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Brown's Chemistry Advanced A Level Notes - Theoretical–Physical
Advanced Level
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 8
Part
8.7 (i) The Solubility of covalent compounds
and (ii) Miscible and
immiscible liquids
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INDEX for Part 8.
Phase equilibria–vapour
pressure, boiling point and intermolecular forces
Index of ALL my chemical equilibrium
context revision notes Index
ALL my advanced A
level theoretical
chemistry revision study notes
What factors control how soluble a
covalent compound is in water or organic solvents.
Why do some liquids
mix (miscible) and others not (immiscible)?
In both situations the role
of intermolecular forces and intermolecular bonding is discussed.
8.7 Solubility
of solid covalent compounds and miscible/immiscible of liquids
An approach to the covalent case studies described
below
(Solubility of
inorganic salt like compounds are considered in
Thermodynamics Part 2)
(solubility product is
dealt with in Equilibria section 4.2 and partition calculations in
section 4.1)
-
The solubility of one
material in another is a consequence of the physical and chemical
interactions involved.
-
A solvation process
may involve intermolecular force interactions or a chemical change
or both.
-
When solids dissolve in
liquids, a resulting saturated solution (maximum solubility) is another equilibrium situation.
-
Similarly mixing two
immiscible liquids is also an equilibrium situation even if only a
trace of one dissolves in the other.
-
It is important that you
have read the
section 8.2.1 on the nature of intermolecular forces because they all
get mentioned again here, but not in as much detail, I will assume a
working knowledge of them in
section 8.7
on solubility and miscibility.
-
In considering whether A
(liquid or solid solute) will dissolve in a liquid (e.g. the
solvent) three possible interactions should be considered i.e. the
intermolecular forces involved and solvation processes.
-
The three 'interactions'
will simply be A...A,
A...B and
B...B
associations via intermolecular forces.
-
If either
A...A
or B...B associations are strong then one of them will
dominate and A and B will be less likely to mix i.e. if two liquids, they are likely to be
immiscible OR
solid A won't dissolve in liquid B.
-
However if an A...B
interaction is strong or comparable to A...A or B...B associations
then two liquids are likely to be miscible and mix with each other
OR solid A dissolves in liquid B.
-
This conceptual way of
thinking is a bit more sophisticated than 'like dissolves like', but
the latter has much to commend it as a rough rule of thumb.
-
If two liquids dissolve
appreciably in each other but still two layers are formed above and
below a meniscus, the two liquids may be referred to as partially
miscible.
-
If you can think like
this, then most cases of what dissolves/mixes with what are usually
quite understandable!
The cases discussed below
are: Case study 8.7.1 Mixing hydrocarbons and water
* Case study 8.7.2 The solubility of alcohols in
water * Case study 8.7.3
The solubility of carboxylic acids in water * Case
Study 8.7.4 Why are pairs of some organic solvents
miscible and others immiscible? * Case Study 8.7.5
Are dissolved molecules always in the same molecular state whatever the
solvent? * 8.7.6
The solubility of some gases in
water
Case study 8.7.1
Mixing hydrocarbons
and water
-
The solubility of
hydrocarbons in water is extremely low and for practical purposes it
is virtually zero.
-
It doesn't matter
whether it is alkanes (lower–small or higher–large), alkenes, arenes
(aromatic hydrocarbons like benzene, methyl benzene etc.) or even
crude oil itself, non of them can associate with water.
-
Hydrocarbon –
hydrocarbon intermolecular forces are only the weak instantaneous
dipole – induced dipole interactions.
-
The polar water – oil
interaction will be of a permanent dipole – induced dipole nature,
but even if this is greater than the hydrocarbon – hydrocarbon
attraction, what will dominate this situation is the strongest of
the intermolecular forces, namely the hydrogen bonds between water
molecules.
-
....δ–:O–Hδ+
llll δ–:O–Hδ+.... rules ok!
-
If a hydrocarbon
molecule is 'shaken' into water, the hydrogen bonds will be
disrupted and will tend to reform and quite literally squeeze out
the hydrocarbon molecules.
-
I would expect methane
will be the most soluble hydrocarbon in water because it is the
smallest possible hydrocarbon and its 'insertion' will cause the
minimum disruption to water's hydrogen bonds.
-
solubility of methane, ethene, ethane etc.?
