Intermolecular Forces

Intermolecular forces (IMFs) (or secondary forces) are forces that mediate interactions between molecules, including the forces of attraction or repulsion that act between atoms and other types of neighboring particles, such as atoms or ions . Intermolecular forces are weak relative to intramolecular forces – the forces that hold a molecule together. For example, covalent bonding , which involves sharing electron pairs between atoms, is much stronger than the forces that exist between neighboring molecules. Both sets of forces are essential parts of force fields frequently used in molecular mechanics .

The investigation of intermolecular forces begins with macroscopic observations that indicate the existence and action of forces at the molecular level. These observations include non-ideal-gas thermodynamic behavior reflected by the virulence coefficient , vapor pressure , viscosity , surface tension and absorption data.

The first reference to the nature of subtle forces is found in Alexis Clairot ‘s work Théorie de la figure de la terre , published in Paris in 1743 . [1] Other scientists who have contributed to the investigation of subtle forces include: Laplace , Gauss , Maxwell and Boltzmann .

Attractive intermolecular forces are classified into the following types:

  • hydrogen solder
  • The ionic bond
  • ion induced dipole force
  • ion-dipole force
  • Van der Waals force – Keysome force , Debye force and London dispersion force

Information on intermolecular forces is obtained by macroscopic measurements of properties such as viscosity, pressure, volume, temperature (PVT) data. The linkage of microscopic aspects is given by the virulence coefficient and the Lenard-Jones potential.

Hydrogen solder

hydrogen bond is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is attached to an electronegative atom, usually nitrogen , oxygen , or fluorine . [2] The hydrogen bond is often described as a strong electrostatic dipole–dipole interaction. However, it also has some characteristics of a covalent bond: it is directional, stronger than the interaction of the van der Waals force , their van der Waals radii .produces inter-atomic distances less than the sum of K, and usually involves a limited number of interaction partners, which can be interpreted as a type of valency . The number of hydrogen bonds formed between molecules is equal to the number of active pairs. The molecule that donates its hydrogen is called the donor molecule, while the molecule having lone pairs participating in the H bond is called the acceptor molecule. The number of active pairs is equal to the common number between the number of hydrogens near the donor and the number of lone pairs near the acceptor.

Intermolecular Forces
Intermolecular Forces

Although both are not depicted in the diagram, water molecules have two active pairs, because the oxygen atom can interact with two hydrogens to form two hydrogen bonds. Intermolecular hydrogen bonding is responsible for the higher boiling point (100 °C) of water compared to other group 16 hydrides , which have a lower hydrogen bond ability. Intramolecular hydrogen bonding is partly responsible for the secondary , tertiary and quaternary structures of proteins and nucleic acids . It also plays an important role in the structure of both synthetic and natural polymers .

The ionic bond

The attraction between the cationic and anionic sites is a non-covalent, or intermolecular interaction commonly known as an ion pairing or salt bridge. [4] This is essentially due to electrostatic forces, although in aqueous medium the association is driven by entropy and often endothermic. Most salts form crystals with specific distances between the ions; Unlike many other non-covalent interactions, salt bridges are not directional and appear in the solid state, usually only having an interaction determined by the van der Waals radii of the anion. Inorganic as well as organic ions exhibit in water at moderate ionic strength. Similar salt bridge as association G values ​​of about 5 to 6 kJ/mol for 1:1 combinations of anion and cation, almost independent of nature ( size, polarization, etc.) ions. [5]The G values ​​are additive and have an approximately linear function of charges, for example the interaction of a doubly charged phosphate ion with a charged ammonium cation is approximately 2×5 = 10 kJ/mol. The G value depends on the ionic strength I of the solution, as described by the Debye–Huckel equation, at zero ionic strength one observes G = 8 kJ/mol.

dipole-dipole and similar interactions

Regular dipole

Dipole–dipole interactions are electrostatic interactions between molecules with permanent dipoles. This interaction is stronger than the London forces but weaker than the ion–ion interaction because only partial charges are involved. These interactions align the molecules to increase attraction (reduce potential energy). An example of a dipole–dipole interaction can be seen in hydrogen chloride (HCl): the positive end of one polar molecule will attract the negative end of another molecule and affect its position. There is a net attraction between polar molecules. Examples of polar molecules include hydrogen chloride (HCl) and chloroform (CHCl ).

