(8.2.3) Permanent dipole – permanent dipole intermolecular bonding

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8.2.3 Intermolecular bonding - permanent dipole – permanent dipole attractions (examples of intermolecular Keesom force bonding)

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Sub-index for this page permanent dipole - permanent dipole intermolecular bonding

(a) Introduction to permanent dipole - permanent dipole attractive forces

(b) Examples of molecules exhibiting permanent dipole – permanent dipole intermolecular bonding

(c) Comparing the intermolecular bonding in alkanes and aldehyde (& ketones)

(d) Comparing 2-methylpropane,  methoxyethane (ethyl methyl ether) and propanone (acetone)

Hydrogen bonding is discussed in detail on a separate page

BUT hydrogen bonding has to be mentioned here too, but details are via above link!

INDEX of notes on intermolecular attractive forces (intermolecular bonding)

(a) 8.2.3 Permanent dipole – permanent dipole interactions of intermolecular bonding

(a) Introduction to permanent dipole - permanent dipole attractive forces

  • Also called Keesom forces or orientation forces.

  • BUT remember, this is where polar molecules contribute an extra attractive force to the intermolecular bond, because all molecules create instantaneous dipole - induced dipole attractions (London dispersion forces).

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

    • Cδ+-Clδ-is a polar bond due to chlorine being much more electronegative than carbon.

    • In some case the effects of polar bonds cancel each other out and the molecule is NOT polar overall e.g. unsymmetrical CH3Cl is a polar molecule, but the symmetrical tetrahedrally shaped CCl4 molecule is NOT.

    • Chloromethane has a permanent dipole and so you get an extra contribution to intermolecular bond from the permanent dipole - permanent dipole attractive forces and this applies to ALL polar molecules.

    • It is importantant to realise that these permanent dipole - permanent dipole attractive forces are extra to the instantaneous dipole - induced dipole attractive forces, the latter applies to all particles - so don't forget Van der Waals forces may have several contributions.

  • Such polar molecules posses what is known as a net dipole moment of over zero.

  • 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 attractive force.

    • Any two molecules, where both have a permanent dipole, will orientate themselves to attract each other's opposite charges, hence contributing to the overall intermolecular bond.

    • This can be simply represented as  δ+moleculeδ-••••δ+moleculeδ- attractive interactions and the electric field effect acts in all directions - this distinguishes it from the permanent dipole - permanent dipole force of hydrogen bonding which is the only specifically directed intermolecular force.

  • This permanent dipole - permanent dipole force intermolecular force acts in all directions between neighbouring molecules (But not in the case of the directional permanent dipole - permanent dipole interaction involved in hydrogen bonding described in section (4).
  • A list of accurate Pauling electronegativities is given below
  • element H Si P C S I Br Cl N O F
    electronegativity 2.20 1.90 2.19 2.55 2.58 2.66 2.96 3.16 3.04 3.44 3.98
  • The greater the difference in negativity between the two atoms of a covalent bond, the greater the polarity of the bond, and tend to produce a polar molecule.

(b) 8.2.3 Permanent dipole – permanent dipole interactions of intermolecular bonding

(b) Examples of molecules exhibiting permanent dipole – permanent dipole intermolecular bonding

BUT again, remember, this is an extra contribution to the intermolecular bond, as well the instantaneous dipole - induced dipole attraction (present in all groups of particles).

  • Hydrogen halide molecule

    • e.g. in the polar molecule hydrogen chloride (•••• represents the intermolecular bond)

    • •••• permanent dipole - permanent dipole attractions
    • from Hδ+-Clδ+ polar bond due to the difference in electronegativity between hydrogen (2.20) and chlorine (3.16).
    • These attractions seem to make little difference to the boiling point, see the graph of group 7/17 hydrides on my page on hydrogen bonding, which shows little anomalous behaviour of HCl (unlike HF).
    • It has been calculated that the 'excess' charge on the hydrogen atom is +0.2e- and the 'excess' charge on the chlorine atom is -0.2e- (where e = the arithmetical electrical charge of an electron).
    • This 'dipole' effect is much smaller, but not insignificant, compared to the minimum charge for a full ionic bond of +e and -e for a simple ionic bond like Na+Cl-.
  • -
  • Halogenated alkane molecules
    • •••• permanent dipole - permanent dipole attractions
    • Halogenoalkane molecules have C-X polar bonds where X = a halogen atom, because the halogen atoms are more electronegative than carbon, so you can permanent dipole - permanent dipole attraction between haloalkane molecules.

