Part 6. The Chemistry of  Carboxylic Acids and their Derivatives

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

Part 6.2 The physical properties of carboxylic acids - odours, melting points, boiling points and solubility

Sub-index for this page

6.2.1 Examples of the melting/boiling points, solubility and smell of carboxylic acids - general trends discussed and explained

6.2.2 More details on the boiling point trend of linear monocarboxylic acids

6.2.3 More details on the solubility of carboxylic acids

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6.2.1 DATA TABLE: Examples of the melting points, boiling points and solubility of carboxylic acids

Abbreviations used: mpt = melting point;  bpt = boiling point; dec = thermally decomposes

Physical state at room temperature

Monocarboxylic acids are colourless liquids or white waxy solids.

All aromatic carboxylic acids are crystalline solids at room temperature.

The odour of carboxylic acids

Aromatic carboxylic acids have relatively faint odours, but aliphatic monocarboxylic acids all have strong odours - think of the dilute aqueous solution of ethanoic acid, known as vinegar!

Butter turns and smells rancid because butanoic acid is formed by the action of bacteria on butter fat.

Notes on the above data table

(a) Names

The first four monocarboxylic acids have retained their trivial names, but after that the name uses pent, hex, hept ... anoic acid for the systematic name.

For more details see Molecular structure and nomenclature of carboxylic acids and derivatives

(b) Structure

For more details see Molecular structure and nomenclature of carboxylic acids and derivatives

(c) Melting points

No obvious pattern with the linear monocarboxylic acids, you don't get a systematic trend as you do with boiling points i.e. a steady increase with increase in electrons in the molecule with increase in carbon chain length.

The melting points of linear aliphatic monocarboxylic acids

One reason for this lack of obvious trend is the ability of carboxylic acids to form 'dimers' via hydrogen bonding between two carboxylic acid group - illustrated in the diagram above.

This is especially so with the lower members of the aliphatic monocarboxylic acids like methanoic acid, ethanoic acid.

It is not a true dimer in the sense that the two molecules are covalently bonded together, but the hydrogen bonding is strong enough to hold them together, not only in the crystalline state, but when liquid or dissolved in most solvents other than water.

This effectively doubles the size of the molecules and roughly doubles the instantaneous dipole - induced dipole intermolecular forces between the neighbouring dimers.

A greater kinetic energy is needed to melt the 'dimer' compared to the 'monomer', hence the higher than expected melting points for the lower members of the series.

However, as the hydrocarbon chain length increases, the hydrogen bonding between he carboxylic acid groups is increasing hindered.

Decanoic acid is the first member of this homologous series to be a solid at room temperature.

It is a white solid with a strong rancid odour - they all have strong odours!

Think of the smell of pure ethanoic acid - 100% concentration of vinegar!

Higher members do follow a more systematic trend of increasing melting point as the instantaneous dipole - induced dipole forces dominate much more than the hydrogen bonding.

The melting points of aromatic carboxylic acids.

Benzoic acid, (64 electrons), mpt 122^oC,  and propylbenzene (66 electrons), mpt -100^oC.

Aromatic acids also form hydrogen bonded dimers, which raises the melting point e.g. comparing the melting points of two aromatic molecules with similar numbers of electrons in the molecule.

The benzoic acid 'dimer' has a much higher melting point than the aromatic hydrocarbon propyl benzene.

Isomer differences

(i) CH₃CH₂CH₂COOH, butanoic acid melts at -6^oC and boils at 164^oC.

(ii) Isomeric (CH₃)₂CHCOOH, 2-methylpropanoic acid melts at -47^oC and boils at 154^oC.

(ii) is a more compact molecule with less surface to surface contact with neighbouring molecules, so the instantaneous dipole - induce dipole forces are reduced - lower kinetic energies and enthalpies required to effect the change of state.

 

(d) Boiling points of linear monocarboxylic acids

The boiling points of aromatic acids.

The boiling points show a steady trend of increase with increase in length of the carbon chain of the molecule (and increase in total electrons).

The boiling points of linear monocarboxylic acids are discussed in detail in 6.2.2

The boiling points of aromatic acids.

Benzoic acid, (64 electrons), mpt 122^oC,  and propylbenzene (66 electrons), mpt -100^oC.

Aromatic acids also form hydrogen bonded dimers, which raises the melting point e.g. comparing the melting points of two aromatic molecules with similar numbers of electrons in the molecule.

The benzoic acid 'dimer' has a much higher melting point than propyl benzene.

 

(e) Solubility in water

For the linear monocarboxylic linear aliphatic acids you get a significant decrease in solubility as the hydrophobic carbon chain gets longer.

However, the solubility of lower members is enhanced by hydrogen bonding with water.

The solubility of carboxylic acids is discussed in detail in 6.2.3

(f) Electrons in molecule

Equals the sum of the atomic numbers in the molecular formula.

