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Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study
Notes for UK IB KS5 A/AS GCE advanced level physical theoretical chemistry
students US K12 grade 11 grade 12 theoretical chemistry
Chemical Equilibrium Revision Notes PART 8.4
8.4 Further case studies comparing boiling point,
intermolecular forces & electrons in a molecule
Further comparisons of the boiling points of
some series of inorganic and organic compounds and relating their
boiling points to the intermolecular forces (intermolecular bonding)
involved.
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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
8.4 Some extra
case studies on boiling point, intermolecular forces and electrons in a
molecule
Note 8 organic molecules are discussed
in detail in section 8.2.1 and continued
in section 8.2.2
Table 8.4.1 Comparing 10
electron species – noble gas, some very small organic/inorganic
molecules |
MOLECULE |
formula |
Mr |
electrons |
Bpt |
ΔHvap/kJmol–1 |
Dipole
moment/D |
neon |
Ne |
20 |
10 |
27K/–246oC |
1.8 |
0.00 |
methane |
CH4 |
16 |
10 |
109K/–164oC |
8.2 |
0.00 |
ammonia |
NH3 |
17 |
10 |
240K/–33oC |
23.4 |
1.48 |
hydrogen fluoride |
HF |
20 |
10 |
293K/20oC |
7.5 |
1.91 |
water |
H2O |
18 |
10 |
373K/100oC |
41.1 |
1.84 |
-
Series 8.4.1
Comparing 10 electron species
- This series offers quite a range
of boiling points!
- Neon, Ne (can consider as
a monatomic molecule')
- The most compact 10 electron
system you can have, but not a good baseline for further
discussion because it is evident from the following four
molecules, as soon as
- Even in the case of
helium,
lowest boiling point of any substance, you still can get
transient dipoles because of the random behaviour of the
electrons ...
-
attractions
- Methane, CH4
- Methane is a highly symmetrical
molecule with no significant bond polarity.
- Clearly, in comparing
methane with water, ammonia or hydrogen fluoride, which all
exhibit hydrogen bonding, the non–polar nature of methane
gives rise to considerably weaker intermolecular forces.
- Ammonia, NH3
- Exhibits hydrogen bonding
increasing the boiling point considerably.
- ....δ–:N–Hδ+llllδ–:N–Hδ+.... etc.
- Hydrogen fluoride, HF
- This is known to display
hydrogen bonding and a greatly increased boiling point would be
expected compared to methane.
- It can only form one hydrogen
bond per molecule (*) and so the boiling point is considerably
raised compared to methane BUT not as much as for water.
- (*) It does form moderately
stable zig–zag
'polymer' molecules in the gaseous phase just above its boiling
point, which is probably why it has a higher boiling point than
hydrogen bonded ammonia.
- ....δ–:F–Hδ+
llll
δ–:F–Hδ+.... etc.
-
More on hydrogen bonding
- Water, H2O
Comparing
some molecules with 18 electrons
Table 8.4.2 Comparing
18
electron species – noble gas, small organic/inorganic
molecules |
MOLECULE |
formula |
Mr |
electrons |
Bpt |
ΔHvap/kJmol–1 |
Dipole
moment/D |
argon |
Ar |
40 |
18 |
87K/–186oC |
6.53 |
0.00 |
fluorine |
F2 |
38 |
18 |
85K/–188oC |
3.16 |
0.00 |
ethane |
CH3CH3 |
30 |
18 |
184K/–89oC |
? |
0.00 |
fluoromethane |
CH3F |
34 |
18 |
185K/–88oC |
? |
0.00 |
phosphine |
PH3 |
34 |
18 |
185K/–88oC |
? |
0.55 |
hydrogen sulphide |
H2S |
34 |
18 |
212K/–61oC |
18.7 |
0.92 |
hydrogen chloride |
HCl |
36.5 |
18 |
188K/–85oC |
16.2 |
1.05 |
methanol |
CH3OH |
32 |
18 |
338K/65oC |
43.5 |
1.71 |
hydrazine |
H2NNH2 |
32 |
18 |
386K/113oC |
? |
? |
-
Series 8.4.2 Comparing 18
electron species
- argon, Ar
(can consider as a monatomic molecule')
- With 18 e's the boiling point is
raised by 60o compared to the 10 electron neon.