    • Trichloromethane (CHCl3) is another haloalkane polar molecule that exhibits permanent dipole - permanent dipole intermolecular bonding δ-Cl3Cδ+H•••δ-Cl3Cδ+H attractions

  • -

  • The >C=O in organic chemistry

    • In organic chemistry, carbonyl compounds with the δ+C=Oδ–....δ+C=Oδ– group exhibit permanent dipole - permanent dipole intermolecular bonding e.g. aldehydes, ketones and carboxylic acids.

    • This explains why aldehydes and ketones have higher boiling points than similar sized alkane molecules (see the comparison table in the next section).

    • ••••permanent dipole - permanent dipole attractions
    • Boiling point of propanone is 56oC (polarizability 6.27) compared to similar sized butane with a boiling point of -1oC (polarizability 8.02).
    • All molecules exhibit instantaneous dipole - induced dipole attraction in their intermolecular bonding.
    • Despite the greater polarizability of butane, the extra intermolecular bonding contribution from the permanent dipole - permanent dipole interactions in polar propane outweigh this effect and give propanone a much higher boiling point.

  • -

  • The Group 7/17 hydrogen halides HX (where X = halogen atom, ignoring F)

    • Here you cannot assume that boiling points are governed by bond polarity as the main factor of the intermolecular bond.

    • The boiling points are: HCl 188K;  HBr 206K;  HI 238K,  HAt 277K

    • This list ignores HF, which exhibits hydrogen bonding (complicating matters), whereas the others don't because the difference in H and X electronegativities is not great enough.

    • Therefore the boiling point order for these polar molecules is HAt > HI > HBr > HCl

    • However the H-X bond becomes less polar: HCl > HBr > HI > HAt,

    • and the dipole moments (D) from HCl to HAt are 1.05 > 0.80 > 0.42 > ? (<0.42)

    • However the polarizability order from HCl to HAt is 2.63 > 3.62 > 5.45 > ? (>5.45)

    • So, the boiling point trend here for HCl to HAt is contradictory to the bond polarity order and principally governed by the number of electrons in the molecule and its (polarizability) which is for the series HCl 18 (0.42) < HBr 36 (0.80) < HI 54 (1.05) < HAt 86 (?)

    • Can add this example to 8a-2?

  • -

(c) 8.2.3 Permanent dipole – permanent dipole interactions of intermolecular bonding

(c)  Comparing the intermolecular bonding of alkanes and aldehyde (& ketones, carbonyl compounds)

Here, I've chosen to compare, two pairs of molecules of similar shape and size and all have a formula mass of 86. They all have 6 carbon atoms or 5 carbons plus an oxygen atom.

The data is tabulated for: hexane & pentanal AND 2-methylpentane & 2-methylbutanal.

ΔHvap = enthalpy of vapourisation (kJ/mol), dynamic viscosity of liquid (mPa.s at 20oC).

The dipole moments of alkanes are zero or very small i.e. <0.2 D (Debye units)

The polarizability is measured in 10-24 cm3 units, but just treat it as a relative number.

Name of compound Molecular formula Structural formula e's Dipole moment Boiling point ΔHvap *Polari-zability Vis-cosity
hexane C6H14 CH3(CH2)4CH3 50 ~0.00 D 69oC 28.9 11.63 0.31
pentanal C5H10O CH3(CH2)3CHO 48 1.37 D 103oC 38.0 ~10? 0.60
2-methylpentane C6H14 CH3(CH2)2CH(CH3)2 50 ~0.00 D 60oC 29.1 ~12? 0.29
2-methylbutanal C5H10O CH3CH2CH(CH3)CHO 48 1.72 D 92oC 34.2 ~10? ~0.6?

For each pair of alkane and carbonyl compound you can say ...

The polarizabilities are not that dissimilar and in fact the hydrocarbon molecules are slightly more polarizable, but this has no significant effect on the differences in physical properties.

The difference in boiling point is very similar at 34oC and 32oC, with the aldehydes having the much higher boiling points.

