O, or F; much stronger than ordinary dipole–dipole attraction, and it raises the boiling point sharply compared to a similarly-sized molecule without it.
Trend 1: Down a Series of Halogens (Same Alkyl Group)
For a fixed R group, boiling point rises as the halogen gets heavier: R−I>R−Br>R−Cl>R−F. This looks surprising at first, since electronegativity (and so bond polarity/dipole moment) actually DECREASES down the group — but boiling point here is dominated by the growing size and polarisability of the halogen atom (stronger dispersion forces), which outweighs the shrinking dipole contribution.
The measured values for the methyl, ethyl and propyl halides show this rise clearly:
Trend 2: Chain Length and Branching
- Longer chains (more carbons) have more surface area for van der Waals contact between neighbouring molecules, so boiling point rises with chain length within a homologous series.
- Branching LOWERS boiling point compared to a straight-chain isomer of the same molecular formula — a more compact, spherical shape has less surface-to-surface contact with neighbouring molecules, weakening the dispersion forces. (E.g. neopentane boils well below n-pentane.)
Trend 3: Hydrogen Bonding Beats Molecular Mass …
Why this formula?
Boiling Point Trends: Why They Happen
Boiling point is the temperature at which a liquid's vapor pressure equals the external atmospheric pressure. To understand why boiling points follow certain trends, we must first understand what determines vapor pressure.
The Core Idea: Intermolecular Forces
A liquid boils when its molecules have enough kinetic energy to overcome the intermolecular forces (IMFs) holding them together in the liquid phase. Stronger IMFs → harder to escape → lower vapor pressure at a given temperature → higher boiling point.
There is no single "formula" for boiling point, but the relationship is captured by the Clausius–Clapeyron equation, which links vapor pressure (P) to temperature (T) and the enthalpy of vaporization (ΔHvap):
lnP=−RΔHvap⋅T1+C
Where:
- P = vapor pressure
- ΔHvap = enthalpy of vaporization (energy needed to vaporize 1 mole)
- R = gas constant
- T = absolute temperature (Kelvin)
- C = constant (depends on substance)
Why this formula makes sense
- ΔHvap is large when IMFs are strong — more energy is needed to separate molecules.
- At boiling point, P=Patm (usually 1 atm). So a substance with larger ΔHvap needs a higher T to reach that pressure.
Thus, boiling point ∝ strength of intermolecular forces.
The Four Key Trends (with Reasoning)
1. Trend across a period (e.g., Period 2: CH₄ → NH₃ → H₂O → HF)
| Molecule | IMFs present | Boiling point (°C) |
|---|---|---|
| CH₄ | London dispersion only | -161 |
| NH₃ | Dispersion + H-bonding | -33 |
| H₂O | Dispersion + H-bonding (2 per molecule) | 100 |
| HF | Dispersion + H-bonding | 19 |
Why?
- CH₄ is nonpolar — only weak London dispersion forces.
- NH₃, H₂O, HF have hydrogen bonding (strongest IMF).
- H₂O forms two H-bonds per molecule (donor + acceptor), while NH₃ forms one and HF forms one — hence H₂O has the highest boiling point.
Key insight: Hydrogen bonding dominates over molecular mass in small molecules.
2. Trend down a group (e.g., Halogens: F₂ → Cl₂ → Br₂ → I₂)
| Molecule | Molar mass (g/mol) | Boiling point (°C) |
|---|---|---|
| F₂ | 38 | -188 |
| Cl₂ | 71 | -34 |
| Br₂ | 160 | 59 |
| I₂ | 254 | 184 |
Why?
- All are nonpolar — only London dispersion forces.
- Dispersion force strength increases with number of electrons (larger molar mass → more polarizable electron cloud → stronger temporary dipoles).
- So boiling point increases down the group.
Key insight: For nonpolar molecules, molar mass (electron count) is the primary factor.
3. Branching in alkanes (e.g., C₅H₁₂ isomers)
| Isomer | Boiling point (°C) |
|---|---|
| n-pentane (straight chain) | 36 |
| 2-methylbutane (branched) | 28 |
| 2,2-dimethylpropane (highly branched) | 10 |
Why?
- All have same molecular formula — same molar mass. …
Concept: Boiling Point Trends — Boiling point depends on the strength of intermolecular forces. Stronger forces require more energy to overcome.
Reasoning:
- Ethanol (CH3CH2OH) has an —OH group, allowing intermolecular hydrogen bonding between molecules.
