`=>` Most of the observable characteristics of chemical systems represent bulk properties of matter, i.e., the properties associated with a collection of a large number of atoms, ions or molecules.

● For example, an individual molecule of a liquid does not boil but the bulk boils.

`=>` Collection of water molecules have wetting properties; individual molecules do not wet.

`=>` Water can exist as ice, which is a solid; it can exist as liquid; or it can exist in the gaseous state as water vapour or steam.

● Physical properties of ice, water and steam are very different.

● In all the three states of water, chemical composition of water remains the same i.e., `H_2O`.

`=>` Characteristics of the three states of water depend on the energies of molecules and on the manner in which water molecules aggregate. Same is true for other substances also.

`=>` Chemical properties of a substance do not change with the change of its physical state; but rate of chemical reactions do depend upon the physical state.

`=>` Many times in calculations while dealing with data of experiments we require knowledge of the state of matter.

● Therefore, it becomes necessary for a chemist to know the physical laws which govern the behaviour of matter in different states.

`color{purple}♣ color{Violet} " Just for Curious"`
•The temperature at which all the three phases exist together is called triple point.
•Besides the three states of matter, there exists fourth state called plasma state(mixture of electrons and positively charged ions formed due to superheating of the gaseous state,e.g., in the sun or stars) and fifth state of supercooled solid in which atoms lose their identity and condense to form a single atom.

`=>` Atoms and non-polar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

● But a dipole may develop momentarily even in such atoms and molecules.

`=>` If two atoms ‘A’ and ‘B’ are in the close vicinity of each other (Fig.1a).

● It may so happen that momentarily electronic charge distribution in one of the atoms, say ‘A’, becomes unsymmetrical i.e., the charge cloud is more on one side than the other (Fig.1 b and c).

● This results in the development of instantaneous dipole on the atom ‘A’ for a very short time.

● This instantaneous or transient dipole distorts the electron density of the other atom ‘B’, which is close to it and as a consequence a dipole is induced in the atom ‘B’.

● The temporary dipoles of atom ‘A’ and ‘B’ attract each other.

`=>` In the same way, temporary dipoles are induced in molecules also.

`=>` This force of attraction was first proposed by the German physicist Fritz London, and for this reason force of attraction between two temporary dipoles is known as `text(London force)`.

● Another name for this force is `text(dispersion force)`.

`=>` These forces are always attractive and interaction energy is inversely proportional to the sixth power of the distance between two interacting particles (i.e., `1//r^6` where `r` is the distance between two particles).

`=>` These forces are important only at short distances (`~500` pm) and their magnitude depends on the polarisability of the particle.

`=>` Dipole-dipole forces act between the molecules possessing permanent dipole.

`=>` Ends of the dipoles possess “partial charges” and these charges are shown by Greek letter delta (`δ`).

● Partial charges are always less than the unit electronic charge (`1.6×10^(-19) C`).

● The polar molecules interact with neighbouring molecules.

`=>` Fig 2 (a) shows electron cloud distribution in the dipole of hydrogen chloride and Fig. 2 (b) shows dipole-dipole interaction between two `HCl` molecules.

`=>` This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved.

`=>` The attractive force decreases with the increase of distance between the dipoles.

`=>` The interaction energy is inversely proportional to distance between polar molecules.

● Dipole-dipole interaction energy :

(i) between stationary polar molecules (as in solids) is proportional to `1//r^3`

(ii) between rotating polar molecules is proportional to `1//r^6,`

where `r` is the distance between polar molecules.

`text(Note :)` (i) Besides dipole-dipole interaction, polar molecules can interact by London forces also.

(ii) Therefore, the total of intermolecular forces in polar molecules increase.

`=>` This type of attractive forces operate between the polar molecules having permanent dipole and the molecules lacking permanent dipole.

`=>` Permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electronic cloud (Fig.3).

● Thus an induced dipole is developed in the other molecule.

`=>` In this case also interaction energy is proportional to `1//r^6` where `r` is the distance between two molecules.

`=>` Induced dipole moment depends upon the dipole moment present in the permanent dipole and the polarisability of the electrically neutral molecule.

● We know hat molecules of larger size can be easily polarized.

● High polarisability increases the strength of attractive interactions.

`=>` In this case also cumulative effect of dispersion forces and dipole-induced dipole interactions exists.

`=>` This is special case of dipole-dipole interaction.

`=>` This is found in the molecules in which highly polar `N–H`, `O–H` or `H–F` bonds are present.

● Although hydrogen bonding is regarded as being limited to `N`, `O` and `F`; but species such as `Cl` may also participate in hydrogen bonding.

`=>` Energy of hydrogen bond varies between `10` to `100 kJ mol^(–1)`.

● This is quite a significant amount of energy; therefore, hydrogen bonds are powerful force in determining the structure and properties of many compounds, for example proteins and nucleic acids.

`=>` Strength of the hydrogen bond is determined by the coulombic interaction between the lone-pair electrons of the electronegative atom of one molecule and the hydrogen atom of other molecule.

● Following diagram shows the formation of hydrogen bond.

`overset(delta+)H - overset(delta-)H ........... overset(delta+)H - overset(delta-)F`


`text(Note :)` Intermolecular forces discussed so far are all attractive.

`=>` Molecules exert repulsive forces on one another.

● When two molecules are brought into close contact with each other, the repulsion between the electron clouds and that between the nuclei of two molecules comes into play.

● Magnitude of the repulsion rises very rapidly as the distance separating the molecules decreases.

● This is the reason that liquids and solids are hard to compress.

● In these states molecules are already in close contact; therefore they resist further compression; as that would result in the increase of repulsive interactions.