Intermolecular Forces
Intermolecular Forces

Often molecules have dipole groups of atoms, but no overall dipole moment on the molecule as a whole. It is when there is a symmetry within the molecule that causes the dipoles to cancel each other out. It occurs in molecules such as tetrachloromethane and carbon dioxide. The dipole–dipole interaction between two different atoms is usually zero, as the atoms rarely carry permanent dipoles. These forces are discussed below in the section about caspase interactions.

Ion-dipole and Ion-induced dipole force

Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and dipole-induced dipole interactions, but involve ions rather than only polar and nonpolar molecules. Ion-dipole and ion-induced dipole forces are stronger than dipole-dipole interactions because the charge of any ion is much greater than the charge of the dipole moment. The ion-dipole bond is stronger than the hydrogen bond. [6]

An ion and a polar molecule interact in an ion-dipole force. They align so that the positive and negative groups are next to each other, allowing maximum attraction. An important example of this interaction is the hydration of ions in water which gives rise to hydration enthalpy. The polar water molecules surround themselves around the ions in the water and the energy released during this process is called hydration enthalpy. The interaction is of utmost importance in justifying the stability of various ions (eg Cu 2+ ) in water. An ion and a nonpolar molecule interact in an ion-induced dipole force. Like the dipole-induced dipole force, the charge of the ion causes a distortion of the electron cloud on the non-polar molecule. 

Van der Waals force

Van der Waals forces arise from interactions between uncharged atoms or molecules, which lead not only to phenomena such as cohesion of condensed phases and physical absorption of gases, but also to a universal force of attraction between macroscopic bodies There are. [8]

Kisome force

The first contributions to van der Waals forces are due to electrostatic interactions between rotating permanent dipoles, quadrupole (all molecules with symmetry less than a cube), and multipole. This is called the kesome interaction , named after the Willem Hendrik kesome. [9] These forces arise from the attraction between permanent dipoles (dipole molecules) and are temperature dependent. [8]

They consist of attractive interactions between dipoles that are the average of the different rotational orientations of the ensemble dipole. It is assumed that the molecules are constantly in motion and are never locked in place. It’s a good assumption, but at some point the molecules are locked in place. The energy of the kesome interaction depends on the sixth power of distance, in contrast to the interaction energy of two spatially stable dipoles, which depend on the inverse third power of distance. Keysome interaction can occur only between molecules that have permanent dipole moments, that is, two polar molecules. Also kisome interactions are very weak van der Waals interactions and do not occur in aqueous solutions containing electrolytes. The mean angle of interaction is given by the following equation:

{\displaystyle {\frac {-m_{1}^{2}m_{2}^{2}}{24\pi ^{2}\varepsilon _{0}^{2}\varepsilon _{r}^ {2}k_{\text{B}}Tr^{6}}}=V,}

where m = dipole moment, = permittivity of free space, = dielectric constant of the surrounding material, T = temperature, = Boltzmann constant, and r = distance between molecules.

\varepsilon _ 0 \varepsilon _ rk_B

Debye (Permanent-Induced Dipole) Force

The second contribution is the induction (also called polarization) or Debye force, which arises between rotating permanent dipoles and from the polarization of atoms and molecules (induced dipoles). These induced dipoles occur when one molecule repels electrons from another molecule with a permanent dipole. A molecule with a permanent dipole can induce a dipole in a similar neighboring molecule and cause mutual attraction. There cannot be Debye forces between atoms. The forces between the induced and permanent dipoles are not temperature dependent like in kisome interactions because the induced dipole is free to move and move around the polar molecule. The Debye induction effect and the keysome orientation effect are called polar interactions. [8]

Induced dipole force is manifested by induction (also called polarization), which is the attractive interaction between a permanent polypole on one molecule to an induced (formerly by di/multi-pole) 31 . [10] [11] [12] This conversation is called Debbie Force , named after Peter JW Debbie.

