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Advanced A/AS Level Organic Chemistry: Boiling points and solubility of alcohols

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

including boiling points, solubility & intermolecular forces and a comparison with ethers

Sub-index for this page

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

INDEX of notes on ALCOHOLS chemistry

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4.3.1 The boiling points of alcohols and intermolecular forces

Lower alcohols in the homologous series CnH2n+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 of boiling point of primary alcohols versus electrons in molecule advanced organic chemistry revision notes doc brown Graph 1 yellow line = alcohols

The red line graph shows the boiling point of alkanes from methane CH4 (boiling point -164oC/109 K)  to tetradecane C14H30 (boiling point 254oC/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 65oC/338 K) to decan-1-ol CH3(CH2)9OH (bpt 230oC/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 (Hδ+-Oδ-) 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.

diagram of intermolecular hydrogen bonding forces between liquid alcohol molecules doc brown A level organic chemistry revision notes R-O–Hδ+llllδ: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 δO–Hδ+ 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 of boiling point of primary alcohols versus molecular mass of molecule advanced organic chemistry revision notes doc brown 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 of boiling point of primary alcohols versus carbon atoms number in molecule advanced organic chemistry revision notes doc brown 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, diols triols and cyclo-alcohols structure and naming (c) doc b, diols triols and cyclo-alcohols structure and naming (c) doc b bpt 198oC

Propane-1,2-diol, C3H8O2, diols triols and cyclo-alcohols structure and naming (c) doc b, diols triols and cyclo-alcohols structure and naming (c) doc b , bpt 188oC

Propane-1,3-diol , diols triols and cyclo-alcohols structure and naming (c) doc b, diols triols and cyclo-alcohols structure and naming (c) doc b  , bpt 213oC

Propane-1,2,3-triol (glycerol), diols triols and cyclo-alcohols structure and naming (c) doc b, diols triols and cyclo-alcohols structure and naming (c) doc b   bpt 290oC 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 198oC, 101o higher than propan-1-ol, bpt 97oC.

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

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

(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 290oC, 152o higher than pentan-1-ol, bpt 138oC.

 


4.3.2 The solubility of alcohols and use as solvents

Reminder: Alcohols are permanently polarised molecule due to the highly polar bond δO–Hδ+ 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

diagram of intermolecular hydrogen bonding forces between liquid water molecules doc brown A level chemistry revision notes 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 Oδ--Hδ+ǁǁǁ:Oδ- 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.

 

diagram of intermolecular hydrogen bonding forces between liquid alcohol molecules doc brown A level organic chemistry revision notes   diagram of intermolecular hydrogen bonding forces between water and alcohol molecules in solution mixture why miscible doc brown A level chemistry revision notes

Left diagram: The hydrogen bonding between alcohol molecules:

Right: The hydrogen bonding between water and alcohol molecules:

Although water - water hydrogen bonds are disrupted (Oδ--Hδ+ǁǁǁ:Oδ-), new alcohol - water bonds are formed (C-Oδ--Hδ+ǁǁǁ:Oδ--Hδ+) 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.

skeletal formula diagram and solubility data for primary alcohols in water doc brown's advanced A level organic chemistry notes

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. 'Glycerol', as it is known as, lowers the freezing point of water, but the boiling point (198oC) is too high to vapourise out of the water, its also very soluble in water.

 

comparing the solubility of hexan-1-ol & hexane-1,2-diol in water hydrogen bonding doc brown's advanced A level organic chemistry revision notes

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..

 

hydrogen bonding between poly(ethenol) PVA polyvinyl alcohol molecules and water advanced A level organic chemistry doc brown's revision notes

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.

 


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

Abbreviations: In both text and diagrams, R-OH refers to alcohols and R-O-R refers to ethers where R = alkyl.

In the diagrams note the significance of the δ+ δ- partial charges and the O:δ+llllδ- H- hydrogen bonding.

(a) Introduction - intermolecular forces between water, alcohol and ether molecules

diagram of stronger intermolecular forces between alcohol molecules is due to instantaneous dipole - induced dipole plus permanent dipole - permanent dipole attractive forces plus the biggest contribution from hydrogen bonding (which cannot happen between ether molecules) diagram of intermolecular hydrogen bonding forces between water and alcohol molecules in solution mixture why miscible doc brown A level chemistry revision notes

Left (i) The  intermolecular forces between alcohol molecules is due to instantaneous dipole - induced dipole plus permanent dipole - permanent dipole attractive forces plus a bigger contribution from hydrogen bonding (δ+llllδ- ,which cannot happen between ether molecules).

Right (ii) The hydrogen bonding (δ+llllδ-) between water molecules and alcohol molecules enabling alcohols to be much more soluble in water than slightly polar ether molecules (and even more so than non-polar alkanes).

