Part 4. The chemistry of ALCOHOLS

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 comparison of boiling points and solubility in water of alcohols and ethers examples theory explanation

Part 4.3 The physical properties of alcohols - boiling points and solubility - including intermolecular forces and a comparison with ethers

Sub-index for this alcohol chemistry page 4.3

4.3.1 The boiling points of alcohols and intermolecular forces

4.3.2 The solubility of alcohols and use as solvents

4.3.3 A comparison of the physical properties of alcohols and their isomeric ethers

4.3.1 The boiling points of alcohols and intermolecular forces

Lower alcohols in the homologous series C_nH_2n+1OH are colourless liquids.

Higher members of the series are white/cream coloured waxy solids.

Here the primary alcohol series discussed is equivalent to:

CH₃(CH₂)_nOH, where n = 0, 1, 2 etc. which I refer to as '1-ols'

For n = 0, methanol CH₃OH, for n = 1 ethanol CH₃CH₂OH, n = 5 hexan-1-ol CH₃(CH₂)₅OH etc.

The boiling point trend of linear primary alcohols (1-ols) are now discussed in detail and compared with other homologous series.

Graph 1 yellow line = alcohols

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 'yellow line' of linear primary alcohols, the graph goes from methanol CH₃OH (bpt 65^oC/338 K) to decan-1-ol CH₃(CH₂)₉OH (bpt 230^oC/503 K).

So, in this discussion we are comparing the red line (linear alkanes) with the line (linear primary alcohols, '1-ols') 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 (H^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 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.

R-O–H^δ+llll^δ–:O-R ... etc.

This is an added effect in attracting the alcohol molecules together much more strongly than just the instantaneous dipole - induced dipole forces - but it isn't necessarily the largest contributor to the total intermolecular force of attraction between molecules.

The total intermolecular force is hydrogen bonding (via the OH group) plus the instantaneous dipole - induced dipole attraction forces (from the whole molecule).

A minor contribution is from permanent dipole - induced dipole attraction.

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

Note that the lone pairs on the most electronegative atom are important to show on a fully detailed diagram (though I haven't always done so on this page).

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 effect of hydrogen bonding on the boiling point is very significant for the lower alcohols, but its effect decreases as the carbon chain length increases.

For methanol: Total intermolecular force = (61.3% instantaneous dipole – induced dipole) + (30.3% perm. dipole – permanent dipole including hydrogen bonding) + (8.4% permanent dipole – induced dipole)

For ethanol: (42.6%  instantaneous  dipole – induced dipole) + (47.6% permanent dipole – permanent dipole including H bonding) + (9.8% permanent dipole – induced dipole)

2–methylpropan–2–ol: Total intermolecular force was: (67.2%  instantaneous  dipole – induced dipole) + (23.1% permanent dipole – permanent dipole including H bonding) + (9.7% permanent dipole – induced dipole)

BUT, the % intermolecular attractive force from hydrogen bonding will tend to decrease, and the % contribution of the instantaneous dipole - induced dipole forces increases, as the alkyl chain gets longer and the boiling points of higher members of the linear alcohols converges towards the boiling point curve of alkanes!

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 yellow line = alcohols

A plot of the molecular mass of the linear primary alcohol 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 yellow line = alcohols

A plot of the carbon number of the linear primary alcohol 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.

 

Note on the boiling points of diols and triols

Ethane-1,2-diol (ethylene glycol, glycol), C₂H₆O₂, , bpt 198^oC

Propane-1,2-diol, C₃H₈O₂, ,  , bpt 188^oC

Propane-1,3-diol , ,   , bpt 213^oC

Propane-1,2,3-triol (glycerol), ,   bpt 290^oC and decomposes

All diols and triols will have relatively higher boiling points because effect of extra hydrogen bonding on the boiling point.

With each extra hydroxy group, there are more sites on the molecules to hydrogen bond with each other.

Comparison examples

(i) Ethane 1,2-diol has a molecular mass of 62 and propan-1-ol 60 (similar numbers of electrons).

Ethane-1,2-diol boils at 198^oC, 101^o higher than propan-1-ol, bpt 97^oC.

(ii) The propane diols have a molecular mass of 76 and butan-1-ol 74 (similar numbers of electrons).

The propane diols boil at 188/213^oC, 70/95^o higher than butan-1-ol, bpt 118^oC

(iii) Propane-1,2,3-triol has a molecular mass of 92 and pentan-1-ol 88 (similar numbers of electrons.

Propane-1,2,3-triol boils at 290^oC, 152^o higher than pentan-1-ol, bpt 138^oC.

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4.3.2 The solubility of alcohols and use as solvents

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 alcohols, 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 or alcohols (and carboxylic acids later) because it is the only specifically spatially directed intermolecular force, all the rest of the other types of intermolecular bonding forces are randomised.

 

Left diagram: The hydrogen bonding between alcohol molecules:

Right: The hydrogen bonding between water and alcohol molecules:

Although water - water hydrogen bonds are disrupted (O^δ--H^δ+ǁǁǁ:O^δ-), new alcohol - water bonds are formed (C-O^δ--H^δ+ǁǁǁ:O^δ--H^δ+) partly compensate for this.  (ǁǁǁ hydrogen bond)

BUT, there are limits to this effect, looking at the diagram below, only the first three alcohols are completely soluble (miscible) in water.

The hydrogen bonding with water enables the first three lower alcohols to be miscible with water (completely soluble in each other, irrespective of proportions), but after that, the solubility of linear primary alcohols 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 makes the alcohol 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 alcohol - water hydrogen bonds.

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

 

A comparison with diols and triols

Use of ethane-1,2-diol as an antifreeze in car engines. 'ethylene glycol', as it is known as, lowers the freezing point of water, but the boiling point (198^oC) is too high to vapourise out of the water, it is also very soluble in water - miscible.

 

The effect of an extra hydroxy group on larger molecules can be seen by comparing the solubilities of hexan-1-ol (very low) and hexane-1,2-diol which is fully miscible (see also a more extreme example with the polymer PVA).

Hexan-1-ol, with its one hydroxy group and a hydrocarbon chain has a very low solubility in water.

Hexane-1,2-diol, with an extra hydroxy group, but with the same length of carbon chain, is much more soluble and is miscible with water.  Hexan-1,2-diol has double the 'molecular' sites to hydrogen bond with water and which greatly enhances its ability to dissolve in water..

 

Even large molecules like poly(ethenol), also known as PVA, poly(vinyl alcohol), can dissolve in water because the regular occurrence of hydroxy groups along the polymer chain that can hydrogen bond with water.

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