Both carbon tetrachloride and water contain strongly polar bonds, yet only water behaves like a polar liquid. The difference is not in the bonds but in the shape. Section 3.2 of the book shows how electronegativity sets the bond type and how geometry decides whether bond dipoles add up or cancel.
Chapter 3 of the PECTAA Chemistry 11 book is Chemical Bonding. Section 3.2 links electronegativity to the type of bond, and section 3.2.1 covers dipole moment and the polarity of molecules. Earlier 1st Year notes covered sections 3.3 to 3.9; this note fills in section 3.2. This note is for FSc Part 1 (1st year, Class 11) students in Lahore, Faisalabad, Sargodha, Rawalpindi, Multan, Gujranwala and the other Punjab boards. Gujrat district comes under the Gujranwala board. PECTAA writes the books. Your own board sets your paper.
What to know
Electronegativity and bond type (section 3.2). Electronegativity was met in Chapter 1. The difference in electronegativity between two bonded atoms gives an approximate measure of bond polarity and hints at the bond type:
| Electronegativity difference | Bond type |
|---|---|
| below 0.4 (including zero) | non-polar (pure) covalent |
| 0.4 to 1.8 | polar covalent |
| above 1.8 | ionic |
As the difference rises, covalent character falls and ionic character rises. The book's examples: H–H, difference 0, non-polar; H–Cl, 0.9, polar covalent; NaCl, 2.1, ionic.
Dipole and dipole moment (section 3.2.1). When bonded atoms are different, electron density is shared unevenly. One end carries a partial positive charge (δ+) and the other a partial negative charge (δ−). Such a molecule is a dipole and has a dipole moment, a quantitative measure of polarity, in debye (D). In the book's convention the dipole moment of HF, HCl, HBr, HI or NO points from the δ+ end to the δ− end; in HF, from H to F.
Molecules with more than two atoms. The net dipole moment depends on the shape, because the bond moments add as vectors.
- H₂O: 1.85 D. A linear H–O–H would have zero, so the non-zero value shows water is non-linear; experiments confirm a V-shape.
- SO₂: 1.61 D, also bent, with its dipole pointing the opposite way to water's relative to the central atom.
- H₂S: non-linear; its polar bonds do not cancel, so it has a dipole moment.
Zero net dipole moment. Symmetry can cancel the bond moments:
- BeCl₂: linear, 180°, two equal and opposite Be–Cl moments.
- BF₃: trigonal planar and symmetrical.
- CCl₄: perfectly tetrahedral; four polar C–Cl bonds cancel.
As a rule, molecules with identical groups around the central atom in a regular geometry have zero net dipole moment and are non-polar, even though the individual bonds may be polar.
Key points
- Bond type follows electronegativity difference: <0.4, 0.4–1.8, >1.8.
- Polar bonds do not guarantee a polar molecule.
- Bent molecules (H₂O, SO₂, H₂S) are polar; symmetrical ones (BeCl₂, BF₃, CCl₄) are not.
- Dipole moments are measured in debye (D).
Worked examples
Example 1 (book Quick Check 3.4 a). Why does CO have a dipole moment but CO₂ does not?
- CO is diatomic with two different atoms, so its one polar bond gives a net dipole.
- CO₂ is linear, O=C=O. The two C=O moments are equal and opposite, so they cancel.
Example 2 (book Quick Check 3.4 b). Are the bonds in CCl₄ polar? Is the molecule polar?
- C and Cl differ in electronegativity, so each C–Cl bond is polar.
- The regular tetrahedral shape cancels the four moments. The molecule is non-polar.
Example 3 (classify bonds). Electronegativity differences of 0, 0.9 and 2.1. Classify.
- 0 is below 0.4: non-polar covalent (H–H).
- 0.9 is between 0.4 and 1.8: polar covalent (H–Cl).
- 2.1 is above 1.8: ionic (NaCl).
Example 4 (book Quick Check 3.4 c). Are HF, CH₂Cl₂, O₂ and H₂S polar or non-polar?
- HF: two different atoms with a large electronegativity difference, so polar.
- CH₂Cl₂: the groups around carbon are not all the same (two H, two Cl), so the bond moments do not cancel; polar.
- O₂: two identical atoms, difference zero; non-polar.
- H₂S: non-linear with polar bonds that do not cancel, as the book states; polar.
The rule from the book does the work in each case: same groups in a regular shape cancel; different groups or a bent shape do not.
Common mistakes in the exam
- Saying any molecule with polar bonds is polar. Check the shape.
- Drawing water as linear. Its dipole moment shows it is bent.
- Forgetting the unit of dipole moment: the debye (D).
- Treating 1.8 as covalent. Above 1.8 the book calls the bond ionic.
- Leaving out the reason for zero dipole moment: symmetry.
Exam technique
- For polarity questions, write: bonds polar or not → shape → moments cancel or not → conclusion.
- Sketch the shape with bond-moment arrows.
- Quote the book's values: H₂O 1.85 D, SO₂ 1.61 D.
- Check your board's notice for the sections in your paper.
Practice questions
- Classify a bond with an electronegativity difference of 1.2. Answer: Polar covalent.
- What is the dipole moment of water? Answer: 1.85 D.
- Is BF₃ polar? Answer: No; its trigonal planar symmetry cancels the bond moments.
- Is Br₂ polar? Answer: No; both atoms are identical, so the difference is zero.
- What does a non-zero dipole moment tell you about a triatomic molecule? Answer: It is not linear and symmetrical; its bond moments do not cancel.
Ready to practise this? Continue in 1st Year Chemistry.