Doc Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK KS5 A/AS GCE IB advanced level organic chemistry students US K12 grade 11 grade 12 organic chemistry

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Part 6.2 The physical properties of aldehydes and ketones - intermolecular forces, physical state, boiling points, solubility and odour!

Sub-index for this page

5.2.1 The boiling points of aldehydes and ketones

5.2.2 The solubility of aldehydes and ketones

5.2.3 The odours of aldehydes and ketones

5.2.1 The boiling points of aldehydes and ketones

Structure reminder

Abbreviations: bpt = boiling point 0^oC), na = not applicable, and get used to:

CO representing C=O in abbreviated carbonyl compound formulae for ketones

and CHO representing H-C=O in abbreviated carbonyl compound formulae for aldehydes.

and note butanone is a simpler and unambiguous name for butan-2-one, but you need the latter if there is a substituent atom or group on one of the non carbonyl carbon atoms in the chain.

 

A comparison of isomeric aldehydes and ketones  (all their full structures are on a separate page)

You can see from the table that, for the same carbon chain number, the boiling points of aldehydes and ketones are quite similar and both follow a steady trend of increasing boiling point with increase in carbon length.

Note the boiling point of methylpropanal is 62, lower than its structural isomer butanal because it is a more compact molecule, which reduces the surface to surface intermolecular contact forces - the instantaneous dipole - induced dipole forces.

Similarly the more compact methylbutanone boils at 94^oC, a bit lower than pentan-2-one .

Simple aliphatic aldehydes and ketones are colourless gases, liquids or white solids for high members in the series.

Pure methanal is a gas at room temperature, ethanal is a very volatile liquid, the rest of the lower members are liquids.

All lower members of the ketone homologous series are liquids.

They are all polar molecules because of the presence of the permanent polar bond, the ^δ+C=O^δ- dipole.

and extra δ+ δ- attractions

You can see that the polar bond adds a significant extra permanent dipole - permanent dipole attraction to the total intermolecular bonding force.

The boiling point trend of linear aldehydes are now discussed in detail and compared with other homologous series.

Graph 1 grey-purple line = aldehydes

The red line graph shows the boiling point of alkanes from methane CH₄ (boiling point -164^oC/109 K)  to tetradecane C₁₄H₃₀ (boiling point 254^oC/527 K). [Remember K = ^oC + 273]

Note: The red line represents linear alkanes in all the graphs 1-3 and is a useful baseline to compare the intermolecular bonding present in other homologous series of non-cyclic aliphatic compounds.

For the 'purple line of linear aldehydes', the graph goes from methanal HCHO to decanal CH₃(CH₂)₈CHO.

It is a similar graph line for methyl ketones (2-ones, RCOCH₃, R = alkyl), but I haven't plotted the data (yet!).

So, in this discussion we are comparing the red line (linear alkanes) with the purple line (linear aldehydes) AND comparing molecules with the same number of electrons.

A plot of number of electrons in any molecule of a homologous series versus its boiling point (K) shows a steady rise with a gradually decreasing gradient.

I consider this the best for comparison of the effects of intermolecular bonding between different functional groups.

REMINDER: Intermolecular forces are all about partially positive (δ+) sites and partially negative (δ–) sites on molecules causing the attraction between neighbouring molecules - though their origin can differ.

I think Graph 1 is the best graph to look at the relative effects on intermolecular forces (intermolecular bonding) on boiling point because it is the distortion of the electron clouds (e.g. in non-polar alkanes), that gives rise to these, weak, but not insignificant forces, known as instantaneous dipole - induced dipole forces.

From Graph 1 you can see the effect of the permanently polar oxygen - hydrogen bond (C^δ+=O^δ-) increases the intermolecular forces of attraction, and raising the boiling point compared to non-polar molecules of similar size in terms of numbers of electrons (clouds).

The bond is highly polar because of the great difference in electronegativity (carbon 2.5, oxygen 3.5).

Even so, for most polar molecules, the majority of the intermolecular force is still due to the instantaneous dipole - induced dipole attractions.

Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

For propanone (14.2%  instantaneous  dipole – induced dipole force) + (78.4% permanent dipole – permanent dipole) + (7.4% permanent dipole – induced dipole)

The increase in intermolecular attractive forces, means the molecules need a higher kinetic energy to escape from the liquid surface i.e. have a higher boiling point.

However, because aldehydes and ketones cannot hydrogen bond with themselves, their boiling points are still less than hydrogen bonded alcohols - check out yellow line for linear primary alcohols on Graph 1 above..

 

Graph 2 grey-purple line = aldehydes

A plot of the molecular mass of the linear aldehyde molecules versus its boiling point (K) shows a steady rise with a gradually decreasing gradient.

More atoms, more electron clouds, more chance of instantaneous dipole - induced dipole forces, so the overall intermolecular force steadily increases with carbon number, the hydrogen bonding is a fairly constant contribution.

