These vapour state of compound such as CH₄, the molecules are well separately and recorded as having little effect of on one another. This consist with one of the postulated of the kinetic theory of gases but in reality, the behaviour of a gas is not ideal because the molecules do interact with one another. We can consider gas law and ideal gases holds for n moles of an ideal gas,
PV= nRT
R= molar gas constant =8. 314k-¹ mol-¹
For a real gas at a given the temperature about PV is a not a constant because, in contrast to the postulated of the kinetic theory of ideal gases.
1) Real gas molecules occupy a volume that can't be ignored, the effective volume of the gas can be corrected from V to (V-nb) where, n is the number of moles of the gas and b is a always constant,
2) Real gas molecules interact with always one another, the pressure has to be corrected from the P to ( P+ an²/v²), where n is the number of moles of the gas and a is a constant.
These are corrections were first proposed by Johannes van der waals in 1873, and this equation for n moles of a real gas. Values of the constant a and b depending on the compound.
(p+an²/v²) (v−nb) =nRT van der waals equation.
Values of the constant a and b for the selected gases. For one mole of the a real gas at temperature T, the pressure of the given volume, V can be calculated as follows,
P= nRT/(V-nb) - an²/v²
The total strength of intermolecular interaction ( van der waals forces or interaction) vary depending upon their precise of nature. The vapour > liquid and liquid >solid phases alsa changes are exothermic and this, in part reflects, and the enthalpy changes associated with formation of intermolecular interaction and must be overcome before phases changes occur. When CH₄ gas, for example is liquefied the molecules come closer together, and when the liquid is solidified and ordered structure is formed in which there are intermolecular interactions between the methane molecules. In this case of Methane is the enthalpy changes associated with fusion and vaporization are small.
CH₄(s) 1kɡ mol−¹→CH₄(l) 8kj mol-¹→CH₄(ɡ)
Above indicating that interaction between methane molecules in the solids and the liquid are weak. This interaction between methane (CH₄) molecules are called London dispersion forces, this weakest type of intermolecular interactions. They are arise from interaction between the electron clouds of adjacent molecules. Now important point to remember is that values of ∆H (melting point) and ∆H (boiling point) as well as boiling and melting points of atomic species, examples : He, Ar, and molecular species example : CH₄, H₂O, N₂ and C₂H₅OH reflect the extent of intermolecular interaction. In ionic solid (eg) NaCl in which ions interact with one another through electrostatic forces, the total amount of the energy is needed to separate the ions is often far greater than that needed to separate covalent molecules, enthalpies of fusion of an ionic solids are significantly higher than those of the molecular solids.
1)London dispersion forces - most molecules, ≤2 kj/mol-¹
2) Dipole - dipole interaction - polar molecules, 2 kj/mol-¹
3) Ion-dipole interaction - Ions and polar molecules, 15kj/mol-¹
4) Hydrogen bonds- electronegative atom usually N, O or F and H atom attached to another electronegative atom, 5-30b, kj/mol-¹
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