State Of Matter (Gases and Liquids):
Matter exists in three different physical states, namely solid, liquid and gaseous states. The important characteristics of three states are:
Solid state: A solid possesses a definite size (volume) and a definite shape. The shape of a solid can be changed but it usually requires considerable force. Therefore, solids are generally hard and rigid. Some common examples of solids are iron, silver, common salt, etc.
Liquid state: A liquid possesses a definite volume but not definite shape. It takes up the shape of the container in which it is placed. Liquid also has a tendency to flow. For example, water, alcohol, milk, oil, etc.
Gaseous state: A gas neither possesses a definite volume nor definite shape. A gas occupies the whole of the volume of the vessel in which it is placed. It also takes up the shape of the container. For example, air, carbon dioxide, oxygen, hydrogen, etc.
Note: The state at which all the three phases exist simultaneously is called Triple point.
Intermolecular Forces:
The forces of attraction and repulsion between interacting particles (atoms or molecules) are called intermolecular forces. Such forces exist in all states of matter and are responsible for many structural features and physical properties of matter. If the particles are atoms these may be termed as interatomic forces, on the other hand, if these are molecules, they are called intermolecular forces. These forces are collectively called Van der Waal’s forces. Van der Waal’s forces are further classified as:
- Dispersion forces or London forces
- Dipole-dipole forces
- Dipole-induced dipole forces
- Hydrogen Bonding
1) Dispersion forces or London forces:
London forces is the interparticle forces among the monoatomic or non-polar molecules such as N2, H2, CO2, etc. Atoms or non-polar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed. However, an instantaneous dipole may develop in such atoms and molecules.
Suppose we have two atoms of neon ‘A’ and ‘B’ in the close vicinity of each other. Rapid movement of electrons in A may cause momentary accumulation of electron density on one side thus causing the charge distribution to become unsymmetrical. In other word, the charge cloud is more concentrated on one side than the other. This will result in the development of temporary instantaneous dipole on the atom ‘A’ for a very short time. This instantaneous dipole distorts the electron density of the atom ‘B’, which is close to it. In other words, a dipole is induced in the atom ‘B’ also.
The temporary dipoles of atom ‘A’ and ‘B’ attract each other. Similar temporary dipoles are also induced in non-polar molecules. Magnitude of such a force of attraction was first calculated by the German physicist Fritz London. For this reason, force of attraction between two temporary dipoles is known as London force.
2) Dipole-Dipole forces:
Dipole-dipole forces operate between the molecules possessing permanent dipole. Ends of the dipole possess “partial charges” and these charges are shown by Greek letter delta (δ). Partial charges are always less than the unit charge (1.6 x 10-19 C) because of electron sharing effect. The polar molecules interact with neighbouring molecules. This interaction is weak as compared to ion-ion interaction because only partial charges are involved. The interaction energy decreases with the increase of distance between the dipoles. Besides dipole-dipole interaction, polar molecules can interact by London forces. The cumulative effect of both the interactions is the total increase of intermolecular forces in polar molecules. The electron cloud distribution and attractive interactions between H—Cl dipoles are shown below.
3) Dipole-Induced Dipole Forces:
This type of attractive forces operates between polar molecules having µ>0 and the non polar molecules having µ = 0. Permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electron cloud. Induced dipole moment depends upon the dipole moment of the permanent dipole and the polarizability of the electrically neutral molecule. Molecules of larger size can be easily polarized. Higher polarizability increases the strength of attractive interactions.
4) Hydrogen Bonds:
This is a special type of dipole-dipole interaction operating between the molecules in which hydrogen atom is covalently bonded to a highly electronegative atom such as fluorine (F), oxygen (O) or nitrogen (N). The bond pair of electrons forming covalent bond is displaced towards electronegative atom. When solitary electron of H atom lies away, it behaves almost as a bare proton and exerts a strong electrostatic force of attraction on the electronegative atom of the other molecule in its vicinity.
This interaction is represented by a dotted line and is called hydrogen bond. Energy of hydrogen bond varies between 10 to100 k J/mol. This is very significant amount of energy; therefore hydrogen bonds are powerful force in determining the structure and properties of many compounds. Formation of hydrogen bond is depicted as follows
………H— F……. H— F…….. H— F……. H— F…… (H-bonding in Hydrogen fluoride)