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Theoretical–Physical Advanced Level Chemistry – Equilibria

 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.

 

(c) doc b KS4 Science GCSE/IGCSE Notes on reversible reactions and chemical equilibrium

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 that 8 organic molecules are discussed in detail in section 8.2.1 and continued in section 8.2.2

  • What contributions do each of the 'sources' of intermolecular force make?

    • 8.2.1 A summary of Van der Waals forces and an introduction to intermolecular forces

    • Although some data has been quoted in section 8.2, its not an easy question to answer! probably because there isn't an easy answer! or one 'simple' pattern.

    • Like many situations in chemistry, there are multiple factors involved and one or factors may be the dominant one!

    • If you look at the boiling point plots in section 8.3 (particularly the electrons versus Bpt plots). you can see that the effect of a permanent dipole becomes increasing reduced i.e. the instantaneous dipole – instantaneous dipole interactions far outweigh the permanent dipole  – permanent dipole interactions for 'longish' molecules.

    • In fact, all the plot lines for the various homologous series of non–polar or polar molecules are actually converging!

    • BUT, if you look at smaller molecules with a permanent dipole, the permanent dipole – permanent dipole interaction, particularly those involving hydrogen bonding may well be the dominant source of the intermolecular attractive force AND it is these sort of molecules which are discussed the most in A level chemistry courses in the context of Bpt versus intermolecular forces.

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+ llll δ: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
      • On the basis of molecular mass and electron number (water < 1/3rd of examples 1–8), the data shows that the bpt and ΔHvap of water are very much higher, and the vapour pressure much lower than expected, even compared to the polar molecules described above.
      • This is due to water exhibiting the strongest intermolecular force hydrogen–bonding
      • ....δ:O–Hδ+ llll δ:O–Hδ+.... etc.
      • BUT, the two lone pairs on the oxygen and the two available hydrogens attached to the oxygen mean that water can form twice as many hydrogen bonds per molecule than alcohols in the liquid state on a 'time averaged basis'.
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)
      • So what proportion of each source of attraction contributes to the total intermolecular force in the case of water?
        • Theoretical data from the web quoted for the total intermolecular force (average of 2 sets of data found)
        • (10.1%  instantaneous  dipole – induced dipole) + (85.4% permanent dipole – permanent dipole) + (4.5% per. dipole – induced dipole)
      • The average kinetic energies of molecules is proportional to the temperature in K (the Kelvin thermodynamic scale).
      • Since the boiling point of water is more than triple that of methane, and I roughly estimate of the boiling point of water from hydride boiling point graphs assuming no hydrogen bonding to be about 200K (real value is less than double this) ...
      • ... we can we 'crudely' argue that ~50% of the attraction comes from the instantaneous dipole – induced dipole forces (~ CH4 ... CH4 attractive forces) and ~50% due to hydrogen bonding?
      • The enthalpy of vaporization water is exceptionally high for that size of molecule.
      • More on hydrogen bonding
 


 

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 liquified, 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 increasing the boiling point considerably, perhaps by nearly doubling it ...
      • so for CH3δO–Hδ+ we have the  δO–Hδ+ llll δO–Hδ+  situation ...
      • ... and the permanent dipole – permanent dipole attractions of the hydrogen bonding probably 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 llll δ+H2NH–NH2δ+ llll δ:N–H...
 


 

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) ? ?
  • Series 8.4.3 Comparing 24–26 electron species – larger linear organic molecules

    • propane, CH3CH2CH3

      • Propane can act as a baseline for this series since it is a totally non–polar molecule with only instantaneous dipole – induced dipole forces operating.

      • Total intermolecular force = (100% instantaneous dipole – induced dipole)

    • methoxymethane, CH3OCH3

      • The presence of the C–O–C does seem to raise the boiling point a bit, despite the fact the two C–O polar bonds tend to cancel each other out, resulting in a relatively non–polar molecule.

      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • chloromethane, CH3Cl

      • Despite the single polar C–Cl bond this doesn't seem to increase the boiling point much above what get for the ether methoxymethane.

      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • ethanal, CH3CHO

      • However the polar C=O bond in ethanal does increase the intermolecular forces an extra amount to raise the boiling point ~30–60o above what expected for a non–polar molecule.

      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • ethanol, CH3CH2OH

      • Here  there is an even greater increase in boiling point due to hydrogen bonding.