- Argon can be liquefied, like all
the noble gases because the random behaviour of electrons in the
orbitals of the electron clouds there is never perfectly
symmetrical distribution of electrical charge so transient
dipole – induced dipole forces only.
- In fact down
Group 0/18 the boiling points steadily increase down the group with increasing
numbers of
electrons and the greater polarizability of the
atom increasing the instantaneous dipole – induced dipole forces.
- i.e. 86Rn >
54Xe > 36Kr > 18Ar > 10Ne
(pre–subscript = proton/atomic number = electrons)
- fluorine, F2
- Its boiling point is very similar
argon despite the greater spread of electron charge over both atoms
of the diatomic molecule.
- With the highest electronegativity,
perhaps the electrons are as compact as in argon?
- ethane, CH3CH3
- Despite having the same number
of electrons as argon, the boiling point is 97o
higher.
- Ethane can act as a baseline for
the rest of this series since the electrons are more spread out
and it is a totally non–polar molecule with only
instantaneous dipole – induced dipole forces operating.
- fluoromethane, CH3F
- Despite the fact that the C–F bond
is quite polar, this permanent dipole seems to contribute virtually
nothing to increasing the intermolecular forces and fluoromethane
has similar boiling point to ethane (in fact it is actually less!).
-
phosphine, PH3
- The electronegativity of P is
2.1, the same as hydrogen, therefore PH3 is a
non–polar molecule with a boiling point similar to the equally
non–polar ethane.
- hydrogen sulphide, H2S
- The electronegativity of S is
2.5 so the S–H bond is slightly polar and there is a small
increase in boiling point compared to the non–polar ethane and
phosphine and has a higher boiling point than hydrogen chloride
with its more polar bond. This may be due to the fact that there
are two S–H bonds and the electrons are more spread out and more
polarisable?
- Total intermolecular force =
(instantaneous dipole – induced dipole) + (permanent dipole
–
permanent dipole including hydrogen bonding) + (permanent dipole
–
induced dipole)
- Hydrogen chloride, HCl
- Similarly, in the case of the
obviously polar hydrogen chloride molecule, the presence of the
permanent dipole has virtually no effect on the boiling point
compared to what you might expect for an 18 electron non–polar
molecule. This suggests that the intermolecular forces operating in
the first three molecules all have their origin in the instantaneous
dipole – induced dipole forces, despite the picture below!
-
attractions
seem to make little difference to the bpt!
- Total intermolecular force =
(88% instantaneous dipole – induced dipole) + (7.5% permanent dipole
–
permanent dipole including hydrogen bonding) + (4.5% permanent dipole
–
induced dipole)
- methanol, CH3OH
- Exhibits hydrogen bonding, so the
extra intermolecular force increases the boiling point considerably, perhaps by nearly
doubling it ...
- so for CH3–δ–O–Hδ+
we have the
δ–O–Hδ+llllδ–:O–Hδ+
situation ...
- ... and the permanent dipole –
permanent dipole attractions of the hydrogen bonding
play a very significant % of the intermolecular forces.
- Total intermolecular force =
(61.3% instantaneous dipole – induced dipole) + (30.3% perm. dipole
–
permanent dipole including hydrogen bonding) + (8.4% permanent dipole
–
induced dipole)
-
More on hydrogen bonding
- hydrazine, N2H4
or H2N–NH2
- Hydrazine's boiling point is another
5oo higher than the hydrogen bonded methanol and with good reason,
there are two sites for hydrogen bonding on the same small molecule
(x–ref water). Both –NH2 groups have the highly N–H
bond giving rise to the permanent dipole – permanent dipole
attractions of hydrogen bonding.
- ...H–N:δ–llllδ+H2NH–NH2δ+llll
δ–:N–H...