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6.2.2 The boiling point trend of linear monocarboxylic acids and intermolecular forces

Lower in this homologous series are colourless liquids.

Higher members of the series are white coloured waxy solids.

Here the monocarboxylic acid series discussed is equivalent to:

C_nH_2n+1COOH where n = 0, or a more structurally correct general formula for linear aliphatic monocarboxylic acids is HCOOH, then CH₃(CH₂)_nCOOH, where n = 0, 1, 2 etc. which I refer to as '1-ols' primary alkanols

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

Graph 1 purple line = RCOOH

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.

So, in this discussion we are comparing the red line (linear alkanes) with the  linear aliphatic monocarboxylic molecule AND comparing molecules with a similar 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 homologous series i.e. carbon chains with a different end functional group.

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, usually weak compared to covalent bonds, but not insignificant forces, known as instantaneous dipole - induced dipole forces.

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

As already mentioned in 6.2.1, hydrogen bonding between the carboxylic acid groups cause the formation of dimers.

However, unlike the lack of a clear melting point trend, the boiling points do follow the expected rising trend as the carbon chain gets longer.

The reason being, on melting, and with increasing rise in temperature, more and more of the hydrogen bonds of the 'dimer' molecules are broken.

So, by the time the liquid boils, most of the hydrogen bonding is randomised between the molecules, without dimer formation - a similar situation to the hydrogen bonding in liquid alcohols.

Don't forget, the hydrogen bonding is in addition to the intermolecular attractive force compared to non-polar molecules.

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

The hydrogen bond is directional i.e. the proton lines up with the lone pair on the oxygen which is effectively the delta minus and this should come out in a full diagram showing the hydrogen bonding between molecules.

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

The increase in intermolecular attractive forces, means the molecules need a higher kinetic energy to overcome the intermolecular forces and escape from the liquid surface, so they have a higher boiling point and increased enthalpy of vapourisation compared to alkanes.

For a broader discussion see on boiling points and intermolecular forces see:

Introduction to Intermolecular Forces

Detailed comparative discussion of boiling points of 8 organic molecules

Boiling point plots for six organic homologous series

and for wider reading on intermolecular bonding forces

Other case studies of boiling points related to intermolecular forces

 Evidence and theory for hydrogen bonding in simple covalent hydrides

 

Graph 2 purple line = RCOOH

A plot of the molecular mass of the  linear aliphatic monocarboxylic 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.

 

Graph 3 purple line = RCOOH

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

For the same carbon number, the primary alcohols have significantly higher boiling points than alkanes, mainly due to the hydrogen bonding.

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.

 

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6.2.3 The solubility of carboxylic acids

Reminder: Alcohols are permanently polarised molecule due to the highly polar bond ^δ–O–H^δ+ caused by the difference in electronegativities between oxygen and hydrogen i.e. O (3.5) > H (2.1). This causes the extra permanent dipole – permanent dipole interaction between neighbouring polar molecules via hydrogen bonding

The intermolecular hydrogen bonding in water

Before looking at the solubility of carboxylic acids, a reminder of the hydrogen bonding in water via the above diagram.

Important note, especially when drawing hydrogen bonding diagrams for any molecule!

You must clearly show the directional linearity of the O^δ--H^δ+ǁǁǁ:O^δ- arrangement of the hydrogen bond including the single O-H covalent bond and the lone pair on the other oxygen too!

You must do this accurately in exams when drawing intermolecular hydrogen bonding diagrams of water, alcohols and carboxylic acids, because it is the only specifically spatially directed intermolecular force, all the rest of the other types of intermolecular bonding forces are randomised.

The lower monocarboxylic acids exist as dimers in the vapour and pure liquid phases and when dissolved in non-polar solvents like benzene of hexane.

However, in aqueous solution, carboxylic acids can also hydrogen bond directly with water - which accounts for why the lower members are so soluble in water.

Although water - water hydrogen bonds are disrupted (O^δ--H^δ+ǁǁǁ:O^δ-), new carboxylic acid - water hydrogen bonds are formed e.g. (C-O^δ--H^δ+ǁǁǁ:O^δ--H^δ+) OR (H-O^δ--H^δ+ǁǁǁ:O^δ-=C^δ+) partly compensate for this.  (ǁǁǁ hydrogen bond)

BUT, there are limits to this effect, looking at the diagram below, only the first four linear carboxylic acids are completely soluble (miscible) in water.

The hydrogen bonding with water enables the first four carboxylic acids to be miscible with water (completely soluble in each other, irrespective of proportions), but after that, the solubility rapidly decreases.

BUT, 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 length of the hydrophobic 'hydrocarbon' chain makes the carboxylic acid 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 carboxylic acid, without compensating alcohol - water hydrogen bonds.

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

 

Benzoic acid, an aromatic carboxylic acid, is much more soluble in hot water than cold water and can be purified and recrystallised in this way using water as the solvent.

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