Doc 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. K_c and K_p 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

• 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.

•  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–H^δ+llll^δ–:N–H^δ+.... 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–H^δ+ llll ^δ–:F–H^δ+.... etc.
    • More on hydrogen bonding
  • Water, H₂O

    • Intermolecular hydrogen bonding in water

    • On the basis of molecular mass and electron number (water < 1/3rd of examples 1–8), the data shows that the bpt and ΔH_vap 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–H^δ+llll^δ–:O–H^δ+.... 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

    • An important exam note:

      • You must clearly show the directional linearity of the X^δ--H^δ+ǁǁǁ:X^δ- arrangement of the hydrogen bond including the single X-H covalent bond and the lone pair on the other X atom too! (X is usually O, N or F)

      • You must do this accurately in exams when drawing intermolecular bonding diagrams of water or alcohols because it is the only spatially directed intermolecular force, all the rest of the other types of intermolecular bonding forces are randomised - the δ+ and δ- electric fields acting in all directions.

• 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 60^o 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 97^o 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–H^δ+ we have the  ^δ–O–H^δ+llll^δ–:O–H^δ+  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–llll^δ+H₂NH–NH₂^δ+llll ^δ–:N–H...

• 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 ~30–40^o 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=O^δ– and ^δ–O–H^δ+ 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.
  • You can see this by comparing the differences in the 24–26 electron species with the 56–58 electron species.
  • 8.2.1 A summary of Van der Waals forces and an introduction to intermolecular forces

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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. K_c and K_p 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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