 

diagram of the weak intermolecular forces between ether molecules due to instantaneous dipole - induced dipole plus permanent dipole - permanent dipole attractive forces  diagram of strong hydrogen bonding between water molecules and the weaker intermolecular forces between water and ether molecules - hence the much lower solubility of ethers in water compared to alcohols

Left (iii) The weaker intermolecular forces between ether molecules due to instantaneous dipole - induced dipole plus permanent dipole - permanent dipole attractive forces, making them slightly polar molecules, but no hydrogen bonding between ether/ether or ether/water interactions.

Right (iv) The strong hydrogen bonding (δ+llllδ-) between water molecules and the weaker intermolecular forces between water and ether molecules - hence the much lower solubility of ethers in water compared to alcohols.

We can now apply these ideas to boiling points and solubility in sections (b) to (d)

 

(b) Comparison of the solubility of alcohols, ethers and alkanes in water

1. Non-polar alkanes are more or less insoluble in water, no permanent dipole - permanent dipole intermolecular forces between solute and solvent.

2. Lower ethers are slightly soluble in water, there are weak permanent dipole - permanent dipole intermolecular forces between ether and water molecules, but hydrogen bonding between water and alcohol molecules aid solvation to dissolve.

Solubilities in water e.g.

ethoxyethane CH3CH2OCH2CH3, 6.0 g/100 cm3 water

butan-1-ol CH3CH2CH2CH2OH, 7.3 g/100 cm3 water

I thought there might be a greater difference, but butan-1-ol has a longer hydrophobic tail.

3. Lower alcohols are very soluble in water and e.g. methanol and ethanol are completely miscible, but no ether is miscible with water.

 

(c) Comparison of solvent uses of alcohols, ethers and alkanes

Non-polar liquid alkanes will readily dissolve many non-polar organic molecules, but lower solubilities for highly polar molecules.

Slightly polar liquid (lower) ethers will dissolve a variety of polar and non-polar organic molecules e.g. solvents for fats, oils, waxes, perfumes, resins, dyes, gums, and hydrocarbons.

Lower alcohols will readily dissolve polar organic molecules e.g. solvents for marker pen inks, medicines, and cosmetics (such as deodorants and perfumes)..

 

(d) Comparison of boiling points of alcohols, ethers and alkanes

The data table below compares the molecular formula, molecular structure, relative molecular mass (Mr), number of electrons in the molecule and boiling point of selected alcohols, ethers and alkanes.

There are three groups of lower members in their respective homologous series and selected for actual/similar molecular mass, actual/similar molecular formula and the same number of electrons in the molecule.

Within each group of molecules (coloured banded green or cyan) the only difference in formula is that alkanes have an extra CH2 group instead of the oxygen atom present in alcohols and ethers.

The lower ether molecules are gases or volatile liquids at room temperature and dangerously flammable.

Molecule Molecular formula Molecular structure Mr Electrons Bpt/oC
Ethanol C2H6O CH3CH2OH 46 26 78oC
Methoxymethane C2H6O CH3OCH3 46 26 -25
Propane C3H8 CH3CH2CH3 44 26 -42
Propan-1-ol C3H8O CH3CH2CH2OH 60 34 97
Propan-2-ol C3H8O CH3CH(OH)CH3 60 34 82
Methoxyethane C3H8O CH3CH2OCH3 60 34 7oC
Butane C4H10 CH3CH2CH2CH3 58 34 -0.5
Butan-1-ol C4H10O CH3CH2CH2CH2OH 74 42 117
Butan-2-ol C4H10O CH3CH2CH(OH)CH3 74 42 99
2-methylpropan-1-ol C4H10O  (CH3)2CHCH2OH) 74 42 108
2-methylpropane-2-ol C4H10O (CH3)3COH) 74 42 82
Ethoxyethane C4H10O CH3CH2OCH2CH3 74 42 34
1-methoxypropane C4H10O CH3CH2CH2OCH3 74 42 38oC
2-methoxypropane C4H10O (CH3)2CHOCH3 74 42 33
Pentane C5H12 CH3CH2CH2CH2CH3 72 42 36

Comments on the data table.

1. For the same molecular formula, the alcohols have significantly higher boiling points than ethers and very much higher for the alkane of similar molecular mass and number of electrons in the molecule.

The extra contribution of hydrogen bonding between alcohol molecules makes all the difference, raising the enthalpy of vaporisation and boiling point.

2. Neither ethers or alkanes can exhibit hydrogen bonding, hence the relatively much weaker intermolecular forces and lower boiling points than alcohols.

3. Ether molecules are slightly polar, but not as much as alcohols, but their boiling points are still higher than the non-polar alkanes - so there is some permanent dipole - permanent dipole interaction between ether molecules, which is absent in alkanes.

 


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