- Methoxymethane (CH3OCH3) is an ether; its molecules are held together only by dipole-dipole interactions and weak London forces — no hydrogen bonding because the oxygen is bonded to two carbons, leaving no H attached to O. …
The key is intermolecular forces: ethanol has strong hydrogen bonding between its molecules, while methoxymethane only has weak dipole-dipole interactions. This makes ethanol's boiling point much higher (78.4∘C vs −24.8∘C).
Why boiling points differ — the core idea
Boiling point is a measure of how much energy you need to pull molecules apart from each other in the liquid state. The stronger the forces holding the molecules together, the higher the temperature required to overcome them. So the question really is: what kind of intermolecular forces exist in each compound?
Ethanol (CH3CH2OH) and methoxymethane (CH3OCH3) are both small molecules with similar molecular masses (46 g/mol). But their structures differ in one critical way: ethanol has an —OH group, while methoxymethane has an ether linkage (—O—). That single —OH group changes everything.
Step-by-step reasoning
1. Identify the intermolecular forces possible in each molecule
For any molecule, the possible forces are:
- London dispersion forces (present in all molecules, increase with size/surface area)
- Dipole-dipole interactions (present if the molecule has a permanent dipole)
- Hydrogen bonding (a special, much stronger dipole-dipole interaction — requires H bonded to N, O, or F)
Ethanol has an O—H bond. The hydrogen is attached directly to oxygen, so it can participate in hydrogen bonding. Methoxymethane has no O—H or N—H bond — its hydrogens are all attached to carbon. So it cannot form hydrogen bonds.
2. Compare the strength of forces in each case
In ethanol:
- London forces (weak, similar to methoxymethane)
- Dipole-dipole interactions (moderate, from the C—O and O—H polar bonds)
- Hydrogen bonding (strong — each ethanol molecule can form up to three hydrogen bonds: two through the lone pairs on oxygen, one through the H of the —OH group)
In methoxymethane:
- London forces (similar magnitude)
- Dipole-dipole interactions (moderate, from the C—O—C polar bonds)
- No hydrogen bonding — the H atoms are all bonded to carbon, not to O or N
A common mistake is to think that because methoxymethane contains oxygen, it can hydrogen bond. It cannot — the hydrogen must be directly bonded to the oxygen (or nitrogen/fluorine) for hydrogen bonding to occur. In methoxymethane, all hydrogens are C—H, which do not form hydrogen bonds.
3. Quantify the difference
The boiling point of ethanol is 78.4∘C. The boiling point of methoxymethane is −24.8∘C. That's a difference of over 100∘C — enormous for molecules of the same mass. …
Method: Intermolecular Force Analysis (Hydrogen Bonding vs. Dipole-Dipole)
This method compares the type and strength of intermolecular forces between molecules to explain boiling point trends.
Step 1: Identify the molecules and their structures
- Ethanol (CH3CH2OH): Contains an –OH group (hydroxyl group).
- Methoxymethane (CH3OCH3): An ether with an –O– linkage, no –OH group.
Step 2: Determine the dominant intermolecular forces
| Molecule | Intermolecular Forces Present |
|---|---|
| Ethanol | Hydrogen bonding (due to O–H bond) + dipole-dipole + London dispersion |
| Methoxymethane | Dipole-dipole (due to C–O bond polarity) + London dispersion only |
Key point: Hydrogen bonding is significantly stronger than dipole-dipole interactions.
Step 3: Compare boiling points based on force strength
- Ethanol molecules must overcome hydrogen bonds (strong) to enter the gas phase → requires more energy → higher boiling point.
- Methoxymethane molecules only need to overcome dipole-dipole and dispersion forces (weaker) → requires less energy → lower boiling point.
Step 4: State the result …
Here are the most common mistakes students make when comparing the boiling points of ethanol and methoxymethane, along with the correct reasoning and how to avoid each error.
The Core Concept (Why Ethanol Boils Higher)
Ethanol (CH3CH2OH) has a significantly higher boiling point (78.4∘C) than methoxymethane (CH3OCH3, −24.8∘C) because ethanol molecules can form intermolecular hydrogen bonds, while methoxymethane cannot.
- Ethanol: Contains an −OH group. The hydrogen atom (bonded to oxygen) is strongly attracted to the lone pairs on the oxygen of a neighboring molecule. This creates a strong, directional intermolecular force.
- Methoxymethane: Contains a C−O−C linkage. While it is polar, the hydrogen atoms are bonded to carbon, not oxygen. It cannot form hydrogen bonds with itself. Its intermolecular forces are limited to dipole-dipole interactions and London dispersion forces.
Common Mistake #1: Confusing Intramolecular vs. Intermolecular Bonds
The Mistake: Students say, "Ethanol has a higher boiling point because it has a strong O-H bond inside the molecule."
