MHT-CET Chemistry · Chemical Bonding and Molecular Structure

Dipole Moment, Polarity and Intermolecular Forces

A polar bond has a dipole; whether the whole molecule is polar depends on shape — symmetric molecules cancel their bond dipoles to zero, bent and pyramidal ones don't. The resulting polarity fixes which intermolecular force acts and hence the boiling point.

All Chemical Bonding and Molecular Structure notes MHT-CET Chemistry strategy

Why this matters

Eleven PYQs, and they cluster into three moves the bank repeats every year. First: spot the molecule with zero (or highest, or lowest) dipole moment — always a symmetry judgement, never just 'does it have polar bonds'. Second: name the intermolecular force between a given pair (dipole-induced dipole is the favourite). Third: the hydrogen-bonding boiling-point question — which molecule can (or cannot) H-bond. Master symmetry, the IMF table and the N/O/F rule and the whole subtopic is yours.

Concept 1 of 4

Dipole moment: definition and comparison

Intuition

When two different atoms share electrons, the more electronegative one pulls the shared pair closer, creating a tiny separation of charge. The dipole moment measures the size of that separation — a bigger electronegativity difference means a bigger bond dipole. It is a vector, pointing from the positive end to the negative end.

Definition

Dipole moment facts the bank tests:

The dipole moment $\mu = q \times d$ — charge separated times the distance between the centres of positive and negative charge. It is a vector (has both size and direction).
Units are the debye (D); $1\,\text{D} = 3.336 \times 10^{-30}\,\text{C·m}$ .
A bond's dipole grows with the electronegativity difference of the two atoms. Down a group EN falls (F > Cl > Br > I), so the bond dipole falls: $\mu(\text{CH}_3\text{F}) > \mu(\text{CH}_3\text{Cl}) > \mu(\text{CH}_3\text{Br}) > \mu(\text{CH}_3\text{I})$ .
To compare whole molecules, take the vector sum of the bond dipoles — a lone pair also contributes (its dipole can add to or oppose the bond dipoles).

Dipole moment

\mu = q \times d

\mudipole moment (debye, D)
qmagnitude of the separated charge
ddistance between the positive and negative charge centres

Worked example

Arrange CH3F, CH3Cl, CH3Br and CH3I in decreasing order of dipole moment, given that the C–H part is the same in all four.

All four have the same shape and the same C–H contribution, so the difference comes from the C–X bond dipole.
The C–X bond dipole tracks the electronegativity of X. Down the halogen group EN falls: $\text{F} > \text{Cl} > \text{Br} > \text{I}$ .
So the C–X dipole falls in the same order, and so does the molecular dipole moment.

Answer:

\text{CH}_3\text{F} > \text{CH}_3\text{Cl} > \text{CH}_3\text{Br} > \text{CH}_3\text{I}

;

\text{CH}_3\text{I}

has the lowest dipole moment.

Practice this conceptself-check · 5 quick reps

Try it yourself

Among H2S (~0.95 D), NH3 (~1.47 D), NF3 (~0.23 D) and CHCl3 (~1.04 D), which molecule is the most polar?

Practice — Level 1 (5 reps)

Quick reps to lock in the method. Try each, then check.

1.
Write the formula for dipole moment.
2.
What is the SI-adjacent unit of dipole moment?
3.
Is dipole moment a scalar or a vector?
4.
Which has the larger bond dipole, C–F or C–I?
5.
Why is NH3 more polar than NF3?

From the bank · past-year question

Example 1Chemical Bonding and Molecular StructureMODERATE

Which from following molecules exhibits lowest dipole moment?

[Q59 · 3rd May Shift 2 · 2023]

Dipole moment is a vector — add directions, not magnitudes

NF3 has very polar N–F bonds yet a tiny dipole moment (~0.23 D) because its lone-pair dipole points opposite to the resultant of the N–F bond dipoles and nearly cancels it. In NH3 the two point the same way and add. Always sum the bond dipoles as vectors, and remember the lone pair contributes too.

Bigger electronegativity difference → bigger bond dipole

Down a group electronegativity falls, so the bond dipole falls:

\mu(\text{CH}_3\text{F}) > \mu(\text{CH}_3\text{Cl}) > \mu(\text{CH}_3\text{Br}) > \mu(\text{CH}_3\text{I})

. Don't rank by molecular mass — the heaviest (CH3I) has the smallest dipole.

