Doc Brown's Advanced A level theoretical chemistry revision notes

8.2.2 Instantaneous dipole – induced dipole intermolecular bonding, examples of London/dispersion forces explained

8.2.2 Instantaneous dipole – induced dipole interactions of intermolecular bonding

1. Introduction - names used and origin of the attractive force

• London forces or dispersion forces are caused by an instantaneous dipole inducing a dipole in a neighbouring molecule, and so causing electrostatic attraction.
• Hence the descriptive term instantaneous dipole - induced dipole intermolecular bonding force, which is the weakest component of the Van der Waals forces.
• This attractive force is named after the German physicist Fritz London (1900-1954).
• Other names used for these forces arising from temporary fluctuations in electron distribution.
• or transient dipole - induced dipole attractive forces
• or temporary dipole - induced dipole forces (sorry, but all 4 descriptors are used)
• In the diagrams •••• represents the intermolecular bonding.

• (Read the 2nd part of the introduction about the polarizability of a molecule)

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

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

  • Note that these partial charges are shown as a delta plus (δ⁺) or a delta minus (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 in crystals of the ionic compound sodium chloride.
  • The London dispersion force or instantaneous dipole - induced dipole intermolecular bonding force acts in all directions between neighbouring molecules.
  • These forces act between all particles whatever their nature whether they be non-polar or polar molecules or even ions (which can be considered as molecular ions).

• The random movement of electrons in their atomic/molecular orbitals creates these temporary (induced or transient or instantaneous) dipoles, which intern induces an equally transient dipole in a neighbouring molecule.

  • Sometimes described as an oscillation producing an asymmetric distribution of the electron charge i.e. a dipole created.

• Therefore e.g. this somewhat random partial positive charge of one dipole will attract the induced partial negative in the neighbouring molecule or vice versa.

• Imagine, initially, an even distribution of the electron clouds, then imagine at random, the electron clouds become spontaneously unevenly distributed to give a temporary dipole moment.

• 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.

• Final introduction notes on instantaneous dipole - induced dipole forces:

  • (i) You come across other words other than instantaneous e.g. 'temporary', 'transient', 'London', dispersion/dispersive 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 or even an ion! It is a universal intermolecular force.

• See comparing 8 organic molecule boiling points

• and homologous series comparison of boiling point trends

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

2. The polarization of a hydrocarbon molecule and the noble gas/halogen polarization trends and its consequences

• HYDROCARBONS

• From symmetrical distribution = polarization => unsymmetrical distribution of electron clouds

• = polarized => δ⁺ dipole formation δ^–

• This asymmetry in one molecule induces a dipole in a neighbouring molecule giving the instantaneous dipole - induced dipole interaction (the London forces).

• Note that the polarised particles are still neutral overall (unless they are ions in the first place).

• So, this results in a weak intermolecular force attraction between adjacent molecules, irrespective of whether the molecules are the same or different or polar or non-polar..

• •••• δ^–=> attraction <=δ⁺

• 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.

• See also section 3. Trends in physical properties of alkanes related to polarizability and intermolecular bonding

• -

• NOBLE GASES
• Even in the case of helium, which has the lowest melting point and boiling point of any substance, you still can get transient dipoles because of the random behaviour of the two electrons in the 1s orbital.
• 
• The result is transient helium atom attractions He••••He

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

  • This shows in the difference in the boiling points of CH₄ -162^oC and Ne -246^oC respectively as the intermolecular forces (ID-ID) are much weaker in neon.
  • However, you get a clear trend in properties if you just look at the Group 0/18 noble gases.