Case study 8.7.2
The solubility of alcohols
in water
linear aliphatic alcohol |
structural formula |
Mr |
solubility g/100g water |
solubility mol/dm3 |
methanol |
CH3OH |
32 |
miscible |
miscible |
ethanol |
CH3CH2OH |
46 |
miscible |
miscible |
propan–1–ol (1–propanol) |
CH3(CH2)2OH |
60 |
miscible |
miscible |
butan–1–ol (1–butanol) |
CH3(CH2)3OH |
74 |
8.0 |
1.08 |
pentan–1–ol (1–pentanol) |
CH3(CH2)4OH |
88 |
2.8 |
0.32 |
hexan–1–ol (1–hexanol) |
CH3(CH2)5OH |
102 |
0.6 |
0.059 |
heptan–1–ol (1–heptanol) |
CH3(CH2)6OH |
116 |
0.093 |
0.008 |
octan–1–ol (1–octanol) |
CH3(CH2)7OH |
130 |
0.054? |
0.004? |
nonan–1–ol (1–nonanol) |
CH3(CH2)8OH |
144 |
~insoluble |
~insoluble |
decan–1–ol (1–decanol) |
CH3(CH2)9OH |
158 |
~insoluble |
~insoluble |
-
The first three lower
members of the alcohol homologous series are completely miscible
with water.
-
This is because all the
A...A, B...B and B...B interactions are all similar, i.e. they can
all hydrogen bond with each other.
-
However, although the
hydroxy O–H group is hydrophilic ('solvent liking') and readily
hydrogen bonds with water, as well as itself, there is a counter
trend of the hydrophobic ('solvent hating') alkyl group which
disrupts the hydrogen bonds in water.
-
Therefore, as the alkyl
hydrocarbon chain increases in length, its greater hydrogen bond
disruption results in a rapidly decreasing solubility so that by the
time you reach octanol all successive alcohols are effectively
insoluble.
-
Note that if more than
one hydroxy group is present then the solubility will be increased
e.g.
-
Liquid
propane–1,2,3–triol (glycerol) has the
structure HOCH2CH(OH)CH2OH (Mr
= 92), so with its triple hydrogen bonding sites the glycerol –
water interactions are strong and both liquids are completely
soluble in each other i.e. fully miscible, despite being as bulky as
pentan–1–ol on a molecular mass basis.
-
Sugars such as glucose
(Mr = 180), despite the higher molecular mass, are highly
soluble in water due to the multiple hydrogen bonding sites on the
molecule (5 O–H groups).
-
Also note the
increase in viscosity ('stickiness') of glucose (solid) > liquid
propane–1,2,3–triol > liquid propane–1,2–diol > propan–1–ol,
-
which is accounted for
in terms of increased hydrogen bonding as well as the increased
instantaneous dipole – induced dipole forces from the increasing
number of electrons.
- Although for linear ..1-ols, from
nonan-1-ol onwards are ~insoluble, they will act as surfactants and
form a monomolecular layer on the surface of water.
Case study 8.7.3
The solubility of carboxylic acids
in water
linear carboxylic acid |
structural formula |
Mr |
solubility g/100g water |
solubility mol/dm3 |
methanoic acid |
HCOOH |
46 |
miscible |
miscible |
ethanoic acid |
CH3COOH |
60 |
miscible |
miscible |
propanoic acid |
CH3CH2COOH |
74 |
miscible |
miscible |
butanoic acid |
CH3(CH2)2COOH |
88 |
miscible |
miscible |
pentanoic acid |
CH3(CH2)3COOH |
102 |
5.0 |
0.49 |
hexanoic acid |
CH3(CH2)4COOH |
116 |
1.1 |
0.095 |
heptanoic acid |
CH3(CH2)5COOH |
130 |
0.24 |
0.017 |
octanoic acid |
CH3(CH2)6COOH |
144 |
0.068 |
0.0047 |
nonanoic acid |
CH3(CH2)7COOH |
158 |
0.03 |
0.0018 |
decanoic acid |
CH3(CH2)8COOH |
172 |
~insoluble |
~insoluble |
-
The solubility trend is
similar to that of the linear alcohols in case study 8.7.2 as are
the reasons for the trend.
-
The first four lower
members of this carboxylic acid homologous series are completely
miscible with water and the solubilities, on a molar basis, tend to
be a bit higher than for alcohols, presumably because they are more
polar and have two sites for H bonding (C=O and O–H).
Case 8.7.4
Why are some pairs of
organic solvents miscible and others immiscible?
(in these examples the
cases are based on mixing equal volumes of each liquid)
-
Propanone and benzene
are miscible
-
(polar) propanone –
propanone intermolecular forces = (instantaneous dipole –
induced dipole) + (permanent dipole – permanent dipole)
-
(non–polar) benzene
– benzene intermolecular forces = (instantaneous dipole –
induced dipole)
-
Although propanone
is polar, most of its intermolecular force is still derived from
instantaneous dipole – induced dipole attractive forces and the
extra permanent dipole – permanent dipole force does not
prevent miscibility.