For a similar molecular mass, the aldehydes have significantly higher boiling points than alkanes, mainly due to the extra force of the permanent dipole - permanent dipole attraction.

 

Graph 3 grey-purple line = aldehydes

A plot of the carbon number of the linear aldehyde molecule versus its boiling point (K) shows a steady rise with a gradually decreasing gradient.

For the same carbon number, the aldehydes have significantly higher boiling points than alkanes, mainly due to the extra force of the permanent dipole - permanent dipole attraction

The increase in intermolecular attractive forces, means the molecules need a higher kinetic energy to escape from the liquid surface i.e. have a higher boiling point.

TOP OF PAGE and sub-index

5.2.2 The solubility of aldehydes and ketones

How soluble in water?

Aldehydes and ketones are highly polar molecule and readily bond with water molecules via permanent dipole - permanent dipole intermolecular forces.

(aldehyde/ketone) ^δ+C=O^δ- llll ^δ+H-O^δ- (water) hydrogen bond (llll)

Note that aldehydes and ketones do not hydrogen bond with themselves, but they can hydrogen bond with water.

The hydrogen bonding with water enables the lower aldehydes/ketones to dissolve in water, but after that, the solubility of them rapidly decreases.

So we need to consider solvent - solvent, solute - solute and solute - solute interactions in terms of intermolecular bonding attractive forces to explain this trend.

An increase in the 'hydrocarbon' chain length makes the carbonyl compound less and less able to disrupt hydrogen bonding - the longer the hydrocarbon chain, the more water - water hydrogen bonds must be disrupted to dissolve the alcohol, without compensating aldehyde/ketone - water hydrogen bonds.

You can also argue that the instantaneous dipole - induced dipole forces between the hydrocarbon chain of neighbouring aldehyde/ketone molecules is stronger than a hydrogen bond, so the longer chain molecules will come together.

As the carbon chain gets longer, the solubility decreases, because the intermolecular bonding described above, becomes less compensated for as more water - water hydrogen bonds are disrupted.

Consequently, methanal, ethanal and propanone are very soluble.

Propanal and butanone are quite soluble.

Butanal, pentan-2-one, pentan-3-one are only slightly soluble as the hydrophobic alkyl group increases in size and 'attempting' to disrupt more the hydrogen bonds between water molecules.

 

Use as solvents themselves

Aldehydes are not used as solvents - unpleasant and can be harmful if inhaled.

The lower ketones are like propanone and butanone are very useful solvents because they dissolve a wide variety of organic compounds.

TOP OF PAGE and sub-index

5.2.3 The odours of aldehydes and ketones!

When first studying organic chemistry at an advanced level, the odour of the compounds is rarely mentioned apart from the pungent 'hydrocarbon' smell of alkanes, alkenes and aromatic hydrocarbons.

Aldehydes and ketones also have strong smelling odours e.g.

Ethanal, like other lower aldehydes, has very strong pungent unpleasant odour, it is also harmful to breathe in - a known carcinogen - and note it is produced in one of the metabolic pathways after drinking an alcoholic beverage!

Propanone and other lower ketones have  moderately pleasant 'sweeter' odour than aldehydes (like 'acetone' nail varnish remover)

The aromatic aldehyde orbenzaldehyde, C₇H₆O, has the smell of almonds and used in flavouring food - I love thick layer of marzipan between a thin slice of sugar icing and a thick layer of rich Christmas cake!

Many more complex molecules containing an aldehyde group or ketone group occur in essential oils from plants, often giving, or contributing to, their characteristic odour e.g. citral and carvone whose skeletal formulae are shown below.

A Citral (lemonal), C₁₀H₁₆O, is an aldehyde with all the usual characteristic reactions of an R-CHO molecule.

Citral also has two alkene groups (C=C) as well as the aldehyde group.

It is found in several species of lemon plants (e.g. lemon grass oil) and contributes to the strong citrus 'lemon-like' odour of the fruit.

Citral can exhibit E/Z stereoisomerism via the top C=C bond (geometric isomers) which is not part of a ring.

Citral cannot exhibit R/S stereoisomerism - no chiral carbon.

Note 'al' in the name 'citral'.

B Carvone, C₁₀H₁₄O, is a cyclic ketone found in spearmint and caraway oil.

It has one ketone group and behaves as a ketone, but it also two C=C double bond groups, so is 'diene'.

Both citral and carvone behaves as a 'double' alkene, reacting quantitatively with two molecules of bromine.

They are also both a good test in reading skeletal formulae - match them up with the molecular formulae!

Carvone cannot exhibit E/Z stereoisomerism via the top C=C bond (geometric isomers) - one C=C has two identical end groups and the top left C=C is part of the ring.

Citral can exhibit R/S stereoisomerism - it has a chiral carbon - the bottom one of the hexagonal ring - see diagram on right for the non-superimposable mirror image forms.

Note 'one' in the name 'carvone'.

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