Why it's wrong: Boiling point depends on the strength of forces between molecules (intermolecular), not the strength of bonds within a molecule (intramolecular). Breaking the O-H bond would require chemical decomposition, not boiling.
How to Avoid:
- Use the "Phase Change" Rule: If you are changing state (liquid to gas), you are overcoming intermolecular forces.
- Memorize the hierarchy: Boiling point is determined by: Hydrogen bonding > Dipole-dipole > London dispersion forces.
Common Mistake #2: Ignoring the "Hydrogen Bonding" Requirement
The Mistake: Students assume any molecule with oxygen and hydrogen can hydrogen bond. They say, "Methoxymethane has oxygen and hydrogen, so it also has hydrogen bonding."
Why it's wrong: For hydrogen bonding to occur between molecules of the same compound, the hydrogen must be directly bonded to a highly electronegative atom (N, O, or F). In methoxymethane (CH3−O−CH3), all hydrogens are bonded to carbon. The hydrogen atoms are not sufficiently δ+ to form a strong hydrogen bond with the oxygen of another molecule.
How to Avoid:
- Check the bonding pattern: Look for an -OH, -NH, or -FH group. If the H is attached to C, it cannot form a hydrogen bond with another molecule of the same compound.
- Draw the structure: Write out the full structural formula. Circle the H atoms. If any H is bonded to C, it is not a hydrogen-bond donor.
Common Mistake #3: Forgetting London Dispersion Forces
The Mistake: Students say, "Ethanol has hydrogen bonding, methoxymethane has dipole-dipole forces, so ethanol wins." This is correct but incomplete.
Why it's a problem: While hydrogen bonding is the dominant factor, both molecules also have London dispersion forces. Methoxymethane is an isomer of ethanol (same molecular formula C2H6O), so their molecular masses are identical. However, the shape matters. Methoxymethane is more compact, while ethanol is a linear chain. A more linear shape can lead to slightly stronger dispersion forces, but this is overwhelmed by the hydrogen bonding in ethanol.
How to Avoid:
- Acknowledge all forces: When comparing boiling points, list all three types of intermolecular forces present in each molecule. …
Showing the 12 most recent of 24 on this concept.
- CBSE 2026Set 56/2/11 markMCQQ.Identify the correct increasing order of boiling points of the given compounds : (A) Propan-1-ol < butan-1-ol < butan-2-ol < pentan-1-ol (B) Pentan-1-ol < butan-1-ol < butan-2-ol < Propan-1-ol (C) Propan-1-ol < butan-2-ol < butan-1-ol < pentan-1-ol (D) Butan-1-ol < Butan-2-ol < Propan-1-ol < Pentan-1-ol
›Reveal solutionSolution
Boiling points of alcohols depend on chain length (more carbons → higher bp) and branching (more branching → lower bp). The correct order is Propan-1‑ol < butan‑2‑ol < butan‑1‑ol < pentan‑1‑ol, which matches option (C).
Why boiling points of alcohols behave this way
Alcohols boil at much higher temperatures than hydrocarbons of similar mass because of hydrogen bonding between the –OH groups. Two factors control the boiling point within a family of alcohols:
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Chain length – A longer carbon chain means more surface area for van der Waals forces. These weak attractions add up, so a larger molecule needs more energy (higher temperature) to escape into the vapour phase. For straight‑chain alcohols, boiling point rises steadily as the number of carbons increases.
-
Branching – When the –OH group is attached to a secondary or tertiary carbon (as in butan‑2‑ol), the molecule becomes more compact. A compact shape reduces the surface area available for van der Waals interactions, so the boiling point drops compared to its straight‑chain isomer. The hydrogen‑bonding ability is roughly the same for all isomers (one –OH per molecule), so the difference comes from the weaker London forces in the branched form.
Watch outA common mistake is to think that branching increases boiling point because the molecule looks “more crowded”. In reality, branching decreases the surface area and therefore weakens the intermolecular forces. Always compare chain length first, then branching.
Step‑by‑step reasoning
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Identify the compounds and their carbon counts
- Propan‑1‑ol: 3 carbons, straight chain.
- Butan‑1‑ol: 4 carbons, straight chain.
- Butan‑2‑ol: 4 carbons, branched (the –OH is on carbon 2).
- Pentan‑1‑ol: 5 carbons, straight chain.
-
Order by chain length
Longer chain → higher boiling point. So the 5‑carbon alcohol (pentan‑1‑ol) should have the highest bp, and the 3‑carbon alcohol (propan‑1‑ol) the lowest. The two 4‑carbon alcohols will sit in between.
