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Theoretical–Physical Advanced Level Chemistry

8.2.2 Intermolecular forces and the boiling points of selected organic molecules

A comparison of the boiling points of a series of selected organic compounds whose molecules have similar molecular mass and a similar number of electrons in the molecule. The role–effect of the intermolecular forces (intermolecular bonding) involved and the their effect on the boiling point is explained and discussed on a comparative basis. The molecules in question are 1. butane (alkane), 2. methoxyethane (ether), 3. chloroethane (halogenoalkane/haloalkane), 4. 1–aminopropane (n–propylamine, primary aliphatic amine), 5. propanone (ketone), 6. propan–1–ol (1–propanol, alcohol), 7. ethanoic acid (carboxylic acid) and 8. ethanamide (acid/acyl amide).

(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


8.2 Survey of 8 selected organic molecules – their boiling points and intermolecular forces contd.

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

8.2.2 Detailed Discussion of the eight individual molecules representing eight homologous series

  1. ALKANE 156sg butane
    • CH3CH2CH2CH3, Mr = 58 and 34 electrons
    • Mpt –138oC, bpt –0.5oC, ΔHvap = 22 kJ mol–1, sat'd pvap = 211316 Pa/1585 mmHg at 20oC when liquified under pressure.
    • Alkanes are non–polar molecules where the only intermolecular force operating is the weakest possible, that is the instantaneous dipole – induced dipole intermolecular forces. These are sometimes called London–dispersion forces and occur between ALL molecules, even single atoms of the noble gases. Van der Waals forces include all types of intermolecular forces which are not due to an actual chemical bond BUT sometimes this name is used just to mean these instantaneous dipole – induced dipole dispersive forces (sorry but it can be confusing!).
      • The electronegativities are: C (2.5) and H (2.1) and produces a virtually non–polar bond and any very small effects will tend to cancel out e.g. H–C–H situations and so alkanes are the least polar organic molecules i.e. as near non–polar molecules you will get.
    • These electrical attractive forces act between ANY atoms or molecules and is primarily a function of the number of electrons in the molecule, though their spatial distribution can be significant.
    • The larger the molecule, i.e. the greater the number of electrons in it, the more polarizable it is and so the attractive force is greater.
    • The force arises from the instantaneous and random asymmetry of the electron fields in the atomic orbitals because of the random behaviour of electrons in the atomic or molecular orbitals.
    • A transient δ+ in one molecule induces a transient δ– in a neighbouring molecule, so causing a very weak and transient electrical attraction.
      • Note that these partial charges are shown as a delta + (δ+) or a delta – (δ–) and they are tiny charges compared to a full single plus charge e.g. on an Na+ sodium ion or a full single minus charge  on a Cl chloride ion.
    • This polarisation can readily occur when particles collide with each other e.g. in liquids or vibrate against each other e.g. in a solid. In this situation electron clouds from neighbouring atoms/molecule will repel each other and the distortion of the charge distribution causes the polarization. Under these circumstances, contact between any two atoms/molecules can produce temporary or transient polarisation.