Drill 2 more on dipole moment: definition and comparison

Concept 2 of 4

Symmetry: when polar bonds give a zero net dipole

Intuition

A molecule can be built entirely from polar bonds and still be non-polar overall — if its shape is symmetric, the bond dipoles point outward evenly and cancel. The moment the shape is bent, pyramidal or otherwise lopsided, the pulls no longer cancel and the molecule has a net dipole.

Definition

The symmetry rule for net dipole moment:

Symmetric shapes cancel to zero. Linear $\text{CO}_2$ , trigonal planar $\text{BF}_3$ , and tetrahedral $\text{CCl}_4$ and $\text{CH}_4$ all have zero net dipole — the identical bond dipoles are arranged so they sum to nothing.
Unsymmetric shapes do not cancel. Bent $\text{H}_2\text{O}$ , pyramidal $\text{NH}_3$ , and $\text{CHCl}_3$ (one C–H breaks the symmetry of CCl4) all have a net dipole.
The test is always: *are the bond dipoles arranged symmetrically?* — not merely *does the molecule contain polar bonds?*
$\text{CHCl}_3$ vs $\text{CCl}_4$ is the classic pair: replacing one Cl with H destroys the tetrahedral cancellation, so CHCl3 is polar while CCl4 is not.

Worked example

Which of CO2, BF3, CCl4 and H2O has a non-zero net dipole moment?

CO2 is linear: the two C=O dipoles point in exactly opposite directions and cancel → zero.
BF3 is trigonal planar and CCl4 is tetrahedral: the identical bond dipoles are symmetric and cancel → zero.
H2O is bent (~104.5°): the two O–H dipoles do not point oppositely, so they add to a net dipole.

Answer:

\text{H}_2\text{O}

— its bent shape prevents the O–H dipoles cancelling; the other three are symmetric and non-polar.

Practice this conceptself-check · 5 quick reps

Try it yourself

Identify the molecule that HAS a dipole moment: BF3, CH4, CHCl3 or CCl4.

Practice — Level 1 (5 reps)

Quick reps to lock in the method. Try each, then check.

1.
Does CO2 have a net dipole moment?
2.
Does BF3 have a net dipole moment?
3.
Does CCl4 have a net dipole moment?
4.
Does CHCl3 have a net dipole moment?
5.
Why is H2O polar but CO2 is not?

From the bank · past-year question

Example 2Chemical Bonding and Molecular StructureEASY

Which of the following molecules has zero dipole moment?

[Q62 · 11th May Shift 2 · 2023]

Polar bonds do NOT guarantee a polar molecule

BF3, CCl4, CH4 and CO2 are built from polar bonds yet have zero net dipole because their symmetric shapes make the bond dipoles cancel. Never answer 'it has a net dipole' just because you see electronegative atoms — check the shape first.

CHCl3 is polar; CCl4 is not

The single most-tested pair here.

\text{CCl}_4

is perfectly tetrahedral, so its four C–Cl dipoles cancel →

\mu = 0

\text{CHCl}_3

swaps one Cl for H, breaking that symmetry, so it has a net dipole. Same atoms, opposite answer — symmetry decides.

Drill 1 more on symmetry: when polar bonds give a zero net dipole

Concept 3 of 4

Types of intermolecular force

Intuition

Whether molecules are polar or not decides which weak force holds them together, and that force in turn sets the melting and boiling point. The bank's job is usually just to name the force acting between a given pair of species — so learn which molecule shows which force and their strength order.

Definition

The intermolecular forces (van der Waals forces, plus hydrogen bonding), weakest to strongest:

London / dispersion forces — momentary induced dipoles between any molecules, including non-polar ones (the only force in $\text{CH}_4$ , $\text{C}_6\text{H}_6$ ). They grow with molecular size, so among HX gases dispersion is largest in HI.
Dipole–induced dipole (Debye) — a polar molecule induces a dipole in a nearby non-polar one (e.g. $\text{NH}_3 + \text{C}_6\text{H}_6$ ).
Dipole–dipole — between two polar molecules; strongest for the largest dipole (among HX, HF has the largest dipole–dipole force).
Hydrogen bonding — a special, strong dipole–dipole force when H is bonded to N, O or F (covered in the next concept).

