Q.Which of the following compounds have polar covalent bond: BeCl2, BF3, CO2 and CCl4?
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🔒 Start your 14-day free trial to unlock the full solution →Concept understanding — Dipole Moment Applications
Dipole Moment Applications: From Intuition to Precision
Imagine you have a magnet. One end pulls, the other pushes. Now imagine a molecule that behaves like a tiny magnet — not because of iron, but because of how its electrons are distributed. That's the idea behind a dipole moment.
A dipole moment arises when there's a separation of charge inside a molecule. One end becomes slightly negative (δ−), the other slightly positive (δ+). This imbalance creates a tiny electric "arrow" pointing from positive to negative. The arrow has both a size (how much charge is separated) and a direction (which way the molecule is polar).
The dipole moment is a vector quantity. Its magnitude is given by μ=q×d, where q is the magnitude of separated charge and d is the distance between the charge centers. The unit is the Debye (D).
Now, why does this matter? Because this tiny electric arrow determines how a molecule behaves around other molecules, around electric fields, and even how it interacts with light. Let's see the key applications.
1. Predicting Molecular Shape (Symmetry Check)
This is the most common exam application. If a molecule has polar bonds (like C–Cl or O–H), does it have a net dipole moment? The answer depends on symmetry.
Consider carbon dioxide, CO₂. Each C=O bond is polar (oxygen pulls electrons). But the molecule is linear: O=C=O. The two dipole arrows point in opposite directions and cancel out. Net dipole moment = zero. The molecule is nonpolar.
Now consider water, H₂O. Each O–H bond is polar. But water is bent (104.5°). The two arrows do not cancel — they add up to a net dipole pointing upward through the oxygen. Net dipole moment = 1.85 D. Water is polar.
Symmetry kills polarity. If a molecule has a center of symmetry or identical polar bonds arranged symmetrically, the net dipole moment is zero. This is how you distinguish between linear CO₂ (nonpolar) and bent SO₂ (polar, 1.63 D).
Exam tip: For molecules like CH₄ (tetrahedral, zero dipole) vs. CH₃Cl (tetrahedral but one C–Cl bond, dipole = 1.87 D), the key is whether the polar bonds are arranged so their vectors cancel.
2. Determining Bond Character (Ionic vs. Covalent)
The dipole moment tells you how "unequal" the sharing of electrons is in a bond. A pure covalent bond (like H–H) has zero dipole. A pure ionic bond (like Na⁺Cl⁻) would have a huge dipole — but in reality, ions are separate.
For a bond like H–Cl, the measured dipole moment is 1.08 D. If the bond were 100% ionic (one full electron transferred), the dipole would be much larger (about 6.1 D for the same bond length). The ratio gives you the percent ionic character:
% ionic character=μcalculated for 100% ionicμobserved×100
For HCl: 6.11.08×100≈17.7%. So the H–Cl bond is about 18% ionic, 82% covalent.
This is a standard numerical problem. Remember: μionic=e×d, where e=4.8×10−10 esu (or 1.6×10−19 C in SI). Convert bond length to cm or m accordingly.
3. Intermolecular Forces and Physical Properties
Polar molecules (with a nonzero dipole) experience dipole-dipole interactions — the positive end of one molecule attracts the negative end of another. This is stronger than the London dispersion forces in nonpolar molecules of similar size.
Consequences:
- Boiling points: Polar molecules have higher boiling points than nonpolar ones of similar molar mass. Example: HCl (polar, bp −85°C) vs. F₂ (nonpolar, bp −188°C). Both have about 38 g/mol, but HCl's dipole adds extra attraction.
- Solubility: "Like dissolves like." Polar solutes dissolve in polar solvents (water, ethanol). Nonpolar solutes dissolve in nonpolar solvents (hexane, CCl₄). The dipole moment explains why NaCl dissolves in water but not in oil.
- Dielectric constant: Polar liquids have high dielectric constants (water = 80), meaning they can weaken the electric field between charges. This is why water is such a good solvent for ionic compounds.
Don't confuse dipole moment with boiling point directly. A molecule can have a large dipole but low boiling point if it's very small (like HF, bp 19.5°C, dipole 1.91 D). Hydrogen bonding (a special case of dipole interaction) is even stronger.
4. Reactivity and Orientation in Electric Fields
In an external electric field, polar molecules align themselves with the field. This is the principle behind microwave heating — water molecules in food rotate to align with the alternating microwave field, generating heat through friction.
In organic chemistry, the dipole moment helps predict reaction sites. The negative end of a dipole (where electrons are concentrated) is where electrophiles attack. The positive end (electron-deficient) is where nucleophiles attack. …
Each of BeCl2, BF3, CO2 and CCl4 contains individual bonds that are polar, even though the overall molecules turn out to be non-polar. …
Every bond in BeCl2, BF3, CO2 and CCl4 is polar (unequal electronegativities), but each molecule's symmetric shape cancels the bond dipoles, making the whole molecule non-polar.