      • attractions
    • Total intermolecular force = 100% (instantaneous dipole – induced dipole force)
    • This is why, in doing comparisons, you should choose molecules of similar molecular mass, and, in particular, the same total number of electros to give a 'base–line' of comparable effects.
    • You can then judge the effects of changing molecular structure e.g. polar bonds increasing the inter–molecular attractive forces between polar molecules with a permanent dipole.
    • In examples 2 to 8, all the molecules will exhibit the effects of the 32–34 electrons in the molecule in terms of instantaneous dipole – induced dipole plus an extra intermolecular force effect due to polar bonds, and in these descriptions it will now be assumed you are aware that instantaneous dipole – induced dipole attractions are ever present!
  2. TOP and LINKSETHER alcohols and ether structure and naming (c) doc b methoxyethane (methyl ethyl ether, 'ether')
    • CH3CH2OCH3, Mr = 60 and 34 electrons.
    • Mpt –139oC, bpt 7oC, ΔHvap =  21 kJ mol–1, sat'd pvap =  160000 Pa/1216 mmHg at 20oC when liquified under pressure.
    • Although there are polar C–O bonds (electronegativities: O 3.5 > C 2.5), δ+C–δOδ–Cδ+, the C–O–C ether linkage means each polar bond cancels the other out so there is only a small polarity effect in the molecule due to the C–O–C bond angle of 109o. The boiling point is only a little above that of non–polar alkanes and a little below halogenoalkanes.
    • Methoxyethane is isomeric with propan–2–ol which has much greater boiling point due to hydrogen bonding.
    • So, for methoxymethane, almost all the intermolecular attraction arises from instantaneous dipoles – induced dipoles.
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)
      • A theoretical figure quoted for ethoxyethane, a similar molecule, for the total intermolecular force was
      • (86.5%  instantaneous  dipole – induced dipole) + (7.4% permanent dipole – permanent dipole) + (4.4% permanent dipole – induced dipole))
      • (86.5% dispersive forces) + (7.4% Keesom forces) + (4.4% Debye forces)
  3. TOP and LINKSHALOGENOALKANE (haloalkanes) Image570b chloroethane (ethyl chloride)
    • CH3CH2Cl, Mr = 64.5 and 34 electrons.
    • Mpt –136oC, bpt 12.5oC, ΔHvap =  25 kJ mol–1, sat'd pvap =  133322 Pa/1000 mmHg at 20oC when liquified under pressure.
    • Halogenoalkanes have a weakly polar Cδ+–Xδ bond (X = halogen) due to the difference in electronegativities (Pauling values) of carbon and halogens, e.g. Cl(3.0) > C(2.5) giving Cδ+–Clδ.
    • This gives rise to a weak, but permanent dipole, hence the extra permanent dipole – permanent dipole intermolecular attractive forces raising the bpt and ΔHvap and lowering the vapour pressure compared to butane.
    • BUT the effect is quite small, so, for chloroethane, despite the C–Cl polar bond, almost all the intermolecular attraction arises from instantaneous dipoles – induced dipoles.
      • attractions
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)
    • One reason why the polar bond, the origin of the permanent dipole, doesn't have as greater effect for a molecule as the same number of electrons as butane is that 8 of the electrons are tightly held in the 2nd inner shell and might not be as polarisable?
  4. TOP and LINKSPRIMARY ALIPHATIC AMINE (c) doc b 1–aminopropane (n–propylamine)
    • CH3CH2CH2NH2, Mr =  59 and 34 electrons.

