Topic 7

Topic 7 on a page

Everything from this topic in one place, for the last read before an exam. If anything here is unfamiliar, go back to the full page.

Definitions to learn word for word

Intermolecular force: a force of attraction between one molecule and another. Boiling and melting overcome it; the covalent bonds inside the molecules stay (Forces between molecules).
London forces: attractions between an instantaneous dipole in one molecule and an induced dipole in a neighbouring molecule (London forces).
Permanent dipole–permanent dipole interaction: an attraction between the δ+ end of one polar molecule and the δ− end of a neighbouring polar molecule (Permanent dipole–dipole).
Hydrogen bond: an attraction between a δ+ hydrogen atom bonded to N, O or F and a lone pair on a N, O or F atom of a different molecule (Hydrogen bonds).
Volatility: how easily a liquid evaporates. Higher boiling temperature means lower volatility (Alcohols against alkanes).
Hydration of ions: the surrounding of ions by water molecules, δ− O towards positive ions and δ+ H towards negative ions (Ionic compounds).

The three forces

	London forces	Permanent dipole–dipole	Hydrogen bonds
Present in	every molecule	molecules with a net dipole	molecules with N–H, O–H or H–F
Strength	weakest; rises with electrons and surface contact	stronger than London at similar electrons	strongest
Examples	alkanes, I₂, noble gases	HCl, propanone, CH₃Cl	H₂O, NH₃, HF, alcohols

The forces add: a polar molecule has dipole–dipole and London forces; a molecule with hydrogen bonds has them and London forces. Count electrons first: a molecule with many more electrons can beat a stronger type of force (octane 126 °C against ethanol 78 °C).

Boiling temperature data

Substance	Electrons	Boiling temperature / °C	Forces
methane to octane	10, 18, 26, 34, 42, 50, 58, 66	−162, −89, −42, −1, 36, 69, 98, 126	London
pentane, 2-methylbutane, 2,2-dimethylpropane	42 each	36, 28, 10	London; less surface contact with more branching
ethane, methanol	18 each	−89, 65	London; London + hydrogen bonds
propane, ethanol	26 each	−42, 78	London; London + hydrogen bonds
butane, propan-1-ol	34 each	−1, 97	London; London + hydrogen bonds
HF, HCl, HBr, HI	10, 18, 36, 54	20, −85, −67, −35	hydrogen bonds; dipole–dipole
H₂O, H₂S, H₂Se, H₂Te	10, 18, 36, 54	100, −60, −41, −2	hydrogen bonds; weaker forces
butane, propanone	34, 32	−1, 56	London; London + dipole–dipole
ethanol, methoxymethane	26 each	78, −24	hydrogen bonds; dipole–dipole

Number of electrons in $\text{C}_n\text{H}_{2n+2}$ : 8n + 2. Water's expected boiling temperature from the trend: about −80 °C; real: 100 °C, about 180 °C higher (page 7).

Explaining boiling temperatures

Name the forces in each substance (London always; dipole–dipole if net dipole; hydrogen bonds if H on N, O or F).
Similar electrons: the stronger type of force boils higher. Same type: more electrons boils higher. Isomers: less surface contact (more branching) boils lower.
More energy is needed to overcome the stronger forces between molecules, so the boiling temperature is higher.

Water and ice

Water boils at 100 °C because each molecule has 2 δ+ H and 2 lone pairs, so it forms up to four hydrogen bonds, the most of any small hydride. H₂S has only London and weak dipole–dipole forces (page 7).
Ice (0.917 g cm⁻³) is less dense than water (1.00 g cm⁻³): hydrogen bonds hold the molecules in an open lattice with empty space; on melting, some bonds break and the molecules move closer (page 8). Water expands by about 9 % on freezing: 1.00 ÷ 0.917 = 1.09.

Drawing a hydrogen bond

Show lone pairs on the N, O or F, δ+ on H, δ− on N, O or F, and a dashed line from H to the lone pair of the other molecule, with X–H···X straight. Label it "hydrogen bond". Predict hydrogen bonds from the structure: an H bonded to N, O or F, plus a lone pair on N, O or F (page 6).