-
Propanone and
butan–1–ol are miscible
-
(polar) propanone –
propanone intermolecular forces = (instantaneous dipole –
induced dipole) + (permanent dipole – permanent dipole)
-
(more polar)
butan–1–ol – butan–1–ol intermolecular forces = (instantaneous
dipole – induced dipole) + (hydrogen bonding)
-
Both molecules are
polar and butan–1–ol can hydrogen bond with propanone as well as
with itself, so miscibility is no problem.
-
Glycerol
(propan–1,2,3–triol) and propanone are immiscible
-
(polar) propanone –
propanone intermolecular forces = (instantaneous dipole –
induced dipole) + (permanent dipole – permanent dipole)
-
(much more polar)
glycerol – glycerol intermolecular forces = (instantaneous
dipole – induced dipole) + (multiple hydrogen bonding)
-
Glycerol has the
structure HOCH2CH(OH)CH2OH, so with its
triple hydrogen bonding sites the glycerol – glycerol attraction
will outweigh any glycerol – propanone interactions.
-
Most organic liquids
readily dissolve in butan–1–ol (1–butanol)
-
Being highly polar
with its hydroxy group, butan–1–ol will readily dissolve many
other polar organic compounds.
-
However, the alkyl
hydrocarbon tail means the hydrogen bonding between the alcohol
molecules doesn't totally dominate the situation and many
relatively non–polar organic liquids will also dissolve in
butan–1–ol.
Case 8.7.5
Are dissolved molecules always
in the same molecular state whatever the solvent?
8.7.6 The solubility of some gases in
water
-
The maximum solubility is quoted
as the volume of gas dissolved in 1 cm3 of water at 0oC
at 1 atm pressure of the gas (101325 Pa).
-
This corresponds to the
equilibrium
X(g)
X(aq)
where X is the gas dissolving in water.
-
This is only quoted specifically where
the gas reacts with water moving this equilibrium further to the
right i.e. increasing solubility.
-
From Avogadro's Law,
equal volumes of gases contain the same number, the data represents
solubility on a molar basis.
-
Solubility of gases
decreases on increase in temperature. and increases with increase in
pressure of the gas.
-
In most cases the
solubility is low because the dissolving molecule is disrupting the
hydrogen bonds between water molecules.
-
*
However, in a few cases,
there is a chemical reaction with water OR strong hydrogen
bonding occurs between solute and the solvent water, and in such
cases the solubility is considerably increased.
-
When comparing any two
of the molecules in the list, you can produce reasonable arguments
to explain either the difference or similarity in solubility and
I've quoted a few such arguments.
Gases in alphabetical order |
Solubility at 0oC |
Structure |
Mr |
electrons |
Comments |
*
Ammonia |
1300 |
NH3 polar molecules |
17 |
10 |
Extremely soluble, can strongly hydrogen bond with water
and a few % of the dissolved ammonia reacts with water –
ionisation.
llll
δ–N–Hδ+
llll δ–O–Hδ+ llll δ–N–Hδ+
llll δ–O–Hδ+ llll
etc.
NH3(g)
NH3(aq)
+ H2O(l)
NH4+(aq) +
OH–(aq) |
Argon |
0.056 |
Ar atoms |
40 |
18 |
Weak
instantaneous dipole – induced dipole interactions, very low
solubility |
Carbon dioxide |
1.71 |
CO2 linear molecules |
44 |
22 |
The effect
of the permanent dipoles (δ+C=Oδ–)
is cancelled out because of the linearity of the molecule,
but its greater solubility compared to molecules like
oxygen, nitrogen etc. may be due to hydrogen bonding? and
the tiny % of the ionisation reaction:
CO2(g)
CO2(aq) +
H2O(l)
H3O+(aq) +
HCO3–(aq) |
Carbon monoxide |
0.035 |
CO diatomic molecules |
28 |
14 |
Not a very
polar molecule, so low solubility via instantaneous
dipole – induced dipole forces |
Chlorine |
4.61 |
Cl2 non–polar
diatomic molecules |
71 |
34 |
A larger
molecule of 34 electrons increasing the instantaneous dipole
– induced dipole forces so increasing the solubility
compared to other molecules of less electrons in the table
and also a tiny % of the following ionisation reaction:
Cl2(g)
Cl2(aq) +
H2O(l)
H3O+(aq) +
Cl–(aq)
+ HClO(aq) |
Ethene |
0.25 |
C2H4
non–polar planar molecules |
28 |
16 |
Weak
interaction but 5 times more soluble than methane, perhaps
because the pi electron clouds of the C=C double bond are
more polarizable increasing the instantaneous dipole – induced
dipole interaction? |
Helium |
0.0094 |
He atoms of only 2 electrons |
4 |
2 |
Only the
very weakest instantaneous dipole – induced dipole forces
despite the minimum disruption of water's hydrogen bonds,
hence extremely low solubility. |
Hydrogen |
0.021 |
H2 non–polar diatomic
molecules of only 2 electrons |
2 |
2 |
Only very
weak instantaneous dipole – induced dipole forces despite
the minimum disruption of water's hydrogen bonds, hence very
low solubility. |
*
Hydrogen chloride |
506 |
HCl polar diatomic molecules |
36.5 |
18 |
Extremely soluble because of a ~100% reaction with water
to give ~10 molar hydrochloric acid.
