Doc Brown's Advanced A level theoretical chemistry revision notes

8.2.1 Introduction to intermolecular bonding (intermolecular attractive forces) including Van der Waals forces, polarizability of molecules, dipole moments of molecules and electronegativity

8.2.1 Introduction to intermolecular bonding (intermolecular attractive forces)

Sub-index for this page

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

(b) (1) Dipole moments and polar bonds

(b) (2) the polarizability of a molecule

(c) Electronegativity: polar bonds, covalent bond and polar molecules (read 1st if not already studied)

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

also referred to as 'intermolecular bonding' forces

The different types of intermolecular force (intermolecular bond) are described, explained and discussed with examples, collectively known as Van der Waals forces i.e. in terms of the various origins of intermolecular bonding we have the attractive forces of:

Instantaneous dipole – induced dipole interaction (London forces, dispersion forces)

Permanent dipole – permanent dipole interactions (Keesom forces, orientation forces)

Hydrogen bonding is often excluded from the definition of Van der Waals forces, but it is the strongest of the permanent dipole - permanent dipole attraction and can be detected in gases.

Permanent dipole – induced dipole interactions (Debye forces, induction forces)

These are all are electrostatic attractions - attractive Coulombic forces where the +ve of a dipole of a polarized molecule attracts the -ve of the dipole of a neighbouring polarized molecule (or vice versa). The dipoles may be permanent, instantaneous or induced.

Reminder: Opposite poles on neighbouring molecules will attract and like poles on neighbouring molecules will repel.

(Do NOT confuse the weak intermolecular bonding Van der Waals attractive forces or hydrogen bonding with the strong chemical bonding between atoms or ions i.e. so-called ionic bond, intramolecular covalent bond or metallic bond)

e.g. it takes 2 x 464 = 928 kJ mol^-1 to break the two O-H intramolecular bonds in the water molecule, but only 40.7 kJ mol^-1 to vaporise liquid water in overcoming the intermolecular bonding.

The above diagram illustrates this point for the group 7/17 halogen molecules

When molecules approach each other, there is both repulsion (of electron clouds) and attractive forces (dipoles) operating between the molecules, so, technically, intermolecular interactions are the sum of both attractive and repulsive components.

Note that each end of a dipole (transient or permanent) is only a fraction of the charge of an electron.

A definition of Van der Waals forces

These can be defined as weak, short-range electrostatic attractive forces between uncharged molecules (i.e. not ions), arising from the interaction of permanent or transient electric dipole moments and their different types and origin are described in detail with lots of explained examples on four separate pages (links further down).

• From the start understand that:

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

  • Throughout these notes the dipoles are depicted as lower case delta plus (δ+ ) and lower case delta minus (δ–)

• 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. the covalent bonds 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 (including Van der Waals force) is quoted as a summation of the various possible dipoles interaction and the principal attractive forces are shown below with links to pages with lots examples and explanations.

  • Irrespective of polar bonds and the polarity of the molecule (factor (iii)), the strength of the Van der Waals intermolecular attractive forces, always depend to some extent on these two factors, which if increased, increase the intermolecular force:

    • (i) The size of the atom or molecule - best thought of as the number of electrons in the electron clouds (atomic or molecular orbitals). This is also linked to the ease of polarizability of the molecule (discussed in the next section).

    • (ii) The surface area contact between the molecules (on a time averaged basis). The larger the surface area, the greater the chance of dipole - dipole attraction.

    • Factor (ii) can override factor (i) and you can see this in the physical properties of the isomers of alkanes.

    • Factor (iii) The permanent polarity of a molecule as result of it having polar bonds due to differences in electronegativity of the two atoms.

• The four types of intermolecular bonding forces are described in details on the following separate pages:

• instantaneous (temporary) dipole - induced dipole intermolecular forces

• permanent dipole - permanent dipole intermolecular forces

• directed permanent dipole - permanent dipole (hydrogen bonding) intermolecular forces

• permanent dipole - induced dipole intermolecular forces

• 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.

• What sort of properties does intermolecular bonding affect?

• Any increase in the intermolecular bonding strength tends to increases bulk properties such as:

  • Melting point - molecules need more kinetic energy to vibrate sufficiently 'violently' to overcome the intermolecular bonding and enthalpy of fusion is increased.