The intermolecular bond for the alkanes is entirely due to the instantaneous dipole - induced dipole forces (London dispersion forces).

London forces instantaneous dipole - induced dipole intermolecular force attraction explained diagram illustrated ••••London forces instantaneous dipole - induced dipole intermolecular force attraction explained diagram illustrated δ–=> attraction <=δ+

BUT, the aldehydes have a stronger intermolecular bond due to the extra contribution of the permanent dipole - permanent dipole attractions from the highly polar >C=O bond which makes the aldehyde (or isomeric ketone) molecule highly polar overall e.g. as with propanone

••••permanent dipole - permanent dipole attractions

and for any carbonyl compound with a >C=O bond i.e. δ+C=Oδ–••••δ+C=Oδ–

The same trends would apply for the isomeric C5H10O ketones: pentan-2-one (2-pentanone), pentan-3-one (3-pentanone) and 3-methylbutan-2-one (3-methyl-2-butanone).

As a result of the increase in the strength of the aldehydes (and the same for isomeric ketones) you expect the carbonyl compounds to have (and do have) ...

... you get higher boiling points and higher enthalpies of vaporization, since a higher molecular kinetic energy is needed to overcome the stronger intermolecular bonding to escape from the liquid surface.

Another physical property affected is the dynamic viscosity, which is higher for the aldehydes (and isomeric ketones) due to their greater strength of the intermolecular bonding.

* Note on polarizabilities

Polarizabilities of pentan-2-one/pentan-3-one (2-pentanone/3-pentanone) are 9.93, so the aldehydes pentanal and 2-methylbutanal are probably similar.

Polarizability of propanal is 6.35 and that of butanal 8.18, so you might expect the polarizability of pentanal to be around 10, which fits in with the 9.93 values above.

(d) 8.2.3 Permanent dipole – permanent dipole interactions of intermolecular bonding

(d) Comparing 2-methylpropane,  methoxyethane (ethyl methyl ether), propanone (acetone) and propan-1-ol (1-propanol)

A rather more diverse set of molecules in terms of their intermolecular bonding

Name of compound Molecular formula Structural formula Mr elect-rons Dipole moment Boiling point Enthalpy of vap'n Polariz'y Dynamic viscosity
methylpropane C4H10 (CH3)3CH 58 34 0.13 D -12oC 25.2 8.0 gas
methoxyethane C3H8O CH3OCH2CH3 60 34 1.17 D 7oC 26.0 7.0 gas
propanone C3H6O (CH3)2C=O 58 32 2.88 D 56oC 32.0 6.3 0.32
propan-1-ol C3H8O CH3CH2CH2OH 60 34 2.27 D 78oC 47.5 6.7 2.3

Notes on data:

Number of electrons = sum of atomic numbers, Dipole moment units in Debye,

Enthalpy of vaporization ΔHvap in kJ mol-1,  Dynamic viscosity of liquid in mPa•s at 20oC,

Polarizability in 10-24cm3 (but just think of it as a relative number).

PLEASE NOTE the use of abbreviations ID-ID, PD-PD and HB for the intermolecular forces.

2-methylpropane: only instantaneous dipole - induced dipole forces (London dispersion forces) (ID-ID)

Alkane dipole moments tend to be zero or very small, a very non-polar molecule, so despite the higher polarizability, the other three compounds have a higher boiling point because of their polar nature and increase in strength of the intermolecular bonding.

London forces instantaneous dipole - induced dipole intermolecular force attraction explained diagram illustrated ••••London forces instantaneous dipole - induced dipole intermolecular force attraction explained diagram illustrated δ–=> attraction <=δ+

methoxyethane (ethyl methyl ether): ID-ID forces plus permanent dipole - permanent dipole (PD-PD)

There is a slight cancelling out of the δ-O-δ+C-Oδ- dipoles, so it is weekly polar molecule.

Never-the-less, the small extra PD-PD attractive force raises the boiling point compared to methylpropane.

propanone (acetone): ID-ID plus PD-PD forces, but a more polar molecule than methoxyethane.

With a more pronounced dipole δ+C=Oδ- from the polar bond the PD-PD attractive force is much stronger, resulting in stronger intermolecular bonding compared to methylpropane and methoxyethane and both the boiling point and enthalpy of vapourisation are raised compared to them too.

