Topic 7 · page 14 of 14
Putting the forces together
One table of the three forces, one method for boiling temperatures and one for solubility. Use this page to practise mixed questions, where you must choose which force and which argument to use.
Read first: The hydrogen halides, Non-aqueous solvents
The three forces side by side 7.1
| London forces | Permanent dipole–dipole | Hydrogen bonds | |
|---|---|---|---|
| Cause | instantaneous dipole induces a dipole in a neighbour | δ+ end of one polar molecule attracts δ− end of the next | δ+ H on N, O or F attracts a lone pair on N, O or F of another molecule |
| Present in | every molecule | molecules with a net dipole | molecules with an N–H, O–H or H–F bond |
| Strength | weakest; rises with electrons and surface contact | stronger than London forces at similar electron counts | strongest of the three |
| Evidence | alkane chain length, branching, HCl to HI | propanone against butane | water, alcohols, HF, ice, solubility |
Forces add. Water has London forces and hydrogen bonds. HCl has London forces and permanent dipole–dipole. Hexane has London forces only.
Boiling temperatures: choose the argument 7.5
When you meet a boiling-temperature question, ask which of these is the difference between the two substances, and argue only from that.
| The two substances differ in... | So the argument is... |
|---|---|
| number of electrons (same type of force) | more electrons: stronger London forces |
| shape (isomers) | more branching: less surface contact: weaker London forces |
| presence of a net dipole | permanent dipole–dipole interactions as well as London forces |
| an O–H, N–H or H–F bond | hydrogen bonds as well as London forces |
Worked exampleRank methane, hydrogen sulfide and water
Boiling temperatures: −162 °C, −60 °C, 100 °C, in some order. Methane, CH₄, has 10 electrons. H₂S has 2 + 16 = 18 electrons. H₂O has 2 + 8 = 10 electrons.
- CH₄ is non-polar: London forces only, and the fewest electrons (10, tied with H₂O). It has the lowest boiling temperature, −162 °C.
- H₂S has more electrons (18) than CH₄, so stronger London forces, and it is polar, so it also has weak permanent dipole–dipole interactions. It has no hydrogen bonds because S is not N, O or F. It boils at the middle temperature, −60 °C.
- H₂O has the same number of electrons as CH₄ (10), so similar London forces. But it also has hydrogen bonds (O–H, two lone pairs on O), which are stronger than the forces in H₂S. It boils highest, 100 °C.
- Order: CH₄ < H₂S < H₂O. Water is highest although it has the fewest electrons, because its hydrogen bonds need more energy to overcome.
Solubility: broken against formed 7.6
Whenever you are asked why something does or does not dissolve, use one table in your head.
| Broken | Formed | |
|---|---|---|
| Solute | forces between solute particles | |
| Solvent | forces between solvent molecules | |
| New | forces between solute and solvent |
Dissolving is favourable when the formed attractions are about as strong as the broken ones.
| Solute in water | Broken | Formed | Result |
|---|---|---|---|
| sodium chloride | ionic attractions, hydrogen bonds | ion–water attractions (hydration) | dissolves |
| ethanol | hydrogen bonds (ethanol and water) | hydrogen bonds with water | dissolves |
| bromoethane | dipole–dipole and hydrogen bonds | weak attractions only (no hydrogen bonds) | poor |
| hexane | hydrogen bonds | London forces only | does not mix |
Check yourself
- Name the intermolecular forces in each of liquid HI, liquid NH₃ and liquid hexane.
- Water (100 °C) boils at a higher temperature than hydrogen fluoride (20 °C), although both have 10 electrons and both have hydrogen bonds. Explain.
- Ammonia, NH₃, boils at −33 °C and phosphine, PH₃, at −88 °C, although PH₃ has more electrons. Explain.
- Predict whether potassium bromide dissolves in (a) water and (b) hexane. Give a reason for each.
- State two differences between how the molecules are arranged in ice and in liquid water, and say what the difference means for density.
- Ethanol (78 °C), propanone (56 °C) and butane (−1 °C) have 26, 32 and 34 electrons. Explain the order of their boiling temperatures.
Answers cover these until you have tried the questions
- HI: permanent dipole–permanent dipole interactions and London forces. NH₃: hydrogen bonds and London forces. Hexane: London forces only.
- Each water molecule has 2 δ+ H and 2 lone pairs, so it can form up to four hydrogen bonds and, on average, 2 bonds per molecule. HF has 1 δ+ H, so each molecule forms on average only 1 (its three lone pairs cannot all be used). Water has more hydrogen bonds per molecule, so more energy is needed to separate its molecules.
- NH₃ has hydrogen bonds (H bonded to N, with a lone pair on N) as well as London forces. PH₃ has no hydrogen bonds, because P is not N, O or F. The hydrogen bonds in NH₃ outweigh the stronger London forces of PH₃, so more energy is needed to separate NH₃ molecules.
- (a) Dissolves in water: polar water molecules hydrate the ions (δ− O towards K⁺, δ+ H towards Br⁻), and the energy released makes up for the energy needed to break the lattice. (b) Does not dissolve in hexane: hexane has only London forces, so it cannot hydrate the ions, and the weak attractions formed do not make up for the strong ionic attractions broken.
- In ice, each molecule is held by four hydrogen bonds in a fixed open lattice with empty space; in liquid water, some hydrogen bonds are broken and the molecules are closer together. So water has more molecules per unit volume and is denser than ice.
- Butane < propanone < ethanol. The electron counts are similar (26 to 34), so compare the type of force. Butane is non-polar: London forces only (−1 °C). Propanone is polar: permanent dipole–dipole interactions and London forces (56 °C). Ethanol has an O–H bond: hydrogen bonds and London forces (78 °C). The stronger the type of force, the more energy is needed to separate the molecules.