Revision notes explaining Van der Waals intermolecular forces - comparing types of intermolecular bonding

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Doc Brown's Chemistry Advanced A Level Notes - Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 8

8.2.1 Intermolecular Forces - Intermolecular Bonding – Van der Waals forces and the boiling point comparison of 8 organic molecules of similar molecular mass

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The different types of intermolecular force (intermolecular bond) are described, explained and discussed with examples

 i.e. Instantaneous dipole – induced dipole interaction (London forces, dispersive/dispersion forces), permanent dipole – permanent dipole interactions (Keesom forces, orientation forces), Permanent dipole – induced dipole interactions (Debye forces, induction forces)

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

8.2.1 Introduction  to intermolecular forces – Van der Waals forces

also referred to as 'intermolecular bonding' forces

(do NOT confuse with chemical bonding between atoms or ions i.e. so-called ionic, covalent or metallic bonds)

A definition of Van der Waals forces

These can be defined as weak, short-range electrostatic attractive forces between uncharged molecules, arising from the interaction of permanent or transient electric dipole moments and the different types and their origin are described below.

  • From the start understand that:

  • Intermolecular forces are all about partially positive (δ+) sites and partially negative (δ) sites on molecules causing the attraction between neighbouring molecules - though their origin can differ.

  • The fact that molecules congregate together to form liquids and solids suggests that there must be attractive forces between the molecules independently from the intramolecular bonds which hold the atoms together in the molecule.

  • The origin of each source of intermolecular force is summarised below and discussed further for particular molecules.

  • In the context of this page, the word dipole means an asymmetric distribution of electron electrical charge to give partially positive (δ+) and partially negative (δ) regions in the same molecule.

  • In a simple sense its a molecule with a partially positive end and a partial negative charge at the other end.

  • Electric dipoles (δ+ and δ) may be permanent or transient (temporary) and the molecules discussed here are electrically neutral  overall.

  • There are always attractive forces operating between ANY particles whatever their particle constitution in gases, liquids or solids composed of atoms, ions or molecules.

    • They are referred to as intermolecular attractive forces or intermolecular bonding.

    • Collectively they are often referred to as Van der Waals forces.

    • DO NOT confuse intermolecular bonds with the very much stronger intramolecular bonds e.g. between atoms in a molecule like the O-H bond holding atoms together in water, or the C-C and C-H bonds holding atoms together in hydrocarbon molecules.

  • The total intermolecular force is quoted as a summation of the various possible dipoles interaction and the principal attractive forces are shown in bold for selected molecules.

  • Wherever possible, albeit just for a few cases, I've quoted % contributions from the three types of intermolecular attractive force that I've been able to obtain from internet searches or textbooks and if I couldn't match the molecule then I may quote percentages for a similar molecule.

  • One source used by writers of research papers is A. L. McClellan, Tables of Experimental Dipole Moments.

  • In pre–university advanced chemistry exams I suggest you use the terms in bold to describe the intermolecular force component

  • Summary of the types of intermolecular bonding forces (Van Waals forces)

    • Instantaneous dipole – induced dipole interaction

      • Also called London forces or dispersive or dispersion forces.

      • The electrons of an atom behave in a random way within the spatial region they occupy for their specific quantum level e.g. in 3s, 2p, 3d atomic orbitals or a bonding molecular orbital.

      • At any given instant in time the electron cloud will randomly distorted, giving rise to a dipole of partial charges which then induces a dipole in a neighbouring molecule.

        • 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.
      • The random partial positive charge of one dipole will attract the partial negative in the neighbouring molecule or vice versa

      • So even with a completely non–polar hydrocarbon molecule (i.e. a molecule with no significant polar bonds like alkanes and alkenes) there are still intermolecular attractive forces.

      • 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
      • Instantaneous dipole – induced dipole interaction increase the more electrons in the molecule.

        • The larger the molecule, i.e. the greater the number of electrons in it, the more polarizable it is and the greater the chance of a random instantaneous dipole occurring to induce a dipole in a neighbouring molecule, so increasing the intermolecular attractive forces.
        • A good example is illustrated by the boiling point plots for various organic homologous series in section 8.3 where the addition of every non-polar –CH2 unit in the carbon chain produces a corresponding incremental rise in the boiling point due to the incremental rise in intermolecular forces (instantaneous dipole – induced dipole).

      • For most molecules, this is the dominating contribution to the total intermolecular force, but the presence of polar bonds can add a significant contribution to this and the consequential affects on the properties of the molecule.

