Bond polarity — AQA A-Level Chemistry
Test yourself on Bond polarity with AQA A-Level practice questions.
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Bond polarity explained
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself.
Read the full explanation
It is measured on the Pauling scale, where fluorine is the highest at 4.0, followed by oxygen at 3.5, then nitrogen and chlorine at 3.0. Across a period it rises because nuclear charge increases while the bonding electrons stay in the same shell with similar shielding, so they are pulled in harder. Down a group it falls because the bonding pair lies in a shell further from the nucleus and is shielded by more inner shells. Only the shared pair is involved; the term says nothing about an atom's attraction for electrons in general.
The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical.
When two atoms of different electronegativity share a pair of electrons, the pair is not held midway. The more electronegative atom pulls it closer, so the electron density around that atom increases and around the other decreases. The result is a partial negative charge, written δ⁻, on the more electronegative atom, and a partial positive charge, δ⁺, on the other. These are fractions of a charge, not full ionic charges; no electron has been transferred. The larger the electronegativity difference the more unsymmetrical the distribution, and at a large enough difference the bond is better described as ionic. When the two atoms are identical, as in chlorine or oxygen, the distribution is symmetrical and the bond is non-polar.
This produces a polar covalent bond, and may cause a molecule to have a permanent dipole.
A polar bond has a dipole: a separation of charge with a direction. Whether the whole molecule has a permanent dipole depends on how those bond dipoles are arranged in space, because they add as vectors. If the polar bonds point symmetrically outwards so their dipoles cancel, the molecule is non-polar overall despite its polar bonds: carbon dioxide is linear, tetrachloromethane is tetrahedral and boron trifluoride is trigonal planar, and all three cancel. If the arrangement is unsymmetrical, usually because of a lone pair on the central atom or different atoms attached, the dipoles do not cancel and the molecule has a permanent dipole. Water and ammonia are polar for exactly this reason.
Students should be able to: use partial charges to show that a bond is polar explain why some molecules with polar bonds do not have a permanent dipole.
To show a bond is polar, draw the molecule and write δ⁺ and δ⁻ at the two ends of each polar bond, then justify the labels by naming the more electronegative element: oxygen and nitrogen are both more electronegative than carbon and hydrogen. To explain why a molecule with polar bonds has no permanent dipole, you must bring in the shape. In a symmetrical arrangement the bond dipoles are equal in size and point outwards in directions that exactly oppose one another, so the vector sum is zero. Ethanol and ethylamine are polar because their shapes are bent or pyramidal; methoxymethane is bent too, but its smaller electronegativity differences make it less polar than ethanol.
Your focus
- Define electronegativity in terms of the bonding pair in a covalent bond.
- Explain why electronegativity increases from sodium to chlorine but decreases from fluorine to iodine.
- Rank a set of atoms by electronegativity using supplied data and justify the order.
Show all 12 objectives
- Deduce which atom in a given bond carries the partial negative charge and justify the choice with electronegativity data.
- Explain why a partial charge is written as δ⁺ or δ⁻ rather than as a whole charge.
- Identify which bonds in a molecule are polar and which are non-polar.
- Predict whether a molecule has a permanent dipole by combining its bond polarities with its shape.
- Explain why carbon dioxide has polar bonds but no permanent dipole while water has both.
- Identify the structural feature that prevents bond dipoles from cancelling in a named molecule.
- Annotate a displayed structure with δ⁺ and δ⁻ and justify each label using electronegativity.
- Explain why tetrachloromethane has four polar bonds but no permanent dipole.
- Rank several molecules by polarity using both their shapes and the electronegativities of the atoms involved.
Bond polarity exam tips
Marking Points
- one mark for the power or ability of an atom to attract electrons
- one mark for specifying the pair of electrons in a covalent bond rather than electrons in general
- one mark for explaining the trend using increasing nuclear charge and decreasing atomic radius
- one mark for identifying the more electronegative atom correctly from supplied data
- one mark for the bonding pair being drawn closer to the more electronegative atom
- one mark for δ⁻ placed on that atom and δ⁺ on the other
- one mark for identifying which atom is more electronegative, using data where supplied
- one mark for stating that a bond between two identical atoms is non-polar because the distribution is symmetrical
- one mark for identifying the individual polar bonds and their direction
- one mark for considering the shape or symmetry of the whole molecule
- one mark for stating whether the bond dipoles cancel or reinforce
- one mark for the conclusion about whether the molecule has a permanent dipole
- one mark for δ⁺ and δ⁻ written on the correct atoms of each polar bond
- one mark for justifying the labels by naming the more electronegative element
- one mark for drawing or describing the correct shape of the molecule
- one mark for stating that the bond dipoles cancel because the molecule is symmetrical, so there is no permanent dipole
Examiner Tips
- 💡Include the words bonding pair or pair of electrons in a covalent bond in your definition.
- 💡Use the data booklet values rather than memory when a question asks you to compare atoms.
- 💡Explain trends with nuclear charge, radius and shielding, the same three factors as ionisation energy.
- 💡Always write δ⁺ and δ⁻ rather than plain plus and minus signs on a polar bond.
- 💡Justify the labels by naming the more electronegative element, not just by asserting the polarity.
- 💡Check diatomic molecules first: identical atoms always give a non-polar bond.
- 💡Draw the shape before deciding polarity; symmetry cannot be judged from a formula.
- 💡Say explicitly that the dipoles cancel or do not cancel, then state the conclusion.
- 💡Watch for molecules where symmetry is broken by one different atom, such as a chlorinated methane.
- 💡A correctly labelled diagram can earn both the polarity marks and the shape mark at once.
- 💡Name the electronegative atom explicitly; diagrams alone without justification often lose a mark.
- 💡When ranking polarity, compare electronegativity differences, not just the presence of polar bonds.
Common Mistakes
- defining it as the attraction of an atom for all electrons rather than for a shared pair
- applying the term to ions instead of to atoms joined by a covalent bond
- claiming electronegativity increases down a group
- confusing electronegativity with ionisation energy or electron affinity
- drawing full positive and negative charges instead of partial charges
- putting δ⁻ on the less electronegative atom
- describing the process as electron transfer, which would produce ions
- assuming any bond between two different elements must be fully ionic
- deciding molecular polarity from the bonds alone without considering the shape
- saying carbon dioxide is polar because each carbon to oxygen bond is polar
- treating a molecule with lone pairs on the central atom as symmetrical
- confusing a permanent dipole with the temporary induced dipole of a London force
- labelling the molecule as a whole rather than the atoms at the ends of each bond
- using full plus and minus signs instead of δ⁺ and δ⁻
- claiming a molecule is non-polar because its bonds are non-polar when the bonds are in fact polar
- forgetting that a lone pair on the central atom destroys the symmetry that would allow cancellation