A covalent bond is polar whenever the two bonded atoms have different electronegativities, since the shared electron pair is then pulled more towards the more electronegative atom, creating a bond dipole.
- BeCl2: Be (EN approximately 1.6) and Cl (EN approximately 3.2) differ significantly, so each Be-Cl bond is polar.
- BF3: B (EN approximately 2.0) and F (EN approximately 4.0) differ significantly, so each B-F bond is polar.
- CO2: C (EN approximately 2.5) and O (EN approximately 3.5) differ, so each C=O bond is polar.
- CCl4: C (EN approximately 2.5) and Cl (EN approximately 3.2) differ, so each C-Cl bond is polar.
So all four compounds do contain polar covalent bonds. However, this does not automatically make the overall molecule polar - the molecular geometry also matters:
- BeCl2 is linear (Cl-Be-Cl, 180 degrees) - the two equal Be-Cl bond dipoles point in exactly opposite directions and cancel.
- BF3 is trigonal planar (120 degrees between B-F bonds) - the three equal B-F dipoles are symmetric and cancel.
- CO2 is linear (O=C=O, 180 degrees) - the two equal C=O dipoles cancel. …
Showing the 12 most recent of 16 on this concept.
- CBSE 2026Set ANNUAL1 markMCQQ.Which of the following is a non-polar compound ?(a) HCl(b) CH2Cl2(c) CHCl3(d) CCl4
›Reveal solutionSolution
CCl4 is non-polar because of its symmetric tetrahedral shape.
Each C−Cl bond is polar, but in CCl4 the four bonds point to the corners of a regular tetrahedron. The bond dipoles cancel exactly, so the net dipole moment is zero — CCl4 is non-polar. HCl, CH2 …
- CBSE 2025Set ANNUAL1 markMCQQ.Which of the following has zero dipole moment?(a) OF2(b) PCl3(c) SiF4(d) CFCl3
›Reveal solutionSolution
SiF4 has zero net dipole moment due to its symmetric tetrahedral shape.
SiF4 has a tetrahedral geometry with four identical Si-F bonds arranged symmetrically, so their individual bond dipoles cancel exactly, giving a net dipole moment of zero. In contrast, OF2 is bent (net dipole), PCl3 is pyramidal (net dipole, du …
- CBSE 2025Set ANNUAL1 markMCQQ.The dipole moment of BeCl2 is(a) less than zero(b) more than zero(c) zero(d) 1.85 D
›Reveal solutionSolution
Individual Be-Cl bonds are polar, but BeCl2's linear, symmetric geometry makes the two bond dipoles cancel exactly, so the molecule as a whole is non-polar (net dipole moment = 0).
Step 1 — Structure: Beryllium (Be) has only 2 valence electrons and forms 2 sigma bonds with chlorine in BeCl2, with no lone pairs on Be. By VSEPR theory, 2 bonding domains and 0 lone pairs arrange themselves linearly to minimise repulsion, so BeCl2 is linear with a bond angle of 180 degrees (Cl-Be-Cl).
Step 2 — Bond polarity: Each individual Be-Cl bond IS polar, since Cl is significantly more electronegative than Be, so each bond has its own dipole moment vector pointing from Be towards Cl.
…
- CBSE 2025Set sz1 markMCQQ.Select the correct one: Hydrogen molecule is:(a) Polar(b) Ionic(c) Covalent(d) Non-Polar
›Reveal solutionSolution
The hydrogen molecule (H2) is a covalent molecule, and because it is homonuclear (both atoms identical), it is specifically a non-polar covalent molecule.
A bond's polarity depends on the electronegativity difference between the bonded atoms. In H2, both atoms are hydrogen, so the electronegativity difference is exactly zero. The single shared electron pair (the H-H sigma bond) is therefore attracted equally by both nuclei, and the electron cloud is distributed symmetrically. This means H2 has zero dipole moment and is classified as a non-polar molecule (while ionic and polar are ruled out, since there i …
- CBSE 2024Set ANNUAL1 markMCQQ.The unit of dipole moment is(a) coulomb(b) debye(c) farad(d) none of these
›Reveal solutionSolution
Dipole moment, being a very small quantity in SI base units (coulomb-metre), is conventionally expressed in debye (D).
Dipole moment (mu) is defined as charge x distance between the charge centres, so its SI unit is coulomb-metre (C m). However, this SI unit gives inconveniently tiny numbers for molecules, so chemists universally report dipole moments in debye:
1 D = 3.33564 x 10^-30 C m.
…
- CBSE 2024Set ANNUAL1 markMCQQ.Which compound do not show zero dipole moment?(a) CO2(b) BF3(c) H2O(d) CCl4
›Reveal solutionSolution
CO₂, BF₃ and CCl₄ are all symmetric molecules whose individual bond dipoles cancel out to zero net dipole moment; H₂O is bent (unsymmetrical), so its bond dipoles do NOT cancel and it has a non-zero dipole moment.
