Doc Brown's Chemistry  Revision Notes

Theoretical–Physical Advanced Level Chemistry

8.2.1 Intermolecular Forces

Intermolecular Bonding – Van der Waals forces

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)

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 and its importance in other molecules * 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

8.2 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'

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

• 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 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.

  • 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.

• 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 Van Waals intermolecular forces or intermolecular bonds

  1. Instantaneous dipole – induced dipole interaction

     • Also called London forces or dispersive/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 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.

     • Note: You come across other words instead of instantaneous e.g. 'temporary', 'transient' or even 'induced' – induced dipole attractions.

     • See comparing organic molecule boiling points and homologous series comparison

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

  2. 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=O^δ–....^δ+C=O^δ– in carbonyl compounds.

     • HYDROGEN BONDING

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

     • Hydrogen bonding only usually occurs when the three most electronegative elements (N, O and F) are bonded to a hydrogen atom (more on hydrogen bonding in section 8.6).

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

     • ^δ–N–H^δ+ e.g. in ammonia, amines, amides

     • ^δ–O–H^δ+ e.g. in water, alcohols, carboxylic acids

     • ^δ+H–F^δ– in hydrogen fluoride

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

     • These are the strongest permanent dipole permanent dipole intermolecular forces

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

     •  ^δ–O–H^δ+ llll ^δ–O–H^δ+ llll ^δ–O–H^δ+ llll ^δ–O–H^δ+ llll  in water (liquid or solid ice)

     • llll ^δ–N–H^δ+ llll ^δ–N–H^δ+ llll ^δ–N–H^δ+  in amines or liquid ammonia,

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

       • ^δ+C=O^δ– llll ^δ+H–O^δ– from molecule to molecule.

     • In hydrogen fluoride, in all physical states you get a 'chain connection' llll ^δ–F–H^δ+ llll ^δ–F–H^δ+ llll ^δ–F–H^δ+

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

       • ^δ–N–H^δ+ llll ^δ–O–H^δ+ in aqueous ammonia solution (NH_3(aq),

       • C-^δ–O–H^δ+ llll ^δ–O–H^δ+ in aqueous carboxylic acid solutions (RCOOH_(aq)),

       • ^δ–H–H^δ+ llll ^δ–O–H^δ+ in hydrofluoric acid solution (HF_(aq)).

     • For detailed case studies see my main hydrogen bonding page

     • and hydrogen bonding in protein structure

     • and comparing organic molecule boiling points and homologous series comparison

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

  3. 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

• 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/centigrade)

• 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

• 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

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. 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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