Comparing
some organic molecules with totals of 24-26 electrons
Table 8.4.3 Comparing
24–26
electron species – smaller linear organic
molecules (3 C/O/N atoms) |
MOLECULE |
formula |
Mr |
electrons |
Bpt |
ΔHvap/kJmol–1 |
Dipole moment/D |
propane |
CH3CH2CH3 |
46 |
26 |
K/–42oC |
? |
0.00 |
methoxymethane |
CH3OCH3 |
46 |
26 |
K/–25oC |
? |
1.30 |
chloromethane |
CH3Cl |
50.5 |
26 |
K/–24oC |
? |
1.87 |
ethanal |
CH3CHO |
44 |
24 |
K/21oC |
? |
? |
ethanol |
CH3CH2OH |
46 |
26 |
K/78.5oC |
38.6 |
1.70 |
methanoic acid |
HCOOH |
46 |
24 |
K/101oC |
? |
1.52 |
methanamide |
HCONH2 |
45 |
24 |
K/193oC
(decomposes) |
? |
? |
Comparing
some organic molecules with totals of 56-58 electrons
Table 8.4.4 Comparing
56–58
electron species – larger linear organic
molecules (7 C/O/N atoms) |
MOLECULE |
formula |
Mr |
electrons |
Bpt |
ΔHvap/kJmol–1 |
Dipole moment/D |
heptane |
CH3CH2CH2CH2CH2CH2CH3 |
100 |
58 |
371K/98oC |
36.6 |
0.00 |
methoxypentane |
CH3CH2CH2CH2CH2OCH3 |
102 |
58 |
– |
? |
? |
1–fluorohexane |
CH3CH2CH2CH2CH2CH2F |
104 |
58 |
365K/92oC |
? |
? |
1–chloropentane |
CH3CH2CH2CH2CH2Cl |
106.5 |
58 |
381K/108oC |
? |
? |
1–methylethyl ethanoate |
CH3COOCH(CH3)2 |
102 |
56 |
K/93oC |
? |
? |
methyl butanoate |
CH3CH2CH2COOCH3 |
102 |
56 |
– |
? |
? |
ethyl propanoate |
CH3CH2COOCH2CH3 |
102 |
56 |
K/99oC |
? |
? |
propyl ethanoate |
CH3COOCH2CH2CH3 |
102 |
56 |
K/102oC |
? |
? |
butyl methanoate |
HCOOCH2CH2CH2CH3 |
102 |
56 |
K/107oC |
? |
? |
hexanal |
CH3CH2CH2CH2CH2CHO |
100 |
56 |
402K/129oC |
? |
? |
hexan–2–one |
CH3COCH2CH2CH2CH3 |
100 |
56 |
400K/127oC |
? |
? |
hexan–1–ol |
CH3CH2CH2CH2CH2CH2OH |
102 |
58 |
429K/156oC |
? |
1.60 |
pentanoic acid |
CH3CH2CH2CH2COOH |
102 |
56 |
459K/186oC |
? |
? |
pentanamide |
CH3CH2CH2CH2CONH2 |
– |
– |
– |
? |
? |
-
Series Table 8.4.4 Comparing 56–58
electron species – larger linear organic
molecules
- heptane,
CH3CH2CH2CH2CH2CH2CH3
- 1–fluorohexane, CH3CH2CH2CH2CH2CH2F
- 1–chloropentane, CH3CH2CH2CH2CH2Cl
- hexanal,
CH3CH2CH2CH2CH2CHO
- hexan–2–one, CH3COCH2CH2CH2CH3
- hexan–1–ol,
CH3CH2CH2CH2CH2CH2OH
- pentanoic acid,
CH3CH2CH2CH2COOH
- Note that, even in the last two
molecules, most of the intermolecular attractive force will still
originate from the temporary dipole – induced dipole interactions.
The base–line' is about 370K for heptane, so the permanent dipoles/H
bonding have increased the boiling point of pentanoic acid by about
23%.
8.4.5 Some after thoughts!
- What we can explain by using
permanent dipole arguments, is why a boiling point of a polar
molecule is raised above
an expected value for a non–polar molecule with no permanent dipole,
but remember that for most molecules the majority of the
intermolecular force is due to transient dipole – induced dipole
attraction.
- What is not easy, is to explain is why
the C–F, C–Cl and H–Cl polar bonds seem to have no real effect on
the intermolecular forces compared to what you expect for a
non–polar molecule?
- Again it should be emphasised that
permanent dipoles like
δ+C=Oδ– and
δ–O–Hδ+ obviously contribute an extra
component to the intermolecular forces.
- BUT it should ALSO be emphasised
that if these dipoles are in these 'bigger' molecules most of the intermolecular force
still originates from transient dipole – induced dipole attraction if the
other bonds are relatively non–polar.
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ignore adverts at top
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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