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Compare the two C₄ isomers …
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- CBSE 2026Set ANNUAL1 markMCQQ.Assertion (A): Boiling point of alkanes decreases with increase in molecular mass. Reason (R): Intermolecular Vander Waals forces increase with increase in molecular size or surface area of the molecules.(a) Both Assertion (A) and Reason (R) are true and Reason (R) is the correct explanation of Assertion (A).(b) Both Assertion (A) and Reason (R) are true, but Reason (R) is not the correct explanation of Assertion (A).(c) Assertion (A) is true, but Reason (R) is false.(d) Assertion (A) is false, but Reason (R) is true.
›Reveal solutionSolution
The Assertion has the trend backwards — alkane boiling points INCREASE with molecular mass — while the Reason correctly describes why (stronger van der Waals forces with larger surface area).
Assertion: 'Boiling point of alkanes decreases with increase in molecular mass' — this is false. In reality, as molecular mass (chain length) increases, boiling point increases (e.g. methane −162°C → butane −0.5°C → octane 126°C).
…
- CBSE 2025Set X11 markMCQQ.Two compounds 'A' and 'B' were being tested for their boiling points. It was observed that 'A' started boiling after 'B', when both were subjected to same conditions. If the compound 'B' is acetone, which of the following can be compound 'A'?(a) Propanal(b) Propan-1-ol(c) Methoxyethane(d) n-Butane
›Reveal solutionSolution
"A boils after B" → A has the higher boiling point; among the options only propan-1-ol (H-bonding) boils higher than acetone, so A = propan-1-ol.
B is acetone (propanone), b.p. ≈56∘C. "A started boiling after B" means A needs a higher temperature, i.e. A has a higher boiling point than acetone. Comparing approximate boiling points of the options (all C3/C4 molecules of similar mass):
Compound Approx. b.p. Reason (a) Propanal ≈49∘C dipole–dipole only (b) Propan-1-ol ≈97∘C strong intermolecular H-bonding - CBSE 2025Set X11 markMCQQ.Select the correct order of melting points of isomeric dichlorobenzenes.(a) o-dichlorobenzene > m-dichlorobenzene > p-dichlorobenzene(b) p-dichlorobenzene > m-dichlorobenzene > o-dichlorobenzene(c) p-dichlorobenzene > o-dichlorobenzene > m-dichlorobenzene(d) m-dichlorobenzene > o-dichlorobenzene > p-dichlorobenzene
›Reveal solutionSolution
Melting point depends on how well molecules pack in the crystal; the symmetrical para isomer packs best (highest m.p.), giving the order para > ortho > meta.
For isomeric dichlorobenzenes, melting point is governed mainly by crystal packing / molecular symmetry rather than by intermolecular force magnitude:
- p-dichlorobenzene is the most symmetrical, so it packs most efficiently into the crystal lattice and has the highest melting point (≈53∘C).
- o-dichlorobenzene (≈−17∘C) packs better than the meta isomer. …
- CBSE 2025Set X11 markMCQQ.Sufficient amount of 2-methylpropan-2-ol heated with 20% phosphoric acid at 358 K gives main product 'X' with the elimination of water and tert-butyl alcohol undergoes dehydration when it is passed over heated copper at 573 K gives 'Y' Pick the correct statement regarding X and Y.(a) The boiling points of 'X' and 'Y' are equal(b) The boiling point of 'X' is greater than the boiling point of 'Y'(c) The boiling point of 'X' is lesser than the boiling point of 'Y'(d) At room temperature both 'X' and 'Y' exists as a solids
›Reveal solutionSolution
Acid dehydration and passing over hot copper both convert 2-methylpropan-2-ol to the same alkene (2-methylpropene), so X = Y and their boiling points are equal — option (a).
2-Methylpropan-2-ol is a tertiary alcohol (tert-butyl alcohol), (CH3)3C–OH.
- With 20% phosphoric acid at 358 K it undergoes acid-catalysed dehydration (elimination of water) to give the alkene: (CH3)3C–OH→(CH3)2C=CH2+H2O, so X = 2-methylpropene (isobutylene). …
- CBSE 2025Set D1 markMCQQ.At room temperature, formaldehyde is(a) gas(b) liquid(c) solid(d) none of these
›Reveal solutionSolution
Formaldehyde, the first member of the aldehyde series, is a gas at ordinary temperature (b.p. about −19 °C).
Formaldehyde (methanal, HCHO) is the lowest aldehyde. It has a very low boiling point (about −19 °C), so at room temperature it exists as a colourless, pungent-smelling gas. Its 40% aqueous solution is calle …
- CBSE 2025Set ANNUAL1 markQ.Fill in the blank: The boiling point of methanol is ________ K.
