INTERMOLECULAR

INTERACTIONS

INTRAMOLECULAR INTERACTION

(http://www.chem.ufl.edu/~itl/4411/lectures/lec_g.html)

INTERMOLECULAR INTERACTIONS

(http://www.chem.ufl.edu/~itl/4411/lectures/lec_g.html)

Modelling of Intermolecular Interactions

Intermolecular interaction

Short range Long range

• < 3 Å• Repulsive

• > 3 Å• Attractive• Van der Waals

interaction

Short Range Repulsive Interactions

• When two non-bonded atoms approach each each other, at some distance overlap of the occupied orbitals results in electrostatic repulsion between the electrons of those atoms.

• This repulsive energy acts over a very short range, but goes up very sharply when that range is violated.

• The repulsion goes up as 1/r12. It is important only when atoms are in very close proximity, but then it becomes very important

• Because this repulsive term rises so sharply as distance decreases it is sometimes reasonable to think of atoms as hard spheres, like small pool balls, defined by van der Waals radii and surfaces.

• When two atoms approach each other their van der Waals surfaces make contact when the distance between them reaches the sum of their van der Waals radii.

• Here we are assuming that bonds do not form. When bonds form van der Waals radii are violated.

• The smallest distance between two non-bonded atoms is the sum of the van der Waals radii of the two atoms.

• The van der Waal radius of carbon is evident from the spacing between the layers in graphite.

• The distance between atoms in different layers of graphite is never less than twice the van der Waals radius of carbon (2 x 1.7 = 3.4 Å).

• The atoms within a graphite layer are covalently linked and so are in violation of the van der Waals radius.

DIPOLE

An electric dipole is a separation of positive and negative charges

(http://ww2.chemistry.gatech.edu/~lw26/structure/molecular_interactions/mol_int.html)

(http://hcxy.wzu.edu.cn/gdwlhx/ShowNews.aspx?ID=6845&Tid=279)

• Due to non-uniform distributions of positive and negative charges on the various atoms, many molecules have dipole moments.

• Such is the case with polar compounds like water (H2O), where electron density is shared unequally between atoms.

• The physical chemist Peter J. W. Debye was the first scientist to study molecular dipoles extensively.

Dipole moment ():

= ql (1)

where q : charge l : distance of positive and negative charges

Dipole energy:

3r0

2

r0

2

44q

U

where 0 : permittivity of vacuum (= 8.854 10-12 F/m)r : relative permittivity or dielectric constant of the medium where the charges are located in.

(2)

For molecules there are three types of dipoles:

1. Permanent dipoles: These occur when two atoms in a molecule have substantially different electronegativity: One atom attracts electrons more than another, becoming more negative, while the other atom becomes more positive. A molecule with a permanent dipole moment is called a polar molecule.

2. Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole.

(http://www.chemprofessor.com/imf.htm)

3. Induced dipoles: These can occur when one molecule with a permanent dipole repels another molecule's electrons, inducing a dipole moment in that molecule. A molecule is polarized when it carries an induced dipole. See induced-dipole attraction.

Dipole moment values of some typical gas phase in debye units are:

Compound Dipole Moment (Debyes)

NaCl 9.0 (measured in the gas phase)

CH3Cl 1.87

H2O 1.85

NH3 1.47

CO2 0

CCl4 0

Ion-Dipole Interaction

(http://www.chem.purdue.edu/gchelp/liquids/iondip.html)

A dipole that is close to a positive or negative ion will orient itself so that the end whose partial charge is opposite to the ion charge will point toward the ion.

This kind of interaction is very important in aqueous solutions of ionic substances.

H2O is a highly polar molecule, so that in a solution of sodium chloride, the Na+ ions will be enveloped by a shell of water molecules with their oxygen-ends pointing toward these ions, while H2O molecules surrounding the Cl– ions will have their hydrogen ends directed inward.

The interaction energy can be constructed from the Coulomb interactions between the bare charge Q and the dipolar charges q:

21r0 r1

r1

4Qq

rU (3)

22

1 sin2

cos2

rr

22

2 sin2

cos2

rr

r2

r1

whereQ : charge of ionr : distance between ion and dipole0 : permittivity of vacuum (= 8.854 10-12 F/m)r : relative permittivity or dielectric constant of the medium where the charges are located inl : length of dipole moment vector : the angle between the dipole moment vector and the vector connecting the ion with the dipoleWhen the dipole is sufficiently far away from the charge

(r >> l), we can approximate:

cos2

rr1

cos2

rr2

The interaction energy is then

cos2

r

1

cos2

r

14

QqrU

r0

2r0 r4

cosQrU

where q

(4)

2

22r0 cos

4r

cos4

Qq

(1)

• When deriving eq. (4) we assumed the dipole to be sufficiently far from the charge (r >> l).

• At about r < 2l the approximation deviate more than 10% from the exact result.

• Taking into account the finite size of atoms and molecules, eq. (4) is actually an excellent approxi-mation for interactions between ions and small polar molecules at all physically relevant distances.