HCl(g) +
H2O(l) ==> H3O+(aq)
+ Cl–(aq)
Similarly hydrogen bromide
(HBr) and hydrogen iodide (HI) are also very soluble for the
same reason. |
Hydrogen sulfide |
4.68 |
H2S slightly polar
bent molecule |
34 |
18 |
Hydrogen
sulphide is slightly polar but the moderate solubility is
due mainly to instantaneous dipole – induced dipole forces
and a tiny % chemical reaction via
H2S(g)
H2S(aq)
+ H2O(l)
H3O+(aq) +
HS–(aq) |
Nitrogen |
0.024 |
N2 non–polar diatomic
molecules |
28 |
14 |
Low
solubility due to being a more bulky molecule than say H2
or He but increased electrons increases instantaneous dipole
– induced dipole forces gives nitrogen a slightly greater
solubility. |
Methane |
0.054 |
CH4 simplest
hydrocarbon molecule |
16 |
10 |
Only very
weak instantaneous dipole – induced dipole forces despite
the minimum disruption of water's hydrogen bonds for a small
molecule, hence very low solubility. |
Oxygen |
0.049 |
O2 diatomic molecules |
32 |
16 |
Low
solubility due to being a more bulky molecule than say H2
or He but increased electrons increases instantaneous dipole
– induced dipole forces gives oxygen a slightly greater
solubility. |
Sulfur dioxide |
79.9 |
SO2 bent polar
molecules |
64 |
32 |
Sulphur
dioxide |
Henry's Law and Gas Solubility
-
Henry's Law states
that the mass of slightly soluble gas that dissolves in a definite
mass of liquid at a given temperature is very nearly proportional to
the partial pressure of that gas (in contact with the liquid).
-
This law holds for gases
that do NOT react with the liquid solvent (e.g. hydrogen
chloride in water) OR have a very strong intermolecular
attraction with the liquid (eg. ammonia in water).
-
One mathematical
expression of Henry's Law is:
px
= KH c
-
where
px
= the partial pressure of the dissolved gas X (i.e. px
above the liquid), KH
is Henry's Constant (which depends on the gas solute, the liquid
solvent and the ambient temperature, and
c is the
concentration of the dissolved gas solute.
-
For relatively non–polar
molecules of low solubility you do get a 'reasonably' linear graph
of concentration (x) versus partial pressure (y)
-
with the axis origin 0,0
and this corresponds to the equilibrium
WHAT NEXT?
INDEX for Part 8.
Phase equilibria–vapour
pressure, boiling point and intermolecular forces
Index of ALL my chemical equilibrium
context revision notes Index
Part 8 sub–index:
8.1 Vapour pressure origin and examples * 8.2.1
Introduction to Intermolecular Forces * 8.2.2
Detailed comparative discussion of boiling points of 8 organic molecules * 8.3
Boiling point plots for six
organic
homologous series * 8.4 Other case studies of
boiling points related to intermolecular forces * 8.5
Steam
distillation – theory and practice * 8.6 Evidence and theory
for hydrogen bonding in simple covalent hydrides *
8.7 Solubility of covalent compounds, miscible and
immiscible liquids
Advanced Equilibrium Chemistry Notes Part 1. Equilibrium,
Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and
calculations * Part 3.
Equilibrium and industrial processes * Part 4.
Partition,
solubility product and ion–exchange * Part 5.
pH, weak–strong acid–base theory and
calculations * Part 6. Salt hydrolysis,
Acid–base titrations–indicators, pH curves and buffers *
Part 7.
Redox equilibria, half–cell electrode potentials,
electrolysis and electrochemical series
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