Topic 7: Intermolecular Forces

Learning objectives

7.1The nature of the following intermolecular forces: i London forces (instantaneous dipole-induced dipole) ii permanentUnderstand the nature of the following intermolecular forces: i London forces (instantaneous dipole-induced dipole) ii permanent dipole-permanent dipole interactions iii hydrogen bonds 7.2The interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bondingUnderstand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding 7.3The following anomalous properties of water resulting from hydrogen bonding: i its high melting and boiling temperature whenUnderstand the following anomalous properties of water resulting from hydrogen bonding: i its high melting and boiling temperature when compared with similar molecules ii the density of ice compared to that of water 7.4Predict the presence of hydrogen bonding in molecules analogous to those mentioned in 7.2Be able to predict the presence of hydrogen bonding in molecules analogous to those mentioned in 7.2 7.5Understand, in terms of intermolecular forces, physical properties shown by substancesUnderstand, in terms of intermolecular forces, physical properties shown by substances, including: i the trends in boiling temperatures of alkanes with increasing chain length ii the effect of branching in the carbon chain on the boiling temperatures of alkanes iii the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons iv the trends in boiling temperatures of the hydrogen halides HF to HI 7.6Factors that influence the choice of solventsUnderstand factors that influence the choice of solvents, including: i water, to dissolve some ionic compounds, in terms of the hydration of the ions ii water, to dissolve simple alcohols, in terms of hydrogen bonding iii water, as a poor solvent for compounds (to include polar molecules such as halogenoalkane), in terms of inability to form hydrogen bonds iv non-aqueous solvents, for compounds that have similar intermolecular forces to those in the solvent

Force	How it forms	Where it occurs
London force	a momentary uneven electron distribution creates an instantaneous dipole, which induces an opposite dipole in a neighbour; the dipoles attract	between all atoms and molecules
permanent dipole–permanent dipole	the $\delta+$ end of one polar molecule attracts the $\delta-$ end of another	between polar molecules
hydrogen bond	a lone pair on N, O or F attracts a strongly $\delta+$ H covalently bonded to N, O or F in another molecule	molecules with a suitable donor and acceptor

London forces strengthen as electron clouds become larger and more polarisable, and as molecular shapes allow more surface contact. They are present even when stronger named interactions also occur.

For comparable small molecules, hydrogen bonding is usually the strongest of these interactions, then permanent-dipole attraction, then London forces; real comparisons must also consider how many contacts and electrons are present.

Intermolecular forces act between particles. Boiling a simple molecular substance overcomes these attractions; it does not normally break the covalent bonds inside each molecule.

In a hydrogen bond, show N, O or F with a lone pair as the acceptor, a dotted line to H $^{\delta+}$ on a neighbouring N–H, O–H or F–H bond, and the donor atom as $\delta-$ .

Liquid	Donor sites	Acceptor feature	Network consequence
H $_2$ O	two O–H hydrogens	two lone pairs on O	each molecule can participate in an extensive network, up to four hydrogen bonds in the ideal arrangement
NH $_3$	three N–H hydrogens	one lone pair on N	hydrogen bonds form, but acceptor availability limits the network compared with water
HF	one H–F hydrogen	lone pairs on F	molecules associate through H–F···H–F chains/networks

The highly electronegative N, O or F atom polarises the X–H bond. The exposed H $^{\delta+}$ can approach a lone pair on another molecule closely, producing a particularly strong intermolecular attraction.

The covalent O–H, N–H or H–F bond is inside a molecule; the dotted hydrogen bond joins different molecules. Do not draw the hydrogen bond to another hydrogen atom.

Water property	Hydrogen-bond explanation
unusually high melting and boiling temperatures	many hydrogen bonds between water molecules must be overcome, requiring more energy than for similar-sized molecules without the same network
ice less dense than liquid water	ice has an open, ordered tetrahedral hydrogen-bond lattice that holds molecules farther apart; on melting, part of the network collapses and molecules pack closer

Water and ammonia have similar electron counts, so their London forces are comparable. Water can form a more extensive hydrogen-bond network, helping explain its much higher boiling temperature.