      • Comparing the boiling points and intermolecular forces operating between molecules with a similar number of electrons does provide important insights into their molecular behaviour.

      • However, you should be aware that the way the electrons are distributed, both in terms of their electronic energy levels, and their spatial distribution, can have significant effects on the strength of instantaneous dipole – induced dipole forces. You will see this particularly in the case studies of section 8.4.

      • Notes on instantaneous dipole - induced dipole forces:

        • (i) You come across other words other than instantaneous e.g. 'temporary', 'transient' or even 'induced' – induced dipole attractions.

        • (ii) The molecule does not have to be polar for this force to exist.

        • (iii) The same force exists between ANY neighbouring molecules, whether they are the same, different, polar or non-polar. It is a universal intermolecular force.

      • See comparing organic molecule boiling points and homologous series comparison

      • and other case studies of boiling points related to intermolecular forces

    • Permanent dipole – permanent dipole interactions

      • Also called Keesom forces or orientation forces.

      • If two atoms constituting a bond have significantly different electronegativities, the bond will be permanently polar and produce a permanently polar molecule.

      • Such molecules posses what is known as dipole moment.

      • Therefore, as result of this permanent dipole, these permanently polarised molecules will attract neighbouring molecules because of this dipole moment as well as the attraction due to instantaneous dipole – induced dipole.

      • e.g.

      • attractions seem to make little difference to the bpt!
      • or in carbonyl compounds δ+C=Oδ....δ+C=Oδ– in organic carbonyl compounds e.g. aldehydes, ketones and carboxylic acids.

      • Trichloromethane (CHCl3) is another polar molecule δ-Cl3Cδ+H


      • There is special sub–category of permanent dipole – permanent dipole interactions called hydrogen bonding.

      • It is a spatially directed permanent dipole - permanent dipole attractive force - the only intermolecular force that is spatially and specifically directional.

      • Hydrogen bonding only usually occurs when the three most electronegative elements (N, O and F) are covalently bonded to a hydrogen atom (intramolecular) AND bonded to a similar neighbouring molecule and an intermolecular bonding force. (more on hydrogen bonding in section 8.6).

      • In these molecules you get one of the following three very polar bonds:–

      • (i)  δ:N–Hδ+ e.g. in ammonia NH3, amines R-NH2, amides RCONH2  (R = alkyl or aryl)

    • diagram of intermolecular hydrogen bonding forces between liquid water molecules doc brown A level chemistry revision notes diagram of intermolecular hydrogen bonding forces between liquid alcohol molecules doc brown A level organic chemistry revision notes

      • (ii)  δ:O–Hδ+ e.g. in water H2O (above), alcohols ROH (above), carboxylic acids RCOOH  (R = alkyl or aryl)

      • (iii)  δ+H–Fδ– in hydrogen fluoride HF

      • and via these highly polar bonds you get molecule to molecule attraction via so called hydrogen bonding.

      • Note the spatially important non-bonding pairs of electrons (:) on the most electronegative atom.

      • These are the strongest permanent dipole permanent dipole intermolecular forces

      • e.g. using llll to indicate a hydrogen bond

      •  δO–Hδ+llllδ:O–Hδ+llllδ:O–Hδ+llllδ:O–Hδ+llll  in water (liquid or solid ice)

      • llll δN–Hδ+llllδ:N–Hδ+llllδ:N–Hδ+  in amines or liquid ammonia,

      •  in the case of carboxylic acids the dominant interaction is the hydrogen bonding via

        • δ+C=O:δllllδ+H–Oδ– from molecule to molecule.

      • In hydrogen fluoride, in all physical states you get a 'chain connection' llll δF–Hδ+llllδF–Hδ+llllδF–Hδ+

      • You can also get hydrogen bonding between these different molecules e.g.

        • δN–Hδ+llllδ:O–Hδ+ or δO–Hδ+llllδ:N–Hδ+ in aqueous ammonia solution (NH3(aq),

        • C-δO–Hδ+llllδ:O–Hδ+ in aqueous carboxylic acid solutions (RCOOH(aq)),

        • δF–Hδ+llllδ:O–Hδ+ or  δO–Hδ+llllδ:F–Hδ+ in hydrofluoric acid solution (HF(aq)).

      • A slightly different and modestly intriguing case of hydrogen bonding!

        • Trichloromethane (CHCl3) and propanone (CH3COCH3) are both polar molecules, but do not hydrogen bond with themselves.

