Electronegativity and bond polarityAQA A-Level Chemistry: Revision notes
Section 1
Electronegativity
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. It is measured on the Pauling scale: fluorine is the most electronegative element (4.0), followed by oxygen (3.5), nitrogen and chlorine (3.0).
Electronegativity increases across a period (more protons, similar shielding, smaller radius) and decreases down a group (larger radius and more shielding, so the bonding pair is further from the nucleus).
Electronegativity is about attracting the shared pair in a covalent bond. Do not confuse it with ionisation energy or electron affinity.
Section 2
Polar and non-polar bonds
If the two atoms in a bond have equal electronegativity (e.g. Cl–Cl, H–H), the bonding pair is shared equally: the bond is non-polar.
If they differ, the electron density is unsymmetrical and the bond is polar covalent: the more electronegative atom is partially negative (δ−) and the other partially positive (δ+). For example in HCl: Hδ+–Clδ−. The larger the difference in electronegativity, the more polar the bond; when it is very large the bonding approaches ionic.
Section 3
Permanent dipoles in molecules
A permanent dipole exists in a molecule when the bond polarities do not cancel, so one end of the molecule is slightly positive and the other slightly negative. Examples: HCl, H₂O, NH₃, CHCl₃, SO₂.
To decide, (1) find which bonds are polar, (2) draw the shape, (3) ask whether the dipoles cancel. A molecule with a lone pair or with different atoms attached is usually not symmetrical, so it has a dipole.
Section 4
Polar bonds but no permanent dipole
Some molecules contain polar bonds but are symmetrical, so the bond dipoles are equal in size and cancel:
- CO₂ and BeCl₂: linear
- BF₃: trigonal planar
- CH₄ and CCl₄: tetrahedral
- SF₆: octahedral
- XeF₄: square planar
By contrast, if one atom is replaced (CHCl₃, CH₃Cl) or lone pairs break the symmetry (H₂O, NH₃, SO₂), the dipoles do not cancel.
Do not write that CO₂ is non-polar because the bonds are non-polar. Say that the bonds are polar but the dipoles cancel because the molecule is linear.
Section 5
Worked example: is NF₃ polar?
Fluorine (4.0) is more electronegative than nitrogen (3.0), so each N–F bond is polar with N δ+ and F δ−. NF₃ has three bonding pairs and one lone pair on nitrogen, so it is pyramidal. A pyramidal molecule is not symmetrical, so the dipoles do not cancel and NF₃ has a permanent dipole.
Compare BF₃: three polar bonds, but trigonal planar and symmetrical, so there is no permanent dipole. The difference in shape comes from the lone pair on nitrogen, so always link polarity to shape.
Section 6
Why it matters
Permanent dipoles lead to permanent dipole-dipole forces between molecules, which raise the boiling point compared with a non-polar molecule of similar size. Very polar bonds, where hydrogen is bonded to fluorine, oxygen or nitrogen, give hydrogen bonding. Polar molecules are also deflected by a charged rod, whereas non-polar liquids are not.
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Electronegativity and bond polarity
- Hydrogen chloride, HCl, is a gas that dissolves in water to form hydrochloric acid. The two atoms in an HCl molecule share one pair of electrons, and chlorine is more electronegative than hydrogen.Explain why the bond in HCl is polar.2 marks
- Carbon dioxide, CO₂, and sulfur dioxide, SO₂, both contain polar bonds between oxygen and a central non-metal atom. Oxygen is more electronegative than both carbon and sulfur. A CO₂ molecule is linear, whereas an SO₂ molecule is bent.Explain why SO₂ has a permanent dipole but CO₂ does not.2 marks
- Chlorine is found in several molecules, including chlorine gas, Cl₂, tetrachloromethane, CCl₄, and trichloromethane, CHCl₃. Pauling electronegativity values are 2.5 for carbon, 2.1 for hydrogen and 3.0 for chlorine.Explain why the bond in Cl₂ is non-polar but the C–Cl bond in CCl₄ is polar, stating the partial charge on each atom in C–Cl.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).