• The atomic radii increase due to the increase in filled inner electron shells as you descend the group.
• The increase in electron number increases the polarizability of the Noble Gases.
• Group 0/18 trend boiling points ^oC (and number of electrons = atomic number) are ...
• He -269 (2);  Ne -246 (10);  Ar -186 (18);  Kr -153 (36); Xe -108 (54);  Rn -62 (86)
• Therefore down group 0/18, as the number of electron shells increases, the outer electrons are further from the nucleus, therefore the increasing number of electrons are less strongly held by the nucleus and so polarization can take place more easily.
• This increase in polarization increases the probability and strength of the  instantaneous dipole - induced dipole attractive forces (London dispersion forces).
• KE note: So, for an increase in the strength of any type of intermolecular bonding, on average, a greater particle vibrational or translational kinetic energy is required to vaporise the atom/molecule and the enthalpy of fusion (melting, data not shown) or vapourisation is increased, reflecting the greater KE needed to escape from the intermolecular forces at the surface of the liquid in the case of the boiling point (or increased KE of vibration to melt the solid).
• The KE note above is a general argument for this page and any other about intermolecular bonding.
• -

• THE HALOGENS

•  

• REMINDER: The diagram also indicates the difference between strong intramolecular covalent bonds (X-X) and the much weaker intermolecular bond (X₂••••X₂), which in this case is entirely due to instantaneous dipole - induced dipole attractive forces (London dispersion forces).

• Data table for down the halogen group 7/17 ===> X2 trends
  	fluorine	chlorine	bromine	iodine	astatine
  Period	2	3	4	5	6
  Appearance of X2 molecule
  	F2  gas	Cl2 gas	Br2  liquid	I2  solid	At2 solid
  Electrons in X2	18	34	70	106	170
  Polarizability	1.16	4.61	6.43	?	?
  Melting point/oC	–219	–101	–7	114 (sub)	302 (sub)
  ΔHfus/melt kJ/mol	0.5	6.4	10.6	15.5	?
  Boiling point/oC	–188	–34	59	184	380
  ΔHvap kJ/mol	6.6	20.4	30.0	41.6	?

• Following on from the noble gas discussion, you see exactly the same pattern in the increase in melting/boiling points and state change enthalpies of the Group 7/17 halogen elements.

• As you descend the group, the diatomic molecules get bigger with the increase in inner filled electron shells per atom, so the halogen molecules become more and more polarizable and increasing the strength of the intermolecular bonding, increasing the melting/boiling points, and the corresponding enthalpies of fusion/vapourisation as the particles need a greater KE to effect a state change.

• In all three cases, all the points made, add up to the same line of argument to explain the trend in physical properties with increase in electron number and polarizability.

3. The trend in physical properties of alkanes related to their polarizability and intermolecular bonding - boiling point, enthalpy of vaporisation and dynamic viscosity.

• As described above, the force of instantaneous dipole – induced dipole interaction increases with more electrons in the molecule (London dispersion forces intermolecular bonding).

  • 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 the linear members of 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) due to more electrons and greater polarizability (actually 8 electrons per -CH₂- unit in the carbon chain).

  • This trend is described and explained for linear alkanes from CH₄ to C₁₄H₃₀ (graph of boiling points in K versus length of carbon atom chain is shown below)

  • 

• The above diagram 'attempts' is my idea to there might several points where dipoles are formed in longer hydrocarbon molecules.

• 

• This change, and increase, in intermolecular forces, gives rise to different, but increasing hydrocarbon boiling points, the basis of separating the components of crude oil by fractional distillation.

• As the organic molecule (e.g. alkanes) gets bigger with more electrons in it, the diagram  is more polarizable and the transient dipole - induced dipole effect increases - the intermolecular forces of attraction

• The boiling point rises as more kinetic energy is needed to effect the state change as the molecules (and electron number) rises with increase in the intermolecular force between neighbouring molecules.

• As the boiling point increase, so will the enthalpy of vaporization e.g. for ...

• C₅H₁₂ pentane ΔH_vap = 26.4 kJmol^-1, rising incrementally to 71.7 kJmol^-1 for C₁₄H₃₀ tetradecane.