For the same mass, the larger volume of the open ice structure gives a lower density. Ice therefore floats on liquid water.

Ice is not less dense because its water molecules become lighter. Molecular mass is unchanged; average spacing and therefore volume change.

For hydrogen bonding between identical molecules, check for both: (1) H directly bonded to N, O or F, which is a donor site; and (2) an available lone pair on N, O or F, which is an acceptor site.

Molecule	Self hydrogen bonding?	Reason
CH $_3$ OH	yes	O–H donor and O lone pairs
CH $_3$ NH $_2$	yes	N–H donor and N lone pair
CH $_3$ OCH $_3$	no between identical molecules	O accepts, but there is no O–H/N–H/F–H donor
CH $_3$ F	no	C–H is not a qualifying donor even though F has lone pairs
CH $_3$ CH $_2$ OH with water	yes	both species provide O–H donors and O acceptors

Draw the full local structure, mark $\delta+$ on the qualifying H, mark the acceptor lone pair, then place the dotted interaction from that lone pair to H. Use the actual bonding, not merely the molecular formula.

A molecule containing N, O or F does not automatically hydrogen-bond to itself. It also needs H directly bonded to one of those atoms.

Comparison	Observed trend	Molecular explanation
straight-chain alkanes as chain length increases	boiling temperature rises	more electrons and a larger, more polarisable contact surface strengthen London forces
branched vs less-branched alkane isomers	more branching usually lowers boiling temperature	compact shapes have less surface contact, weakening total London attraction
alcohol vs alkane with similar electron count	alcohol has higher boiling temperature and lower volatility	alcohol molecules hydrogen-bond; alkane molecules have London forces only
HCl to HBr to HI	boiling temperature rises down the group	larger, more polarisable electron clouds strengthen London forces
HF compared with HCl, HBr and HI	HF is anomalously high	HF forms hydrogen bonds; the other hydrogen halides do not

First list every force present, then compare the factor that changes—electron number/polarisability, contact surface or hydrogen-bond capability. More energy needed to separate molecules means a higher boiling temperature and lower volatility.

Do not say 'larger molecules have stronger bonds' when explaining boiling. The covalent bonds remain intact; it is the total intermolecular attraction that changes.

Solute and solvent	New interactions on mixing	Solubility reasoning
some ionic compounds in water	ion–dipole hydration: O $^{\delta-}$ points towards cations and H $^{\delta+}$ towards anions	hydration can compensate for separating lattice ions and water molecules
simple alcohol in water	hydrogen bonds form between alcohol –OH and water	favourable new hydrogen bonds give good solubility; a longer hydrocarbon chain reduces it
halogenoalkane in water	it may be polar but cannot form enough strong hydrogen bonds with water	new attractions do not compensate well for disrupting water's hydrogen-bond network, so solubility is low
non-polar solute in a non-aqueous non-polar solvent	both rely mainly on London forces	similar intermolecular forces make mixing more favourable

Dissolving requires separation of some solute particles and solvent molecules, followed by formation of solute–solvent attractions. 'Like dissolves like' is a summary of that energetic competition, not a replacement for naming the forces.

Not every ionic compound is soluble: the balance between lattice attraction and hydration differs. For dissolved ions, describe the orientation of water dipoles around each charge.

Polarity alone does not guarantee water solubility. A polar halogenoalkane can still be poorly soluble because it cannot replace the strong water–water hydrogen bonds effectively.

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Topic 7: Intermolecular Forces

Learning objectives

Three intermolecular forces arise from charge separation

H2O, NH3 and HF form hydrogen-bond networks

Hydrogen bonding makes water thermally and structurally anomalous

Predict hydrogen bonding from donor and acceptor sites

Intermolecular forces explain boiling and volatility trends

A solvent works when new attractions compensate for old ones

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