• In the case of larger molecules, where the charges comprising the dipole may be several ångströms apart, we need to exercise good judgement whether to use eq. (4) or not.

Dipole-Dipole Interaction

(http://2012books.lardbucket.org/books/general-chemistry-principles-patterns-and-applications-v1.0m/section_15_02.html)

(http://2012books.lardbucket.org/books/general-chemistry-principles-patterns-and-applications-v1.0m/section_15_02.html)

• As two dipoles approach each other, they will tend to orient themselves so that their oppositely-charged ends are adjacent.

• Two such arrangements are possible: the dipoles can be side by side but pointing in opposite directions, or they can be end to end.

• It can be shown that the end-to-end arrangement gives a lower potential energy.

• Dipole-dipole attraction is weaker than ion-dipole attraction, but it can still have significant effects if the dipole moments are large. The most important example of dipole-dipole attraction is hydrogen bonding.

cossinsincoscos2r4

U 21213r0

21 (5)

1 and 2 are the polar orientation angles of 1 and 2, respectively, and is the azimuthal orientation angle of 2 in reference to 1

for r > 3 l:

• Dipole-dipole interaction is comparatively weak (for dipole moment of 1 Debye at 0.35 nm in vacuum, the interaction is already weaker than kT).

• In certain molecules (small size and large dipole moment O-H, N-H, and F-H), dipole-dipole interaction can lead to short range association in liquid (part of H-bond).

• Dipole-dipole interaction is strongest when the two dipoles mutually orient themselves in line.

• At large separation or in a medium of high , when interaction falls below kT, dipoles can now rotate more freely.

• The angle averaged potentials are not zero because of Boltzmann weighting factor, the energy (Keesom interaction or orientation interaction) becomes

6

2

r0

21

B r1

4Tk1

32

rU

(6)

for

3r0

21B r4Tk

Ion-Induced Dipole Interactions

Polarizability ()

• Polarizability is the ease of distortion of the electron cloud of a molecular entity by an electric field (such as that due to the proximity of a charged reagent).

• It is experimentally measured as the ratio of induced dipole moment (ind) to the field E which induces it:

Eind

0

(7)

• Nearby ions can distort the electron clouds even in polar molecules, thus temporarily changing their dipole moments.

• The larger ions (especially negative ones such as SO22–

and ClO42–) are highly polarizable, and the dipole

moments induced in them by a cation can play a dominant role in compound formation.

• The most significant induced dipole effects result from nearby ions, particularly cations (positive ions).

Dipole-Induced Dipole Interactions(induction or Debye interaction)

(http://textbook.s-anand.net/ncert/class-11/chemistry/5-states-of-matter)

• A permanent dipole can induce a temporary one in a species that is normally non-polar, and thus produce a net attractive force between the two particles.

• This attraction is usually rather weak, but in a few cases it can lead to the formation of loosely-bound compounds.

• This effect explains the otherwise surprising observation that a wide variety of neutral molecules such as hydrocarbons, and even some of the noble gas elements, form stable hydrate compounds with water

(8) 62r0

221

Debye r1

42U

Dispersion or London Forces

(http://www.chem.ufl.edu/~itl/4411/lectures/lec_g.html)

Noble gas elements and completely non-polar molecules such as H2 and N2 can be condensed to liquids or solids.

There must be another source of attraction between particles that does not depend on the existence of

permanent dipole moments in either particle.

• A molecule is “nonpolar” the time-averaged dipole moment is zero.

• On a very short time scale, however, the electron must be increasingly localized.

• As a consequence, there is no guarantee that the distribution of negative charge around the center of an atom will be perfectly symmetrical at every instant; every atom therefore has a weak, fluctuating dipole moment that is continually disappearing and reappearing in another direction.

• Although these extremely short-lived fluctuations quickly average out to zero, they can still induce new dipoles in a neighboring atom or molecule, which helps sustain the original dipole and gives rise to a weak attractive force known as the dispersion or London force

• Dispersion is applicable to all atoms or molecules (unlike Keesom or Debye interaction).

• It is responsible for certain phenomena in macro-scopic scale (adhesion, surface tension, physical adsorption, wetting, properties of gases and liquid, structures of condensed macromolecules,... ).

• It is a long range force that can be effective at large distance (>10 nm) to interatomic spacings.

• Dispersion is non-additive. The dispersion of two molecules is affected by the presence of the third molecules

Fritz London (1937) proposed a theory based on quantum mechanics to explain dispersion

21

2162

0

21London II

IIr1

423

U

(9)

where I is the first ionization potential I = h

van der Waals Interaction

• Many kinds of molecules possess permanent dipole moments, so liquids and solids composed of these species will be held together by a combination of dipole-dipole, dipole-induced dipole, and dispersion forces.

• These weaker forces (that is, those other than coulombic attractions) are known collectively as van der Waals forces.

• These are short-ranged and weak interactions existing between all types of atoms and molecules.