An example of an induction interaction between a permanent dipole and an induced dipole is the interaction between HCl and Ar. In this system, Ar experiences a dipole because its electrons are attracted (to the H side of HCl) or repelled (from the Cl side) by HCl. [10] [11] The average interaction of the angle is given by the following equation:

{\displaystyle {\frac {-m_{1}^{2}\alpha _{2}}{16\pi ^{2}\varepsilon _{0}^{2}\varepsilon _{r}^{2}r^{6}}}=V,}

where is = polarization.\alpha

Such an interaction can be expected between any polar molecule and a non-polar/symmetric molecule. The induction–interaction force is much weaker than the dipole–dipole interaction, but stronger than the London dispersion force.

London dispersion force (dipole interaction induced by fluctuations)

The third and major contribution is the dispersion or London force (dipole induced by fluctuations), which arises due to the non-zero instantaneous dipole moments of all atoms and molecules. Such polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in non-polar molecules. Thus, the London interaction is caused by random fluctuations of electron density in an electron cloud. An atom with a large number of electrons will have a greater associated London force than an atom with fewer electrons. The dispersion (London) force is the most important component as all materials are polarizable, whereas the Keysome and Debye forces require permanent dipoles. The London interaction is universal and also exists in atomic-atomic interactions. For various reasons, The London interaction (dispersion) has been shown to be relevant to interactions between macroscopic bodies in condensed systems. Hamemaker developed the van der Waals theory between macroscopic bodies in 1937 and showed that the extravasation of these interactions gives them much longer distances.

Relative strength of forces

bond typeDissociation energy
(kcal/mol) [13]
dissociation energy(kJ/mol)pay attention
ionic lattice२५०-४००० [१४]1100-20000
covalent bond30-260130-1100–
Hydrogen bonding1-12४-५०About 5 kcal/mol (21 kJ/mol) in water
london dispersion force< 1 to 15< 4 to 63Estimated from the enthalpy of vaporization of hydrocarbons [15]
Intermolecular Forces

This comparison is approximate. The actual relative strength will vary depending on the molecules involved. Ionic bonds and covalent bonds in any substance will always be stronger than intermolecular forces.

Effect on the behavior of gases

Intermolecular forces are repulsive at short distances and attractive over long distances (see Lenard-Jones potential). In a gas, the repulsive force mainly prevents two molecules from occupying the same volume. This gives a real gas the tendency to occupy a greater volume than an ideal gas at the same temperature and pressure. The attractive force draws the molecules closer together and gives a real gas the tendency to occupy less volume than an ideal gas. Which interaction is more important depends on temperature and pressure (see compression factor).

In a gas, the distance between the molecules is generally large, so intermolecular forces have only a small effect. The force of attraction is not caused by the force of repulsion, but by the thermal energy of the molecules. Temperature is a measure of thermal energy, so the effect of the attractive force decreases as temperature increases. In contrast, the effect of the repulsive force is essentially unaffected by temperature.

When a gas is compressed to increase its density, the effect of attractive force increases. If the gas is made sufficiently dense, the attraction may be large enough to overcome the tendency for thermal motion to disperse the molecules. The gas can then condense to form a solid or a liquid, that is, the condensed phase. The low temperature favors the formation of a condensed phase. In a condensed phase, there is an almost equilibrium between attractive and repulsive forces.

Quantum mechanical theory

As explained above, the observed intermolecular forces between atoms and molecules can be described as phenomena occurring between permanent and instantaneous dipoles. Alternatively, one may seek a fundamental, unified theory capable of explaining different types of interactions such as hydrogen bonding, van der Waals forces, and dipole–dipole interactions. Typically, this is done by applying the ideas of quantum mechanics to molecules, and the Rayleigh–Schrodinger perturbation theory has been particularly effective in this regard. When applied to existing quantum chemistry methods, such quantum mechanical interpretation of intermolecular interactions provides an array of approximate methods that can be used to analyze intermolecular interactions. [ citation needed ]One of the most useful methods for visualizing such intermolecular interactions that we can find in quantum chemistry is the non-covalent interaction index, which is based on the electron density of the system. London dispersion forces play a big role in this.