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About Various Chemical Bonds
The chemical properties of a substance do not change with the change of its physical state, but the rate of chemical reactions does 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 behavior of matter in
In this unit, we will learn more about these three physical states of matter, particularly liquid and gaseous states.
To begin with, it is necessary to understand the nature of intermolecular forces, molecular interactions, and the effect of thermal energy on the motion of particles because a balance between these determines the state of a substance.
Intermolecular forces are the forces of attraction and repulsion between interacting particles (atoms and molecules).
This term does not include the electrostatic forces that exist between the two oppositely charged ions and the forces that hold atoms of a molecule together i.e., covalent bonds.
Dipole-dipole forces act between the molecules possessing permanent dipole. Ends of the dipoles possess “partial charges” and these charges are shown by the Greek letter delta (8).
Partial charges are always less than the unit electronic charge (1.610-1¹9 C).
The polar molecules interact with neighboring molecules. Fig 5.2 (a) shows electron cloud distribution in the dipole of hydrogen chloride and Fig. 5.2 (b) shows dipole-dipole interaction between two HC1 molecules.
This interaction is stronger than the London forces but is weaker than the ion-ion interaction because only partial charges are involved.
The attractive force decreases with the increase of distance between the dipoles.
As in the above case here also, the interaction energy is inversely proportional to the distance between polar molecules.
The dipole-dipole interaction energy between stationary polar molecules (as in solids) is proportional to 1/r and that between rotating polar molecules is proportional to 1/r 6, where r is the distance
between polar molecules.
Besides dipole-dipole interaction, polar molecules can interact by London forces also. Thus the cumulative effect is that the total intermolecular forces in polar molecules increases.
Dipole-Induced Dipole Forces
This type of attractive force operates between the polar molecules having permanent dipole and the molecules lacking permanent dipole.
The permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electronic cloud (Fig. 5.3).
Thus an induced dipole is developed in the other molecule.
In this case, also interaction energy is proportional to 1/r6 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 have already learned in Unit 4 that molecules of larger size can be easily polarized. High polarisability increases the strength of attractive interactions.
As already mentioned in section (5.1); this is a special case of dipole-dipole interaction. We have already learned about this in Unit 4. 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; species such as Cl may also participate in hydrogen bonding.
The energy of the hydrogen bond varies between 10 to 100 kJ mol–1.
This is quite a significant amount of energy; therefore, hydrogen bonds are a powerful force in determining the structure and properties of many compounds, for example, proteins and nucleic acids.
The 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 molecules.
The following diagram shows the formation of the hydrogen bond.
Assumptions or postulates of the kineticmolecular theory of gases are given below.
These postulates are related to atoms and molecules which cannot be seen, hence it is said to provide a microscopic model of gases.
Gases consist of large number of identical particles (atoms or molecules) that are so small and so far apart on the average that the actual volume of the molecules is negligible in comparison to the empty space between them.
They are considered as point masses. This assumption explains the great compressibility of gases.
There is no force of attraction between the particles of a gas at ordinary temperature and pressure.
The support for this assumption comes from the fact that gases expand and occupy all the space available to them.
Particles of a gas are always in constant and random motion.
If the particles were at rest and occupied fixed positions, then a gas would have had a fixed shape which is not observed.
Particles of a gas move in all possible directions in straight lines. During their random motion, they collide with each
other and with the walls of the container.
Pressure is exerted by the gas as a result of collision of the particles with the walls of the container.
Collisions of gas molecules are perfectly elastic. This means that total energy of molecules before and after the collision remains same.
There may be exchange of energy between colliding molecules, their individual energies may change, but the sum of their energies remains constant.
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State of Matter Class 11 Textbook PDF Free Download