Force	Acts between	Strength	Example pair
London / dispersion	Any molecules (even non-polar)	Weakest (grows with size)	CH4 + C2H6 Present in every substance; the ONLY force in non-polar molecules. Largest among HX in HI (biggest, most polarisable).
Dipole–induced dipole (Debye)	One polar + one non-polar molecule	Weak	NH3 + C6H6Q
Dipole–dipole	Two polar molecules	Moderate (bigger dipole → stronger)	HF, HCl (polar HX) Strongest dipole–dipole among the hydrogen halides is HF, because F gives the largest bond dipole.
Hydrogen bonding	H on N/O/F, near a lone pair on N/O/F	Strongest of these	H2O, NH3, HF, alcohols

Polar–polar → dipole–dipole; polar–non-polar → dipole-induced dipole; non-polar only → dispersion; H on N/O/F → hydrogen bond.

Practice this conceptself-check · 5 quick reps

Try it yourself

Which intermolecular force acts (a) between NH3 and benzene, and (b) which hydrogen halide has the strongest dipole–dipole force?

Practice — Level 1 (5 reps)

Quick reps to lock in the method. Try each, then check.

1.
Which force acts between a polar and a non-polar molecule?
2.
Which is the only intermolecular force in a non-polar molecule like CH4?
3.
Among HF, HCl, HBr, HI, which has the strongest dipole–dipole force?
4.
Among HF, HCl, HBr, HI, which has the largest dispersion force?
5.
Which force acts between two polar molecules?

From the bank · past-year question

Example 3Chemical Bonding and Molecular StructureEASY

Which among the following forces of attraction is developed between polar and non-polar molecules?

[Q83 · 9th May Shift 1 · 2023]

Dipole–dipole vs dispersion point to different HX

For the hydrogen halides, the strongest dipole–dipole force is in HF (largest bond dipole), but the largest dispersion force is in HI (biggest, most polarisable). The question wording decides — read whether it asks for dipole–dipole or for total van der Waals / dispersion.

Match the force to the pair's polarity

Two polar molecules → dipole–dipole. One polar + one non-polar → dipole-induced dipole. Both non-polar → dispersion only. Check the polarity of both species before naming the force.

Drill 2 more on types of intermolecular force

Concept 4 of 4

Hydrogen bonding and boiling point

Intuition

Hydrogen bonding is an unusually strong attraction that forms when a hydrogen atom is tied to a small, very electronegative atom — nitrogen, oxygen or fluorine. It behaves like a weak extra bond between molecules, so substances that can hydrogen-bond have abnormally high boiling points; those that can't boil lower.

Definition

The rule and its consequence:

Hydrogen bonding needs H covalently bonded to N, O or F (the small, highly electronegative atoms) AND a nearby lone pair on another N, O or F to attract it. Memorise it as the N/O/F rule.
Molecules with an O–H or N–H group (water, alcohols, phenol, ammonia, carboxylic acids, primary and secondary amines) form intermolecular hydrogen bonds and boil high.
Molecules with no H on N/O/F (hydrocarbons like butane; ethers like CH3–O–CH3; tertiary amines with no N–H) cannot donate a hydrogen bond, so they boil low.
$\text{H}_2\text{S}$ barely hydrogen-bonds — S is large and not electronegative enough, which is why $\text{H}_2\text{O}$ boils far higher than $\text{H}_2\text{S}$ .

Worked example

Which of these does NOT form intermolecular hydrogen bonding: ethanol (C2H5OH), butane (C4H10), phenol (C6H5OH) or butan-1-ol (C4H9OH)?

Hydrogen bonding needs H attached to N, O or F.
Ethanol, phenol and butan-1-ol all carry an O–H group → they hydrogen-bond.
Butane is a hydrocarbon: its only bonds are C–C and C–H, with no H on N/O/F.

Answer:Butane — it has no O–H or N–H, so it cannot form intermolecular hydrogen bonds (the others all can).

Practice this conceptself-check · 5 quick reps

Try it yourself

Which has the lowest boiling point: (C2H5)2NH, C2H5N(CH3)2, n-C4H9OH or C2H5COOH?