Check the geometry of each option:
- CO₂: linear (O=C=O), the two C=O bond dipoles point in exactly opposite directions and cancel → μ = 0.
- BF₃: trigonal planar, the three B–F dipoles are symmetric at 120° and cancel → μ = 0.
- CCl₄: tetrahedral, the four C–Cl dipoles cancel by symmetry → μ = 0. …
- CBSE 2024Set ANNUAL1 markMCQQ.Which of the following molecules has zero dipole moment?(a) NH3(b) H2O(c) CHCl3(d) CO2
›Reveal solutionSolution
CO2 has zero net dipole moment because its linear, symmetric shape makes the two polar C=O bond dipoles cancel each other, even though the individual bonds are polar.
Dipole moment (mu) is a vector sum of all individual bond dipoles in a molecule; it is zero only when the geometry is symmetric enough for the bond dipoles to cancel.
- NH3: pyramidal shape (one lone pair on N pointing away from the three N-H bonds) — the bond dipoles and lone pair dipole do not cancel, so NH3 has a net dipole moment (~1.47 D).
- H2O: bent/angular shape (two lone pairs on O) — the two O-H bond dipoles do not cancel, giving a significant net dipole moment (~1.85 D).
- CHCl3: tetrahedral but NOT symmetric (three C-Cl bonds and one C-H bond, which are not identical) — the dipoles do not cancel, giving a net dipole moment (~1.04 D). …
- CBSE 2023Set ANNUAL1 markMCQQ.The dipole moment of BeCl2 is(a) less than zero(b) more than zero(c) zero(d) 1.85 D
›Reveal solutionSolution
BeCl2 is sp hybridised and linear (Cl-Be-Cl bond angle 180°); the two identical, oppositely-directed Be-Cl bond dipoles cancel, giving a net dipole moment of zero despite each individual bond being polar.
Beryllium (Group 2) has 2 valence electrons and forms two sigma bonds to two Cl atoms, with no lone pairs on Be. VSEPR predicts a linear shape (bond angle 180°) since there are only 2 electron domains around Be. Each Be-Cl bond is individually polar (Cl more electronegative), but because the mo …
- CBSE 2023Set ANNUAL1 markMCQQ.Which of the following has the highest dipole moment?(a) NH3(b) NF3(c) BF3(d) BeCl2
›Reveal solutionSolution
NH3 has the highest dipole moment of the four because its bond dipoles and lone-pair dipole add up, unlike the other three.
Consider each molecule's shape and how bond dipoles combine:
- BeCl2: linear (sp hybridised Be, no lone pair) — the two Be-Cl bond dipoles point in opposite directions and cancel exactly. Net dipole moment = 0.
- BF3: trigonal planar (sp2 B, no lone pair) — three identical B-F bond dipoles at 120 degrees cancel by symmetry. Net dipole moment = 0.
- NF3: pyramidal (sp3 N, one lone pair) — the lone pair points away from the three N-F bonds. F is more electronegative than N, so the N-F bond dipoles point toward F, while the lone pair's dipole points the other way, partly cancelling the resultant bond dipole. Net dipole moment is small (~0.24 D). …
- CBSE 2022Set ANNUAL1 markMCQQ.Which compound do not show zero dipole moment ?(a) CO2(b) BF3(c) H2O(d) CCl4
›Reveal solutionSolution
Only H2O among the four is a bent, asymmetric molecule, so only it has a non-zero net dipole moment.
A molecule's net dipole moment is the vector sum of its individual bond dipoles.
- CO2: linear (O=C=O), the two C=O dipoles point in opposite directions and cancel exactly → net dipole = 0.
- BF3: trigonal planar, the three B–F dipoles are arranged symmetrically at 120° and cancel → net dipole = 0.
- CCl4: tetrahedral, the four C–Cl dipoles cancel by symmetry → net dipole = 0. …
- CBSE 2022Set ANNUAL1 markMCQQ.The value of dipole moment _______ with increasing bond angle.(a) increase(b) decrease(c) no change
›Reveal solutionSolution
For a bent (AX2-type) molecule with two equal bond dipoles mu at bond angle theta, the resultant dipole moment is 2 mu cos(theta/2); as theta increases, cos(theta/2) decreases, so the net dipole moment decreases.
Consider a molecule with two identical polar bonds, each of bond dipole mu, meeting at bond angle theta. Their vector sum (net molecular dipole moment) is given by:
net dipole = 2 mu cos(theta/2)
…
- CBSE 2022Set ANNUAL1 markMCQQ.In which of the following has dipole moment zero?(a) ClF(b) PCl3(c) SiF4(d) CFCl3
›Reveal solutionSolution
SiF₄ is symmetric tetrahedral, so bond dipoles cancel → μ = 0 — option (c).
From NCERT Class 11 Chemistry (Chemical Bonding and Molecular Structure):
- ClF: heteronuclear diatomic → permanent dipole.
- PCl₃: pyramidal (lone pair on P) → net dipole. …
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