      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)

      • A theoretical figure quoted for 2–methylprpan–2–ol (tert–butanol), a similar molecule, for the total intermolecular force was: (67.2%  instantaneous  dipole – induced dipole) + (23.1% permanent dipole – induced dipole including H bonding) + (9.7% permanent dipole – induced dipole)

      • Note that the hydrogen bonding only contributes ~1/4 of the total intermolecular force (Van der Waals forces) but this is sufficient to raise the boiling point significantly above that of a non–polar molecule of similar molecular mass (particularly if similar number of electrons).

      • However in the case of ethanol you would expect a greater contribution from the permanent dipole – permanent dipole interactions (including the hydrogen bonding).

      • The quotation for methanol is: (61.3%  instantaneous  dipole – induced dipole) + (30.3% permanent dipole – induced dipole including H bonding) + (8.4% permanent dipole – induced dipole)

      • This would suggest that the contributions to the intermolecular forces for ethanol are approximately: (~64%  instantaneous  dipole – induced dipole) + (~27% permanent dipole – induced dipole including H bonding) + (~9% permanent dipole – induced dipole)

      • Textbooks tend to indicate that hydrogen bonding is the predominant intermolecular force in molecules such as alcohols, but this is very rarely the case (only a few other molecules like methanol (methyl alcohol), methanamide (formamide) and ethanamide (acetamide) really), for the vast majority of molecules the predominant intermolecular attraction is the instantaneous dipole – induced dipole interaction (the weakest of the forces!).

    • methanoic acid, HCOOH

      • For carboxylic acids the two permanent dipoles C=O and O–H, and both sites available for hydrogen bonding means the boiling point is raised even further and linear or cyclic dimers formed. (see ethanoic acid)

      • Here  there is an even greater increase in boiling point due to hydrogen bonding.

      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)

      • I would estimate that 70–80% of the intermolecular attractive force in methanoic acid comes from the permanent dipole – permanent dipole attractions including hydrogen bonding.

    • methanamide, HCONH2

      • Like carboxylic acids, amides have two permanent dipoles, in this case C=O and N–H, and both sites available for hydrogen bonding means the boiling point is raised even further. (see ethanoic acid)

      • So again there is an even greater increase in boiling point due to hydrogen bonding.

      • Total intermolecular force = (3.2% instantaneous dipole – induced dipole) + (93.3% permanent dipole – permanent dipole including hydrogen bonding) + (3.5% permanent dipole – induced dipole)

      • Its interesting here that the boiling point of methanamide is raised another ~100o compared to the hydrogen bonded methanoic acid. I'm not quite sure why? Can more hydrogen bonds be formed? Are much larger polymeric species formed in liquid methanamide?  Not sure on this one but ethanamide behaves in the same way with respect to ethanoic acid!

 


 

 

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
      • Heptane can act as a baseline for this series since it is a totally non–polar molecule with only instantaneous dipole – induced dipole forces operating.
      • Total intermolecular force = (100% instantaneous dipole – induced dipole)

    • 1–fluorohexane, CH3CH2CH2CH2CH2CH2F
      • The polar C–F bond in 1–fluorohexane seems to have little impact on the boiling point, which is slightly below that of heptane (as was the similar case with fluoromethane in Table 2). So again, most of the intermolecular force of attraction originates from temporary dipole – induced dipole interactions.
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • 1–chloropentane, CH3CH2CH2CH2CH2Cl
      • Despite the polar C–Cl bond this doesn't seem to increase the boiling point much above what might be expected for a non–polar molecule.
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • hexanal, CH3CH2CH2CH2CH2CHO
      • However the polar C=O bond in hexanal does increase the intermolecular forces an extra amount to raise the boiling point ~30–40o above what expected for a non–polar molecule.
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • hexan–2–one, CH3COCH2CH2CH2CH3
      • Isomeric with hexanal and with the same C=O permanent dipole for the same chain length, not surprisingly, a similar boiling point is observed similar arguments apply.
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)

    • hexan–1–ol, CH3CH2CH2CH2CH2CH2OH
      • The increase from hydrogen bonding is greater in the case of hexan–1–ol (1–hexanol).
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)

    • pentanoic acid, CH3CH2CH2CH2COOH
      • For carboxylic acids the two permanent dipoles C=O and O–H, and both sites available for hydrogen bonding means the boiling point is raised even further and linear or cyclic dimers formed. (see ethanoic acid)
      • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)

    • 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 δO–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.

 

WHAT NEXT?

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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