  • Boiling point - molecules need more kinetic energy to overcome the cohesive effect of intermolecular forces and escape from the surface of the liquid and enthalpy of vaporisation is increased.

  • Dynamic viscosity of a liquid increases if the intermolecular bonding forces increase - the molecules stick together more strongly.

  • Physical strength of the solid - tensile strength and flexibility.

  • Where appropriate, all these effects are described on the four page links above.

  • Deviation from ideal gas behaviour are not discussed in this section

    • See Non–ideal real gas behaviour and Van der Waals Equation of State

(b) Dipole moments and the polarizability of a molecule

(b1) DIPOLE MOMENTS

The electric dipole moment is a measure of the separation of partial positive and negative electrical charges within a molecule and is a measure of the molecule's overall polarity.

The dipole moment is a vector quantity whose magnitude is equal to the charge multiplied by the distance between the charges and the direction is from negative charge to positive charge. It is measured in Debye units (D)

dipole moment µ = q x d, for dipole charges q+ and q- and distance d apart

1 Debye  1 D = 3.336 x 10^-30 C m

Molecules may have a permanent dipole moment due to unevenness or lack of symmetry in the distribution of charge in the molecule due to a polar bond e.g. H^δ+-Cl^δ+ i.e. one end of the molecule has a partial positive charge compared to the other end of the molecule with a partial negative charge due to differences in electronegativity. See section below on electronegativity.

Data table of dipole moments (in D, Debye) and comments

Examples of simplified molecule diagrams with their dipoles and shape:

bent ^δ+H₂O^δ-,  trigonal pyramid ^δ-NH₃^δ+,  bent ^δ+SO₂^δ-  and tetrahedral  ^δ+CHCl₃^δ-

The symmetrical molecules of carbon dioxide (linear ^δ-O=C^δ+=O^δ-) and tetrachloromethane (perfect tetrahedral molecule with 4 ^δ+C-Cl^δ- bonds) have zero dipole moments because the dipole charges cancel each other out.

Despite the increase in bond length, the dipole moment increase for the group 7/17 hydrogen halides ^δ+H-X^δ- is from HI to HF, which is in line with the increase in electronegativity of the halogen and the increasing polar nature of the hydrogen halide due to an increasing difference in electronegativity.

^δ+H-O^δ-H^δ+ Water is very polar molecule with a relatively high dipole moment.

Carbon monoxide is weakly polar molecule with a small dipole moment, but despite having two polar bonds, carbon dioxide has a zero dipole moment.

^δ+SO₂^δ- Sulfur dioxide is a highly polar molecule with two very electronegative oxygen atoms acting on the sulfur atom at the centre of the bent molecule.

Unlike tetrachloromethane, the less symmetrical trichloromethane ^δ+CHCl₃^δ- is an overall polar molecule with a dipole moment.

Alkanes have no polar bonds and no double bond of extra electron density, unlike unsymmetrical alkenes, so their dipole moments are zero or very small.

R-O^δ-H^δ+ R = alky. The lower alcohols are very polar molecules with relatively high dipole moments.

(b2) POLARIZABILITY

An extremely important concept for understanding the origin and effects of intermolecular bonding forces and will be discussed on the four pages on intermolecular bonding forces describing the trends and effects on selected physical properties of molecules.

Even without a permanent dipole moment, at any instant in time it molecules will almost certainly have a transient dipole moment because of the 'random' behaviour of the electron clouds. Although on average these transient dipole moments will average out to zero, at a given instant in time they can give rise to intermolecular attractive forces between nearby molecules.

Any molecule with a permanent or transient dipole produces an electric field and so can induce a dipole moment in neighbouring molecules and always causes the molecules to attract each other.

The magnitude of the induced dipole moment will depend on how easy it is to distort the electron cloud in the molecule and is called the polarizability of the molecule.

Polarizability is defined as the tendency of an atom or molecule's electron cloud to be distorted by an external electric field, enabling the formation of a temporary dipole moment.

Its all about the ease of deformation of the electron clouds into a more asymmetric distribution of electron charge.

