Topic 7 · page 4 of 14

Permanent dipole–permanent dipole interactions

A molecule with polar bonds that do not cancel has a permanent dipole. The δ+ end of one molecule attracts the δ− end of the next. This force acts in addition to London forces.

12 min read7.17.1(ii)

Read first: London forces

Polar molecules 7.17.1(ii)

Atoms of different elements pull on a shared pair of electrons by different amounts. Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. If the two atoms differ, the electrons sit closer to the more electronegative atom. That atom is slightly negative (δ−) and the other is slightly positive (δ+). The bond is polar and has a dipole.

In hydrogen chloride, Cl is more electronegative than H, so H is δ+ and Cl is δ−.

A molecule has a permanent dipole only if the bond dipoles do not cancel. Think of each bond dipole as a pull towards the δ− atom, and add the pulls. Water is bent, so its two O–H pulls add up to a net dipole pointing towards oxygen. Carbon dioxide is linear, so its two C=O pulls point in opposite directions and cancel.

Three molecules with their dipoles. Hydrogen chloride: a polar bond with delta plus on H and delta minus on Cl, so the molecule has a net dipole. Water: two polar O–H bonds in a bent shape whose dipoles add to a net dipole pointing towards oxygen. Carbon dioxide: two polar C=O bonds in a straight line, whose dipoles point opposite ways and cancel, so there is no net dipole.
HCl and H₂O have a net dipole. CO₂ has polar bonds but no net dipole.

MethodIs this molecule polar?

  1. Draw the shape of the molecule (from the number of bonding and lone pairs, Topic 3).
  2. Mark each polar bond: the atom that is more electronegative is δ−.
  3. Ask whether the pulls cancel. A symmetrical shape with identical outer atoms cancels: CO₂ (linear), BF₃ (trigonal planar), CCl₄ (tetrahedral). Otherwise the molecule has a net dipole: HCl, H₂O, CH₃Cl.

The attraction between polar molecules 7.17.1(ii)

Polar molecules turn so that the δ+ end of one faces the δ− end of the next. The oppositely charged ends attract. This is the permanent dipole–permanent dipole interaction.

Two arrangements of hydrogen chloride molecules. Left: two HCl molecules end to end, with the delta minus chlorine of the first facing the delta plus hydrogen of the second, joined by a dashed line. Right: two HCl molecules side by side pointing opposite ways, with dashed lines joining each delta plus hydrogen to the delta minus chlorine of the other molecule.
In both arrangements, every dashed attraction joins a δ+ end to a δ− end.

The word permanent is what separates this from London forces. The dipole is always there, so it does not need the electrons to move. But London forces are still present in a polar molecule, because every molecule has them. A polar molecule has both forces.

The effect shows up when you compare molecules with a similar number of electrons, so that their London forces are about equal. Butane, C₄H₁₀, is non-polar and has 34 electrons. Propanone, CH₃COCH₃, has a polar C=O bond and 32 electrons: 3 × 6 + 6 + 8 = 32. Butane has slightly more electrons, but propanone boils much higher.

Butane Propanone
Electrons 34 32
Polar? no yes (C=O)
Forces London only London + permanent dipole–dipole
Boiling temperature −1 °C 56 °C

The extra 57 °C for propanone comes from the permanent dipole–dipole interactions, which butane does not have.

Check yourself

  1. Define a permanent dipole–permanent dipole interaction.
  2. State which of these molecules have a net dipole: HBr, CO₂, CH₃Cl, CCl₄, BF₃, H₂O.
  3. CO₂ contains polar bonds. Explain why its molecules have no permanent dipole–dipole interactions.
  4. Name the intermolecular forces in liquid HCl.
  5. Propanone (56 °C) boils at a higher temperature than butane (−1 °C), although butane has more electrons. Explain.
  6. Which force is present in every molecule in question 2, including those with no net dipole?

Answers cover these until you have tried the questions

  1. An attraction between the δ+ end of one polar molecule and the δ− end of a neighbouring polar molecule. The dipoles are permanent.
  2. Net dipole: HBr (polar bond, two atoms), CH₃Cl (the C–Cl dipole is not cancelled by three C–H bonds) and H₂O (bent). No net dipole: CO₂ (linear), CCl₄ (tetrahedral) and BF₃ (trigonal planar), because their dipoles cancel.
  3. The two C=O bond dipoles point in opposite directions in a linear molecule, so they cancel. The molecule has no net dipole, so it has only London forces.
  4. Permanent dipole–permanent dipole interactions and London forces.
  5. Propanone has a polar C=O bond that does not cancel, so it has permanent dipole–dipole interactions as well as London forces. Butane is non-polar and has London forces only. Butane's 34 electrons give London forces that are only slightly stronger than propanone's (32 electrons), so the extra dipole–dipole interactions in propanone outweigh them. More energy is needed to separate propanone molecules.
  6. London forces.