    • Mpt oC, bpt 48oC, ΔHvap =  30 kJ mol–1, sat'd pvap =  33060 Pa/248 mmHg at 20oC.
    • Amines are permanently polar molecules because of the polarised N–H bond giving rise to hydrogen bonding.
    • Electronegativities N (3.5) > H (2.1), giving the polar bond δN–Hδ+ (permanent dipole).
    • It appears that the tiny partially positive proton can get relatively close to the lone pair of electrons on the very electronegative nitrogen atom and produce the strongest possible permanent dipole – permanent dipole attraction and this special case is called 'hydrogen bonding'.
    • This only happens, bar a few exceptions, with the three most electronegative elements namely, nitrogen, oxygen and fluorine.
      • δN–Hδ+  δO–Hδ+  δ+H–Fδ– 
    • Hydrogen bonding is the strongest type of intermolecular force arising from permanent dipole – permanent dipole attraction and the partially positive proton end of one molecule attracts the partially negative nitrogen (or oxygen/fluorine) lone pair end of a close neighbouring molecule.
    •  llll δN–Hδ+llll δN–Hδ+llll δN–Hδ+  etc. (a hydrogen bond is often denoted by llll)
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)
    • On a delta(+) ... delta(–) charge basis, a highly polar bond outweighs an individual instantaneous – induced dipole BUT there may several temporary dipoles occurring at the same time, particularly in a larger molecule of many electrons, so the permanently polar bond dipole is in most cases an extra contribution to the total intermolecular force and not necessarily the dominant contributor..
    • but the hydrogen bonding effect is not as great as for alcohols probably because oxygen is more electronegative than nitrogen and there are two lone pairs of electrons on the oxygen atom
    • More notes on 'hydrogen bonding' in section 8.6
  5. TOP and LINKSKETONE Image1901 propanone (acetone)
    • CH3COCH3, Mr = 58 and 32 electrons.
    • Mpt –95oC, bpt 56oC, ΔHvap =  29 kJ mol–1, sat'd pvap = 24131 Pa/181 mmHg at 20oC.
    • Ketones are permanently polarised molecules due to the highly polar bond δ+C=Oδ caused by the difference in electronegativities of carbon and oxygen i.e. O(3.5) > C(2.5).
    • This causes the permanent dipole – permanent dipole interaction between neighbouring polar molecules
    • via δ+C=Oδ....δ+C=Oδ– interactions further illustrated below
    • extra attractions
    • The vapour pressure is reduced, and the bpt. and ΔHvap increased compared to butane because there is a 2nd intermolecular force operating.
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole) + (permanent dipole – induced dipole)
      • Theoretical data from the web quoted for the total intermolecular force
      • (14.2%  instantaneous  dipole – induced dipole force) + (78.4% permanent dipole – permanent dipole) + (7.4% permanent dipole – induced dipole)
    • The effect is also greater than for chloroethane because the electronegativity difference between C and O is greater than C and Cl, giving a more polar molecule, so stronger permanent dipole – permanent dipole intermolecular attractive forces.
    • Note that you can't get hydrogen bonding here because no H attached to a very electronegative atom O.
    • Aldehydes isomeric with ketones will have a similar boiling point i.e. the similar intermolecular forces (instantaneous dipole – induced dipole plus permanent dipole – permanent dipole intermolecular forces e.g.
      • aldehydes and ketones nomenclature (c) doc b CH3CH2CHO, propanal (propionaldehyde), Mr = 58 and 32 electrons.
      • Mpt –81oC, bpt 49oC, ΔHvap =  31.5 kJ mol–1, sat'd pvap = 31300 Pa/235 mmHg at 20oC.
  6. TOP and LINKSALCOHOL Image248 propan–1–ol (1–propanol, propyl alcohol)
    • CH3CH2CH2OH, Mr = 60 and 34 electrons.
    • Mpt –127oC, bpt 97oC, ΔHvap = 45 kJ mol–1, sat'd pvap = 1933 Pa/14.5 mmHg at 20oC.
    • Alcohols are permanently polarised molecule due to the highly polar bond δO–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
      • δ:O–Hδ+ llll δ:O–Hδ+ ... etc.    (llll represents the directional hydrogen bond)
      • 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 which is effectively the delta minus and this should come out in a full diagram showing the hydrogen bonding between molecules.
      • extra attractions
    • So called hydrogen bonding is the strongest of the permanent dipole – permanent dipole intermolecular forces, but it is NOT a true ionic or covalent chemical bond in the sense that electrons are not transferred or shared to form the bond.
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)
      • A theoretical figure quoted for 2–methylpropan–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 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!).
    • The vapour pressure is reduced, and the bpt. and ΔHvap considerably increased compared to butane because of this 2nd intermolecular force operating – hydrogen bonding.
    • The effect is also greater than for the polar chloroethane/propanone for two reasons
      • (i) the electronegativity difference between O and H is greater than for  C and Cl or C and O giving a more polar bond
      • (ii) the small size of the hydrogen atom of one molecule allows it to get nearer to the oxygen of a neighbouring molecule increasing the strength of this particular intermolecular force.