        • BUT, if you mix the two liquids, they readily dissolve in each other via hydrogen bonding!

        • Cl3C-Hδ+llllδ-:O=C(CH3)2

        • Note that the three electronegative chlorine atoms have such an effect on the carbon atom (δ+), that the hydrogen atom also acquires a sufficient (δ+) to hydrogen bond with the oxygen atom (via a lone pair of electrons).

    • An important exam note:

      • You must clearly show the directional linearity of the Xδ--Hδ+ǁǁǁ:Xδ- 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.

      • The spatial directional nature of the hydrogen bond is very important when studying e.g. the crystalline structure of ice or the double helix of DNA - the latter is held together by base pair hydrogen bonds.

    • For detailed case studies see my main hydrogen bonding page

    • Permanent dipole – induced dipole interactions

      • Also called Debye forces or induction forces.

      • The permanently polar bond in one molecule can induce a dipole in a neighbouring molecule, whether the other molecule is polar or non–polar, it makes no difference, induction happens!

      • e.g. attractions

      • This applies to any pure polar molecule or in a mixture of a polar and non-polar molecules as illustrated above.

      • This tends to be the minority contribution to a molecule's total intermolecular bonding forces.

  • The discussions below primarily focus on intermolecular forces 1. and 2. and 3. is usually less significant and gets a brief mention.

  • Data abbreviations used: Mpt = melting point, Bpt = boiling point (given in Kelvin and Celsius)

  • Here in section 8.2 examples are chosen from the eight homologous series listed below and all the molecules have 32–34 electrons and four larger C, N or O atoms plus hydrogen (except chloroethane C + C + Cl, but still has 34 electrons).

  • See comparing organic molecule boiling points and homologous series comparison

  • and other case studies of boiling points related to intermolecular forces

The survey and a preliminary summary table

hopefully justified by the arguments outlined after the table and in on a separate page in section 8.2.2

  • Ins = instantaneous (temporary) dipole – induced dipole attraction (a sort of baseline force since it applies to all molecules, in fact it operates between ANY adjacent particles - atoms, ions or molecules).

  • WP = weaker permanent dipole – permanent dipole attraction (doesn't seem to have much effect on the boiling point)

  • SP stronger permanent dipole – permanent dipole attraction (NOT H bonding, but has a definite effect on the boiling point)

  • HB = hydrogen bonding attraction - the strongest permanent dipole – permanent dipole attractive force, i.e. the strongest SP and has the largest effect on the boiling point)

  • MHB multiple hydrogen bonding attraction sites on the molecule (i.e. where there are at least two 'functional' groups capable of two permanent dipole – permanent dipole interactions including hydrogen bonding, hence producing an even bigger effect on raising the boiling point)

  • Note that permanent dipole – induced dipole attractive forces are not mentioned much and generally only contribute a small portion of the total intermolecular force.

  • Also, where I can obtain data, I've indicated the percentage contribution of the three types of intermolecular attraction which contribute to the total intermolecular force i.e. the % contributions to Van der Waals force.

  • D = Debye dipole moment  units

8.4 Table 1a. Comparing 32–34 electron species – linear organic molecules (4 C/O/N atoms)

1.to 8. are discussed in detail on a separate page

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
2. methoxyethane CH3OCH2CH3 60 34 280K/7oC 21 1.23 Ins, WP
3. chloroethane CH3CH2Cl 64.5 34 285.5K/12.5oC 25 2.06 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
6. propan–1–ol CH3CH2CH2OH 60 34 370K/97oC 45 1.69 Ins, HB
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
  • As you go down the table the boiling point increases clearly reflecting the increasingly strong intermolecular forces operating, all of which have been individually discussed above.
  • The table shows that increasingly polar bonds will increase the intermolecular forces between polar molecules and raise the boiling point compared to a non–polar molecule of similar size and number of electrons and in particular where hydrogen bonding occurs – the strongest of the permanent dipole – permanent dipole interactions.
  • It should however, be pointed out, that in most cases, most of the intermolecular attractive force originates from the transient instantaneous dipole – induced dipole interactions, but the extra effect of polar bonds is significant in discussing and accounting for differences in physical properties such as melting points and boiling points, and also solubility (see section 8.7 on solubility)
  • The enthalpies of vaporisation do not show as clearer a pattern, though the higher values are in bottom half of the table.
  • Each molecule 1. to 8.  is discussed in detail on the next page.
  • See comparing organic molecule boiling points and homologous series comparison

  • and other case studies of boiling points related to intermolecular forces



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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