Electronegativity & polarity
Electronegativity & polarity
Electronegativity and polar bonds
🔓 CRACK IT
In a covalent bond, both atoms attract the shared electrons. Electronegativity describes how strongly an atom attracts that bonding pair.
If one atom attracts the pair more strongly, the electrons are shared unequally. That atom becomes δ− and the other becomes δ+. These are partial charges, not the full charges of ions.
Use the difference between the supplied electronegativity values. A larger difference generally gives a more polar bond, but there is no single numerical cut-off that works for every bond.
- Electronegativity
- The power of an atom to attract the pair of electrons in a covalent bond.
- Polar covalent bond
- Unequal sharing of bonding electrons produces δ+ and δ− ends.
🧪 TEST IT
Work on paper. Attempt each question before opening its answer. These are original CrackingChem questions, written in AQA style.
- Define electronegativity. [2]
- Define a polar covalent bond. [2]
- State the three types of intermolecular force named in the AQA specification. [3]
- State which type acts between all atoms and molecules. [1]
Total: 8 marks
Show answer & marking points
- The power of an atom to attract the bonding pair of electrons [1] in a covalent bond [1].
- A covalent bond with unequal distribution of bonding electrons [1], producing δ+ and δ− ends [1].
- Induced dipole–dipole (van der Waals) forces [1]; permanent dipole–dipole forces [1]; hydrogen bonding [1].
- Induced dipole–dipole (van der Waals) forces [1].
Use Pauling electronegativity values: H 2.1, C 2.5, Br 2.8, Cl 3.0, O 3.5, F 4.0.
- Add δ+ and δ− to H–Cl. [1]
- Add δ+ and δ− to C=O. [1]
- Add δ+ and δ− to O–H. [1]
- Explain how partial charges arise in a polar covalent bond. [2]
- Place H–F, H–Cl and H–Br in decreasing bond polarity. Justify the first position. [2]
Total: 7 marks
Show answer & marking points
- Hδ+–Clδ− [1].
- Cδ+=Oδ− [1].
- Oδ−–Hδ+ [1].
- The more electronegative atom attracts the bonding pair more strongly [1], producing unsymmetrical electron distribution and partial charges [1].
- H–F > H–Cl > H–Br [1]; H–F has the largest electronegativity difference [1].
From bond polarity to molecular polarity
🔓 CRACK IT
A molecule can contain polar bonds and still be non-polar overall. The answer depends on whether the bond dipoles cancel in its three-dimensional shape.
- Identify which bonds are polar and mark their dipole directions.
- Work out the molecular shape, including the effect of lone pairs.
- Check whether the bond dipoles balance in all directions.
- If they cancel completely, the molecule has no permanent dipole. If they do not, the molecule is polar.
CO₂ is linear, so its equal C=O dipoles cancel. H₂O is bent, so its O-H dipoles do not cancel. Draw the shape before deciding.
SO3 and SO2
SO₃ and SO₂ both contain polar S–O bonds. Explain why SO₂ is polar overall but SO₃ is not.
Worked answer
SO3 is symmetrical trigonal planar, so equivalent dipoles cancel. SO2 is bent, so its dipoles do not cancel.
🧪 TEST IT
Work on paper. Attempt each question before opening its answer. These are original CrackingChem questions, written in AQA style.
The figure shows A: CO2, B: BF3, C: H2O and D: CH2Cl2. Every bond to O, F or Cl shown is polar.
- State whether each molecule A–D has a permanent dipole. [4]
- Explain why A and B do not. [2]
- Explain why C and D do. [2]
Total: 8 marks
Show answer & marking points
- A: no; B: no; C: yes; D: yes. One mark each.
- CO2 is linear with equal opposite dipoles that cancel [1]. BF3 is symmetrical trigonal planar, so its three dipoles cancel [1].
- H2O is bent, so dipoles do not cancel [1]. CH2Cl2 has unlike outer atoms in a tetrahedral arrangement, so its dipoles do not cancel [1].
Symmetry and cancellation of dipoles
🔓 CRACK IT
To decide whether a molecule is polar, look at the whole arrangement of bond dipoles. Their sizes and directions must balance for complete cancellation.
BF₃, CO₂ and CCl₄ have equivalent surrounding atoms arranged so the dipoles cancel. NH₃ and H₂O have shapes in which they do not.
A tetrahedral shape alone does not guarantee cancellation. In CH₂Cl₂ the surrounding atoms are different, so its bond dipoles do not balance and the molecule is polar.
| Molecule | Shape | Permanent dipole? |
|---|---|---|
| BF3 | Symmetrical trigonal planar | No: equal dipoles cancel |
| CO2 | Symmetrical linear | No: equal dipoles cancel |
| CCl4 | Symmetrical tetrahedral | No: equal dipoles cancel |
| H2O | Bent | Yes: dipoles do not cancel |
| NH3 | Trigonal pyramidal | Yes: dipoles do not cancel |
| CH2Cl2 | Tetrahedral with unlike surrounding atoms | Yes: dipoles do not cancel |
🧪 TEST IT
Work on paper. Attempt each question before opening its answer. These are original CrackingChem questions, written in AQA style.
The figure shows A: CF4, B: NF3, C: SF4 and D: SF6. Each bond to F is polar.
- Identify the two molecules without permanent dipoles. [2]
- Identify the two with permanent dipoles. [2]
- Explain cancellation in the non-polar molecules. [2]
- Explain why dipoles do not cancel in NF3 and SF4. [2]
- Explain why polar bonds alone cannot establish molecular polarity. [1]
Total: 9 marks
Show answer & marking points
- A CF4 [1] and D SF6 [1].
- B NF3 [1] and C SF4 [1].
- Symmetrical arrangements of identical polar bonds [1] give cancellation [1].
- The lone pair produces an asymmetric molecular shape [1], so bond dipoles do not cancel [1].
- The shape and vector directions of all bond dipoles must also be considered [1].