    • Propan–2–ol is isomeric with methoxyethane which has a much lower boiling point due to lack of hydrogen bonding.
    • More on hydrogen bonding
  7. TOP and LINKSCARBOXYLIC ACID Image1672 ethanoic acid (acetic acid)
    • CH3COOH, Mr = 60 and 32 electrons.
    • Mpt 17oC, bpt 118oC, ΔHvap = 58 kJ mol–1, sat'd pvap = 1520 Pa/11.4 mmHg at 20oC.
    • Of the five compared organic molecules, ethanoic acid has the highest bpt and ΔHvap and lowest vapour pressure at 20oC.
    • This is because as well as the instantaneous dipole – induced dipole attraction, there are two sources of permanent dipole – permanent dipole interactions.
      • Permanent dipole – permanent dipole interaction between neighbouring polar molecules via the carbonyl group δ+C=Oδ....δ+C=Oδ– (see propanone above for more details),
      • and the even stronger hydrogen bonding via the hydroxy group δO–Hδ+ llll δO–Hδ+ (see propan–1–ol above for more details) ...
      • BUT there are two principal structures which can arise from hydrogen bonding via
      • the δ+C=Oδ llll δ+H–Oδ– linking molecules of ethanoic acid together, a linear dimer (below) or a cyclic dimer (further down) and the latter is the more predominant hydrogen bonding interaction.
      • extra attractions
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)
    • I've recently found a research paper from 1999 that liquid ethanoic acid consists of 'polymers' of ethanoic acid held together by hydrogen bonding via δ+C=Oδ llll δ+H–Oδ– δ+C=Oδ llll δ+H–Oδ– etc. linkages. They may be linear dimers or longer species.
    • Then I was sent another research paper from 2001 which claimed the principal species in liquid ethanoic acid is the cyclic dimer shown below.
    • Whatever species exist in liquid ethanoic acid, and its probably a mixture of all the species mentioned above, quite simply, in boiling ethanoic acid you are effectively vaporising a larger molecule prior to vapourisation and therefore more energy needed to vapourise it i.e. a higher boiling point where the particles at a higher temperature have enough kinetic energy to 'escape' the intermolecular forces at the surface..
    • Also, whatever the structure of liquid ethanoic acid at the molecular level, hydrogen bonding is primarily responsible for the creation of a larger molecular species of (CH3COOH)n, with n>1.
    • In the cyclic dimer structure, the hydrogen bonds create a larger molecule with a total of 64 electrons and this alone would considerably increase the boiling point of ethanoic acid.
    • As an 'electron' point of comparison, octan–1–ene (1–octene), C8H16, has 64 electrons and boils at 121oC close to the 118oC bpt of ethanoic acid.
    • There is plenty of evidence from vapour pressure/density measurement that dimer does exist as a high % mole fraction of ethanoic acid vapour as well as in the liquid.
    • More on hydrogen bonding
  8. TOP and LINKSACID AMIDE  (c) doc b ethanamide (acetamide)
    • CH3CONH2, Mr =59 and 32 electrons.
    • Mpt 82oC, bpt 221oC, ΔHvap = 46 kJ mol–1, sat'd pvap = ? Pa/? mmHg but very low at 20oC since it is solid.
    • Permanent dipole – permanent dipole interaction between neighbouring polar molecules via the carbonyl group δ+C=Oδ....δ+C=Oδ– (see propanone above for more details) ..
    • ... and ethanamide is even higher boiling than ethanoic acid because there are more hydrogen bonds possible.
    • There are two sites on the molecule for hydrogen bonding
    • (i) The –NH2 group i.e.  δN–Hδ+ llll δN–Hδ+ as well as the hydrogen bonding due to
    • (ii) δ+C=Oδ llll δ+H–Nδ– interactions, the latter (ii) are probably the dominant inter molecular force if ethanoic acid provides a model.
    • Total intermolecular force = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole including hydrogen bonding) + (permanent dipole – induced dipole)
      • A theoretical figure quoted for methanamide (formamide), a similar molecule, for the total intermolecular force was
      • (3.2%  instantaneous  dipole – induced dipole) + (93.3% permanent dipole – induced dipole including H bonding) + (3.5% permanent dipole – induced dipole)
      • This is one of the very few molecules where the majority of the intermolecular force originates from hydrogen bonding.
    • I don't know if acyl amides can form polymers or dimers via hydrogen bonding in the way that carboxylic acids do?
    • Its worth noting, in view of the high boiling point of ethanamide, a reflection of the strong intermolecular forces, that the δ+C=Oδ llll δ+H–Nδ– interaction is the most important hydrogen bond that holds together the secondary structure of proteins to give the sheet and helical structures AND about 50% of the hydrogen bonds that hold together the single helix structure of RNA molecules and the double helix structure of DNA molecules.
    • Its interesting here that the boiling point of ethanamide is raised another ~100o compared to the hydrogen bonded ethanoic acid. I'm not quite sure why? Can more hydrogen bonds be formed? Are much larger polymeric species formed in liquid ethanamide? Not sure on this one but methanamide behaves in the same way with respect to methanoic acid!