The external electric field can be due to a neighbouring polar molecule with a permanent dipole or a molecule exerting a transient instantaneous dipole.

Polarizability depends on the size of the atoms making up a molecule are and the number of atoms in the molecule - collectively you can think of the volume of the electron clouds and the number of electrons in the molecule - the greater the volume of the electron clouds, the more polarizable is the molecule and therefore an increase in the intermolecular forces.

The larger an atom the outer electrons are held less tightly and the more readily the electron clouds can be distorted by a neighbouring electric field of another molecule i.e. even in non-polar molecules, you get attraction between a transient instantaneous dipole moment in one molecule and the induced dipole in another molecule (dispersion force, see below).

Data table of polarizability values and comments for non-polar molecules

The polarizability of a molecule is a measure of easily the electron clouds of a molecule are distorted by an external electric field to create a dipole. The units are 10^-24 cm³.

The units of polarizabilty are not important at pre-university level, just treat them as relative values and most of the polarizability data was obtained from the internet and values do vary a little depending on the source, but the NIST source is perhaps the best at https://cccbdb.nist.gov/pollistx.asp

Comments on the above polarizability data

You need to consider number of electrons, polar bonds and the idea of relative dipole contact due to differences in surface area, all of which can significantly affect physical properties such a boiling point - for more details see the four pages on the types of intermolecular force.

helium has the lowest polarizabilty of any other atom or molecule.

Comparing the polarizability of linear and branched alkanes (the latter polarizability not tabulated)

C₄H₁₀: butane (8.02) isomeric with 2-methylpropane (isobutane) (8.01)

C₅H₁₂: linear pentane (9.88) and branched isomers 2-methylbutane (8.77) and  2,2-dimethypropane (10.24).

C₈H₁₈: linear octane (15.24) and most branched isomer tetramethylbutane (15.61).

These isomer examples show that the polarizability seems to be more governed by the number of electrons in the molecule, rather than the shape i.e. linear and highly branched are similar.

With the linear alkanes, as the number of electrons increases, with increase in carbon chain length, there is a 'larger overall electron cloud' that becomes easier to distort i.e. become polarized, perhaps in several place simultaneously in a longer carbon chain.

In other sequences of homologous series in organic chemistry e.g. linear alcohols, you see exactly the same pattern as with alkanes with increasing polarizability with increase in size of molecule with increase in number of electrons in the molecule.

In the series of Group 7/17 hydrogen halides there is a steady increase in polarizability as an extra electron shell is added as you descend the group from HF to HI, but this trend is the opposite in terms of the polarity of the bond.

Even the most electronically stable Group 0/18 noble gas atoms become more polarizable with increase in atomic number because of the increasing number of full electron shells as you descend the group.

CH₄ and Ne (10 electrons) have very different polarizabilities of 2.59 and 0.38 respectively because the electrons in the Group 0/18 noble gas are the most strongly held of the elements

(c) ELECTRONEGATIVITY - another extremely important concept for understanding the origin and effects of intermolecular bonding forces

Electronegativity - its relationship to covalent bond character, polar bonds and polar molecules

• Electronegativity is a measure of the power of an atom to attract electron charge from another atom it is covalently bonded to.

• Some Pauling values of electronegativity are quoted 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

• Generally speaking electronegativity increases from left to right across a period of the periodic table and decreases down a group of the periodic table.

• The greater the difference in electronegativity between two atoms of a covalent bond, the more polar the bond and often results in a molecule being polar overall.

• The presence of polar bonds tends to increase the net total intermolecular force acting between neighbouring molecules.

• For a covalent bond, the most electronegative element atom carries the δ- charge and the least electronegative element atom the δ+ charge.

  • e.g. ^δ-Cl-C^δ+, ^δ-O-C^δ+ and ^δ+C-N^δ- polar bonds.

• Note that the most electronegative elements are nitrogen, oxygen and fluorine and this is a significant factor in the formation of intermolecular hydrogen bonds via the ^δ-N-H^δ+, ^δ-O-H^δ+ and ^δ+H-F^δ- polar bonds.

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 on this page

(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

The total intermolecular bonding can arise from between one and four of these forces.

(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)

ALL my advanced A level theoretical chemistry revision study notes

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. 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