Solvents

Solute in solvent	Broken	Formed	Result
ionic solid in water	ionic lattice, some hydrogen bonds	ion–water attractions (hydration)	dissolves if hydration makes up for the lattice
simple alcohol in water	hydrogen bonds in each	hydrogen bonds with water	dissolves
halogenoalkane in water	hydrogen bonds between water molecules	weak attractions only	poor solvent: no hydrogen bonds with water
London-force solute in hexane	London forces	London forces	dissolves: similar forces
ionic solid in hexane	ionic lattice	nothing as strong	insoluble

The ten traps

Saying covalent bonds break when a substance boils.
Saying London forces only exist in non-polar molecules, or that their dipoles are permanent.
Forgetting that a molecule needs a net dipole (CO₂ has none).
Giving hydrogen bonds to HCl, H₂S or anything with only C–H bonds, or to a molecule with N, O or F but no H on it (CH₃F, propanone).
Drawing the hydrogen bond to an H atom instead of a lone pair, or leaving out the lone pair, the partial charges or the straight line.
Explaining branching with electrons or bond strength instead of surface contact.
Comparing an alcohol with an alkane without saying they have similar electron counts.
Explaining ice floating by stronger hydrogen bonds or expanding molecules instead of the open lattice.
Saying "like dissolves like" with no named forces, or "polar, so it dissolves in water" for a halogenoalkane.
Saying hydration happens to the water, not the ions; or putting the wrong end of water towards an ion.

Mixed test

Name all the intermolecular forces in liquid HBr, liquid CH₃OH and liquid C₆H₁₄.
Explain why hexane (69 °C) boils at a higher temperature than butane (−1 °C).
Explain why 2,2-dimethylpropane boils at a lower temperature than pentane, although both are C₅H₁₂.
Draw two molecules of hydrogen fluoride with a hydrogen bond between them.
Explain why ethanol boils at a higher temperature than propane.
Explain why HF boils at a higher temperature than HCl, and why HI boils at a higher temperature than HCl.
Explain why ice floats on water, and calculate the volume of ice formed when 360 g of water freezes (ice 0.917 g cm⁻³).
Explain why sodium chloride dissolves in water but not in hexane.
Explain why ethanol dissolves in water but bromoethane does not.
Predict, with a reason, whether I₂ dissolves in hexane.

Answers cover these until you have tried the questions

HBr: permanent dipole–permanent dipole interactions and London forces. CH₃OH: hydrogen bonds and London forces. C₆H₁₄: London forces only.
Hexane has 50 electrons against 34 in butane, so stronger London forces. More energy is needed to separate hexane molecules.
Both have 42 electrons. 2,2-Dimethylpropane is branched and compact, so it has less surface contact and weaker London forces. Less energy is needed to separate its molecules.
Two HF molecules, each F with three lone pairs, δ+ on H and δ− on F, and a dashed line from the H of one molecule straight to a lone pair on the F of the other, with F–H···F in a straight line, labelled "hydrogen bond".
Both have 26 electrons, so similar London forces. Propane has only London forces. Ethanol has an O–H group and forms hydrogen bonds, so more energy is needed to separate its molecules.
HF has hydrogen bonds (H on F with lone pairs on F), stronger than the dipole–dipole forces and London forces in HCl, which has no hydrogen bonds. HI has 54 electrons against 18 in HCl, so stronger London forces, which outweigh the weaker dipole–dipole forces in HI.
Ice is less dense (0.917 against 1.00 g cm⁻³): in ice, hydrogen bonds hold the molecules in an open lattice with empty space. When it melts some hydrogen bonds break and the molecules pack closer together. Volume of ice = 360 ÷ 0.917 = 392.6 cm³ (393 cm³), against 360 cm³ for the water, an increase of 32.6 cm³, or 9.05 %.
Water is polar: its δ− O attracts Na⁺ and its δ+ H attracts Cl⁻, so the ions are hydrated and the energy released makes up for breaking the lattice. Hexane has only London forces and cannot hydrate the ions, so nothing makes up for the strong ionic attractions broken.
Ethanol forms hydrogen bonds with water (O–H and lone pairs on O), about as strong as the hydrogen bonds broken. Bromoethane has no O–H, N–H or H–F bond, so it cannot form hydrogen bonds with water, and the hydrogen bonds between water molecules are not replaced by attractions of similar strength.
Yes, it dissolves. Iodine has London forces only, as does hexane, so the attractions formed with hexane are about as strong as the London forces broken.