TOP and LINKSSummary table of the 8 organic molecules discussed above plus some others

In the table the following abbreviations for the different contributory Van der Waals intermolecular forces: Ins = instantaneous dipole – induced dipole attraction, WP = weaker permanent dipole – permanent dipole attraction, SP stronger permanent dipole – permanent dipole attraction, HB = hydrogen bonding attraction, MHB multiple hydrogen bonding attraction sites on the molecule, D = Debye dipole moment  units

8.4 Table 1b.  REPEAT OF SUMMARY of 1. to 8. and  * other molecular data to do more comparisons

MOLECULE formula Mr electrons bpt ΔHvap/kJmol–1 Dipole moment/D Intermolecular forces
1. butane CH3CH2CH2CH3 58 34 272.5K/–0.5oC 22 0.00 Ins
* cyclobutane cyclo C4H8 56 34 286K/13oC ? ? Ins
* but–1–ene CH3CH2CH=CH2 56 32 267K/–6oC ? ? Ins
* but–2–ene CH3CH=CHCH3 56 32 3.7oC(cis) 1oC(trans) ? ? Ins
* buta–1,3–diene CH3CH=CH=CH2 56 32 269K/–4oC ? ? Ins
* but–1–yne CH3CH2C≡CH 54 30 281K/8oC ? ? Ins
* but–2–yne CH3C≡CHCH3 54 30 300K/27oC ? ? Ins
2. methoxyethane CH3OCH2CH3 60 34 280K/7oC 21 1.23 Ins, WP
* 1–fluoropropane CH3CH2CH2F 62 34 270.5K/–2.5oC ? Ins Ins, WP
3. chloroethane CH3CH2Cl 64.5 34 285.5K/12.5oC 25 2.06 Ins, WP
* methyl methanoate HCOOCH3 60 32 304K/31oC ? ? Ins, WP?
4. propylamine CH3CH2CH2NH2 59 34 321K/48oC 30 1.17 Ins, HB
5. propanone CH3COCH3 58 32 329K/56oC 29 2.88 Ins, SP
* propanal CH3CH2CHO 58 32   ? ? Ins, SP
6. propan–1–ol CH3CH2CH2OH 60 34 370K/97oC 45 1.69 Ins, HB
* propan–2–ol CH3CHOHCH3 60 34 355K/82oC ? ? Ins, HB
* propanenitrile CH3CH2C≡N 55 30 370K/97oC ? ? Ins, SP?
7. ethanoic acid CH3COOH 60 32 391K/118oC 58 1.74 Ins, SP, MHB
8. ethanamide CH3CONH2 59 32 494K/221oC 46 3.60 Ins, SP, MHB

Further discussion points

  • 1–fluoropropane was added for comparison showing that the highly polar C–F bond does not appear to increase the intermolecular force for the same chain length (counting an F equal to a C) and same number of electrons and I'm not sure how to explain this?

In section 8.4 I've discussed more examples of comparative boiling points, intermolecular forces and number of electrons in the molecule which I hope will be of interest to the more 'inquisitive' students and teachers.

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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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