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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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Chemistry Notes
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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 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.
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
δ–O–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 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 198oC
Propane-1,2-diol,
C3H8O2,
 ,
 ,
bpt 188oC
Propane-1,3-diol ,
 ,
, bpt 213oC
Propane-1,2,3-triol
(glycerol),
 ,
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
δ–O–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
 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.
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.
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.
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.
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
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).
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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INDEX of notes on ALCOHOLS
chemistry
All Advanced Organic
Chemistry Notes
Index of GCSE/IGCSE Oil - Useful Products
Chemistry Revision Notes
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