• See the table of polarizabilities in section 8. on the alkane trend in viscosity because the steady increase in the strength of the intermolecular bond, also steadily increases the dynamic viscosity of the liquid alkane hydrocarbons.

  • For viscosity, think about the difference in 'stickiness' between petrol and lubricating oil.

• Apart from fractional distillation, the chromatographic separation of hydrocarbons can be related to the differences in intermolecular bonding strength.

• Gas-liquid chromatography is a way of analysing mixtures such as a mixture of liquid hydrocarbons.

• 

• The components are separated due to their difference in temporarily adhering to the coating in the separating column.

• The main factor controlling this is the temporary intermolecular bonds between the mixture components and the immobile phase in the separating column.

• The larger the hydrocarbon molecule, the greater the temporary intermolecular bonding, hence the smaller and least easily polarized molecules are eluted by the carrier gas from the column first with smaller retention times.

• e.g. in terms of retention time: heptane > hexane > pentane

• More details on gas chromatography

4. Trends in sequences of Group 4/14 carbon group hydrides and polyhalogenated methane molecules

• The group 4/14 carbon group hydrides

  • Relevant data table for the Group 4/14 carbon group hydrides
    Period	XH4	polarizability	electrons	bpt/K	ΔHvap kJ/mol
    2	CH4	2.45	10	112	8.2
    3	SiH4	4.78	18	161	13.0
    4	GeH4	4.77	36	184	23.0
    5	SnH4	?	54	221	32..6
    6	PbH4	?	86	260	?

  • Note: Electrons = sum of atomic numbers in the molecular formula.

  • These are all symmetrical non-polar molecules with zero dipole moments, so the intermolecular bonding is due to the instantaneous dipole - induced dipole forces (London dispersion forces).

  • Know that as the size of the hydride increases, the molecule has more electrons and a greater polarizability, increasing the strength of the intermolecular bonding via the instantaneous dipole - induced dipole forces and increasing the boiling points and enthalpy of vaporisation.

• The same electron number and polarizability arguments apply to the series of increasing melting and boiling points for methane and its multi-halogen substitute compounds.

  • The polarizabilities are in the () and the order of increasing melting/boiling points are ....

  • CH₄ (2.45) < CF₄ (2.82) < CCl₄ < (10.00) CBr₄ (?) < CI₄ (?)

  • From left to right, the size of the molecule with an increasing number of electrons steadily increases the polarizability of the molecule and leads to increase in the dispersion attractive forces (transient dipole - induced dipole forces), hence the increase in melting/boiling point as more kinetic energy is needed to effect the state change.

  • All five molecules are symmetrical with zero dipole moments, so no complications from polar bonds or polar molecules.

  • Obviously, down group 7/17 the halogens have an increasing number of filled inner electron shells, hence a significant number of electrons and a larger volume of electron clouds that are more readily polarized increasing the intermolecular bonding.

• Same argument applies to some extent to the melting/boiling points for the series increasing chlorination or bromination of methane, as each successive substitution greatly increases the number of electrons in the molecule.

  • The polarizabilities are in the () and the order of melting/boiling points are ....

  • CH₄ (2.45) < CH₃Cl (4.42) < CH₂Cl₂ (6.66) < CHCl₃ (8.13) < CCl₄ (10.00)

  • The middle three are polar molecules with a dipole moments of 1.90,1.62 and 1.02 D respectively, but this seems to make little difference to the trend, which is similar for the following too.

  • CH₄ (2.45) < CH₃Br (5.61) < CH₂Br₂ (8.68) < CHBr₃ (11.87) < CBr₄ (?)

  • Again, the middle three are polar molecules with a dipole moments of >0 D.

  • In all cases, the enthalpy of fusion/vaporization will increase in the same order, as increasingly greater kinetic energies are needed to overcome the increasing intermolecular bonding to effect a state change.

  • However for the 'middle' three molecules there will be a small contribution to the total intermolecular bonding from the permanent dipole - permanent dipole attractive forces.

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)

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