›Reveal solutionSolution
Methanol (CH3OH), the smallest alcohol, boils at about 338 K (64.7 degrees C) at atmospheric pressure.
Methanol's boiling point of ~338 K is relatively low among common alcohols because of its small molecular size (weaker van der Waals/London forces), even though, like other alcohols, it is capable of intermolecu …
- CBSE 2025Set ANNUAL1 markMCQQ.Which of the following compounds has highest melting point?(a) 1,2-dichlorobenzene (ortho-dichlorobenzene, structure drawn)(b) 1,3-dichlorobenzene (meta-dichlorobenzene, structure drawn)(c) 1,4-dichlorobenzene (para-dichlorobenzene, structure drawn)(d) All have same melting point.
›Reveal solutionSolution
Melting point depends on how efficiently molecules pack into a crystal lattice, not just molecular weight — the highly symmetric para isomer packs far better than the less symmetric ortho and meta isomers, giving it a much higher melting point.
All three dichlorobenzenes have the same molecular formula and molecular weight, so their melting-point difference comes purely from crystal packing efficiency. p-Dichlorobenzene is linear and symmetric, so molecules stack very closely and regularly in the solid lattice, maximising van der Waals contact — this raises its melting point sharply (~53°C). The or …
- CBSE 2025Set ANNUAL1 markQ.Arrange the following compounds in increasing order of their boiling points: CH3CHO, CH3CH2OH, CH3OCH3, CH3CH2CH3
›Reveal solutionSolution
Boiling point here tracks the strength of intermolecular forces: propane (only weak van der Waals forces) boils lowest, dimethyl ether (weak dipole-dipole, no H-bonding) next, acetaldehyde (stronger dipole-dipole from the polar C=O) next, and ethanol (hydrogen-bonded) boils highest.
All four compounds have comparable molar mass (propane 44, dimethyl ether 46, acetaldehyde 44, ethanol 46 g mol−1), so the boiling-point order is decided almost entirely by the type of intermolecular attraction available, not by size:
- CH3CH2CH3 (propane): a non-polar hydrocarbon; molecules are held together only by weak instantaneous dipole–induced dipole (London/van der Waals) forces. Lowest boiling point (real value ≈ −42 °C).
- CH3OCH3 (dimethyl ether): the C–O–C linkage gives the molecule a small permanent dipole, so molecules attract each other by dipole–dipole forces, stronger than propane's dispersion forces alone but the ether oxygen has no O–H bond, so no hydrogen bonding is possible. Boils higher than propane (real value ≈ −24 °C). …
- CBSE 2025Set ANNUAL1 markMCQQ.The correct order of boiling points of alcohols having the same number of Carbon atoms is ...................... .(a) 2° > 1° > 3°(b) 1° > 2° > 3°(c) 3° > 1° > 2°(d) 3° > 2° > 1°
›Reveal solutionSolution
Among isomeric alcohols, boiling point falls as branching increases, because branching reduces the effective surface area available for intermolecular hydrogen bonding and van der Waals interactions.
All isomeric alcohols with the same molecular formula can hydrogen-bond through their –OH group, but a straight-chain (primary) alcohol packs more efficiently and has a larger surface area for van der Waals contact between molecules than a branched (te …
- CBSE 2025Set ANNUAL1 markMCQQ.Which of the following has the highest melting point?(a) o-xylene(b) m-xylene(c) p-xylene(d) Toluene
›Reveal solutionSolution
p-Xylene has the highest melting point because its high molecular symmetry allows the most efficient crystal packing.
Among the three xylene isomers (o-, m-, p-) and toluene, melting point depends heavily on how symmetrically the molecules can pack into a solid lattice (not just on molecular weight or boiling point). p-Xylene, with its methyl groups symmetrically placed at opposite (1,4) positions on the ring, packs most efficiently into a crystal lattice, giving it a distinct …
- CBSE 2024Set ANNUAL1 markQ.Why has propanol higher boiling point than propane?
›Reveal solutionSolution
Boiling point depends on the strength of intermolecular forces that must be overcome; propanol's -OH group enables hydrogen bonding between molecules, a much stronger force than the weak van der Waals (London dispersion) forces that are all propane has.
Propane (CH3-CH2-CH3) is a non-polar hydrocarbon with no functional group capable of hydrogen bonding. Its molecules are held together only by weak van der Waals (induced-dipole) forces, so relatively little energy is needed to separate them - it boils at a very low temperature (-42 degree C) and is a gas at room temperature.
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