• All atoms and molecules, even non-polar and uncharged ones, exert attractive forces on each other.

• This is a result of the atomic polarizability 0 of atoms.

• The constant motion of electrons in atoms results in the fact that at any given instant in time, any atom actually has a finite electric dipole moment.

• Despite the quantum-mechanical uncertainty of position, even particles as light as electrons have to occupy some region of space at a given time.

• In the absence of an external influence, the average dipole moment of an atom is zero.

In general, the van der Waals forces arise from three different contributions:

(1) orientation or Keesom interaction; (2) induction or Debye interaction; and (3) dispersion or London interaction

Thus, in general the total van der Waals energy is given by

LondonDebyeKeesomvdW UUUU (10)

• The dispersion term is the most important of the three contributions, as it is always present, regardless whether permanent dipoles take part in the interaction or not.

• Moreover, usually the dispersion term is also the strongest, contributing around 80 - 100% to the total van der Waals interaction energy.

• A notable exception to this is water, where in fact the Keesom interaction dominates (about 70% of the total interaction energy)

Summary

Intermolecular forces are responsible most properties of all the phases:

1. Gas: Vapor pressure, critical point, and boiling point

2. Liquid: viscosity, diffusion, and surface tension.

3. Solid: melting and sublimation.

• The calculation of the potential energy involves assumptions concerning the nature of attraction and repulsion between molecules.

• Intermolecular interaction is the result of both short- and long-range effects.

• Electrostatic, induction, and dispersion effects are examples of long range interactions.

• In these cases, the energy of interaction is proportional to some inverse power of intermolecular separation. Electrostatic interactions result from the static charge distribution between molecules.

• The effect can be either attractive or repulsive and it is exclusively pairwise additive.

• Induction effects are always attractive, resulting from the distortions caused by the molecular fields of neighbouring molecules

• The specification of intermolecular potential, representing interaction between molecules, is a critical step in Molecular Dynamics simulations.

• Generally, a two-body potential of the form U(rij) is used, where rij is the distance between the centers of molecules i and j.

• This form neglects multibody interactions.

• Once the potential is prescribed, the intermolecular force is obtained from Fij = -dU(rij)/drij.

• In the following, we will discuss intermolecular potential models.

Hard-Sphere Potential

In this model, the molecules move freely and do not interact with one another except when they collide. The intermolecular potential is given by

rrU

r0rU(11)

Here (σ) is the diameter of the molecule.Thus the molecules exert force on one another only when they collide

Intermolecular potential for the hard-sphere modelwww.uic.edu/eng/ems/MEng/ChEME494/pdf/L8pt1.pdf

Square-Well Potential

The square-well potential is the simplest inter-molecular potential that is capable of representing the properties of liquids

r0r

rrU (12)

where is some multiple of the hard-sphere diameter and is a measure of the attractive interaction.

The square-well potential represents a mathematically idealised model of molecular interactions.

U(r)

Square well potential model

Yukawa Potential

The square-well potential can be made more realistic by changing the variation of attractive interactions. There have been many such variations of which the Yukawa potential is an important example.

r1r

expr

rrU (13)

where is an attractive term (depth of potential well), is the hard-sphere diameter and is an adjustable parameter.

The inverse power dependence of this potential means that it can be applied to ionic systems.

Hard-core Yukawa potential with various interaction ranges(Naresh and Singh, 2000)

Lennard-Jones Potential:

The Lennard-Jones potential (also referred to as the L-J potential, 6-12 potential, or 12-6 potential) is a mathematically simple model that approximates the interaction between a pair of neutral atoms or molecules.

612

rr4rU (14)

where is the depth of potential well, is the hard-sphere diameter (the finite distance at which the inter-particle potential is zero), and r is the distance between the particles.

http://what-when-how.com/molecular-biology/van-der-waals-interactions-molecular-biology/

Lennard-Jones potential

www.uic.edu/eng/ems/MEng/ChEME494/pdf/L8pt1.pdf

• The parameter σ is the zero energy separation distance, and defines a molecular length scale related to the particle diameter, while ε is the minimum energy and controls the strength of the interaction.

• As an example, σ= 0.41 nm, and ε/kB = 221K for xenon, where kB is the Boltzmann's constant.

• For a pure substance, σ equals the particle diameter.

• The parameter σ is also related to the critical volume of the fluid, while ε to its critical temperature.

• This dependence is particularly strong for small molecules.

For simulations involving two fluids, σ and ε are expressed using the Lorentz-Berthelot mixing rules (Allen and Tildesley, 1984), as:

jiij 21

jiij

(15)

(16)

Intermolecular ForceThe force between the two L-J molecules is given by

(17)

By convention, repulsive (short-range) forces are positive while attractive (long-range) forces are negative. i.e.,

713

rr224

drdU

F

612rfor0F:pulsionRe

612rfor0F:Attraction (18)

Lennard-Jones intermolecular forcewww.uic.edu/eng/ems/MEng/ChEME494/pdf/L8pt1.pdf