••••permanent dipole - permanent dipole attractions

diagram of intermolecular hydrogen bonding forces between liquid alcohol molecules doc brown A level organic chemistry revision notespropan-1-ol (1-propanol): ID-ID plus PD-PD forces plus hydrogen bonding (HB) and the latter is the strongest of the intermolecular attractive forces.

Although the dipole moment is smaller than for propanone, the total strength of the intermolecular bonding is considerably increased by the extra contribution from hydrogen bonding, so the boiling point and enthalpy of vapourisation are greater than the previous three molecules.

In the right diagram R = CH3CH2CH2

and Hδ+llll:Oδ– the directed PD-PD attractive force of the hydrogen bond

Detailed notes on hydrogen bonding

Summarising the trends

The order of intermolecular bond strength is

propan-1-ol  >  propanone  >  methoxyethane  >  2-methylpropane

The order of boiling point and the accompanying enthalpy of vaporization is ...

propan-1-ol  >  propanone  >  methoxyethane  >  2-methylpropane

Is this the expected order? YES from the arguments above, because of the increase in the strength of the intermolecular bonding from different intermolecular force contributions. You need to be able to make this sort of prediction and justify it.

The order of dipole moment is ...

propanone  >  propan-1-ol  >  methoxyethane  >  2-methylpropane

Fits the general trend of: highly polar > weakly polar > ~non-polar

The order of polarizability is ...

2-methylpropane  >  methoxyethane  >  propan-1-ol  >  propanone

The polarizabilities are all quite similar, but I've noticed that for molecules of similar 'electron size', with a very electronegative atom, tend to have lower polarizabilities - due to holding on to the electron clouds a bit more tightly?

The order of dynamic viscosity is ...

propan-1-ol  >  propanone (others are gases at 20oC)

The increase in the strength of the intermolecular bonding trend also explains why propan-1-ol is a more viscous liquid than propanone.

See for a fully detailed description of hydrogen bonding - many examples explained

Some learning objectives

Know that permanent dipole - permanent dipole attractive forces intermolecular attractive forces are also called London dispersion forces.

Know that a polar bond results because of the difference in electronegativity between the two atoms constituting a covalent bond.

Know that a polar bond can cause a molecule to be polar molecule and have a dipole moment greater than zero.

Know that polar molecules will always attract each other due to the permanent dipole - permanent dipole attractive forces attractive forces.

Know that the intermolecular bond may be due to several types of dipole attraction.

Know that the intermolecular bond between hydrocarbons is only due to instantaneous dipole - induced dipole forces

Know that for polar molecules the intermolecular bond is due instantaneous dipole - induced dipole forces plus permanent dipole - permanent dipole attractive forces (e.g. aldehydes and ketones) and may be hydrogen bonding too (e.g. alcohols).

Given a set of molecules, be able to reason the order of boiling point, based on the strength of the intermolecular bond and the contributions of the different intermolecular forces quoted above.

Be able to argue why an increased intermolecular bond will tend to also increase the enthalpy of vaporization and the dynamic viscosity of the liquid.

WHAT NEXT?

8.2 INDEX for intermolecular bonding (intermolecular attractive forces) and examples explained

Sub-index for this page on intermolecular forces an intermolecular bonding

(1) Three introductory sections important to read (a) to (c) first

(a) Introduction to intermolecular forces – Van der Waals forces

(b) Dipole moments and the polarizability of a molecule

(c) Electronegativity, covalent bond character and polar molecules

Four types of intermolecular attractive force sections (2) to (4) are the most important

(2) Instantaneous dipole – induced dipole interactions (London forces or dispersive forces)

(3) Permanent dipole – permanent dipole interactions (Keesom forces/orientation forces)

(4) Hydrogen bonding intermolecular forces

(5) Permanent dipole – induced dipole interactions (Debye forces or induction forces)

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INDEX Part 8. Phase equilibria–vapour pressure, boiling/melting points, solubility and intermolecular forces

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Part 8 sub–index: 8.1 Vapour pressure, nature, origin and examples explained * 8.2.1 Introduction to the types of intermolecular forces and examples explained (index) * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecule of similar molecular mass * 8.3 Boiling point plots of six organic homologous series - graphs and explanation * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * Evidence and theory for hydrogen bonding in simple covalent hydride * 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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