Practice — Level 1 (5 reps)

Quick reps to lock in the method. Try each, then check.

1.
Hydrogen bonding forms only when H is bonded to which three atoms?
2.
Can butane form intermolecular hydrogen bonds?
3.
Can dimethyl ether (CH3–O–CH3) donate a hydrogen bond?
4.
Why does a tertiary amine like C2H5N(CH3)2 have a low boiling point?
5.
Why does H2O boil higher than H2S?

From the bank · past-year question

Example 4Chemical Bonding and Molecular StructureEASY

Which among the following compounds does NOT form intermolecular hydrogen bonding?

[Q74 · 26 April Shift I · 2025]

No H on N/O/F means no hydrogen-bond donor

A molecule can contain N, O or F and still not hydrogen-bond as a donor if there is no H directly on that atom. Ethers (CH3–O–CH3) and tertiary amines (no N–H) cannot donate a hydrogen bond, so they boil lower than alcohols and primary/secondary amines of similar size.

H2S does not hydrogen-bond like H2O

Students extend H-bonding to

\text{H}_2\text{S}

by analogy with water, but sulphur is large and not electronegative enough —

\text{H}_2\text{S}

shows only weak dipole–dipole/dispersion forces, which is exactly why it is a gas while water is a liquid.

Drill 2 more on hydrogen bonding and boiling point

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Formulas (1)

Dipole moment: definition and comparison
Dipole moment
$\mu = q \times d$

Reference tables (1)

Types of intermolecular force4 rows

Force	Acts between	Strength	Example pair
London / dispersion	Any molecules (even non-polar)	Weakest (grows with size)	CH4 + C2H6 Present in every substance; the ONLY force in non-polar molecules. Largest among HX in HI (biggest, most polarisable).
Dipole–induced dipole (Debye)	One polar + one non-polar molecule	Weak	NH3 + C6H6Q
Dipole–dipole	Two polar molecules	Moderate (bigger dipole → stronger)	HF, HCl (polar HX) Strongest dipole–dipole among the hydrogen halides is HF, because F gives the largest bond dipole.
Hydrogen bonding	H on N/O/F, near a lone pair on N/O/F	Strongest of these	H2O, NH3, HF, alcohols

Polar–polar → dipole–dipole; polar–non-polar → dipole-induced dipole; non-polar only → dispersion; H on N/O/F → hydrogen bond.

Watch out for (8)

Dipole moment is a vector — add directions, not magnitudes
→ Dipole moment: definition and comparison
Bigger electronegativity difference → bigger bond dipole
→ Dipole moment: definition and comparison
Polar bonds do NOT guarantee a polar molecule
→ Symmetry: when polar bonds give a zero net dipole
CHCl3 is polar; CCl4 is not
→ Symmetry: when polar bonds give a zero net dipole
Dipole–dipole vs dispersion point to different HX
→ Types of intermolecular force
Match the force to the pair's polarity
→ Types of intermolecular force
No H on N/O/F means no hydrogen-bond donor
→ Hydrogen bonding and boiling point
H2S does not hydrogen-bond like H2O
→ Hydrogen bonding and boiling point

Mastery check — 5 interleaved questions

Try each one before clicking. Questions are interleaved across the concepts above, not grouped — interleaving sharpens transfer.

Example 1Chemical Bonding and Molecular StructureMODERATE

Which of the following is more polar?

[Q79 · 25 April Shift II · 2025]

Example 2Chemical Bonding and Molecular StructureEASY

Identify the molecule having dipole moment.

[Q75 · May Shift 1 · 2021]

Example 3Chemical Bonding and Molecular StructureMODERATE

Which of the following has dipole-induced dipole interaction as inter molecular force?

[Q91 · 11th May Shift 1 · 2023]

Example 4Chemical Bonding and Molecular StructureEASY

Which of the following molecules can form hydrogen bonding with itself?

[Q60 · Shift 1 · 2023]

Example 5Chemical Bonding and Molecular StructureEASY

What is the value of electronegativity of oxygen?

[Q97 · 4th May Shift 1 · 2023]

Drill every past-year question on this subtopic

11 questions from the bank — paginated, with cart and Word-export support.

Drill the 11 questions

Related notes

Ionic and Covalent Bonding (NDA) →