Why the intermolecular force becomes attractive with the increase of the Intermolecular Distance?
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Forces involved
Physisorption involves weak intermolecular forces such as London–van der Waals forces and hydrogen bonding. These forces act simultaneously in combinations permitted by the structures of the interacting species. London–van der Waals forces include dipole–dipole, dipole–induced dipole (attraction of a permanent dipole with the dipole it induces in its neighbor), and induced dipole–induced dipole (mutual attraction of momentary dipoles produced by the synchronization of electronic motion in interacting neighbors). Additionally, the interaction of a charge with a neutral molecule may involve charge–dipole and charge–induced-dipole forces. The strengths of London–van der Waals interactions depend on the molecular area of contact, separation distance (to the inverse sixth power), and applicable molecular properties of the interacting species such as dipole moment, ionization potential, and polarizability.
The most common type of hydrogen bond is between an acidic proton and a lone pair of electrons (i.e., XH⋯ : Y) which are located on O, N, or S atoms. This bond is highly oriented (XHY angle, 180° ± 15°), ranges in enthalpy from 12 to 35 kJ mol−1, and involves a combination of dipole–dipole attraction and molecular orbital overlap. Other much weaker hydrogen bonds are known, such as the one formed between an acidic proton and the π system of an aromatic ring, but whether they are important in soil systems in competition with hydrogen bonds from water molecules is unclear.
Physisorption involves weak intermolecular forces such as London–van der Waals forces and hydrogen bonding. These forces act simultaneously in combinations permitted by the structures of the interacting species. London–van der Waals forces include dipole–dipole, dipole–induced dipole (attraction of a permanent dipole with the dipole it induces in its neighbor), and induced dipole–induced dipole (mutual attraction of momentary dipoles produced by the synchronization of electronic motion in interacting neighbors). Additionally, the interaction of a charge with a neutral molecule may involve charge–dipole and charge–induced-dipole forces. The strengths of London–van der Waals interactions depend on the molecular area of contact, separation distance (to the inverse sixth power), and applicable molecular properties of the interacting species such as dipole moment, ionization potential, and polarizability.
The most common type of hydrogen bond is between an acidic proton and a lone pair of electrons (i.e., XH⋯ : Y) which are located on O, N, or S atoms. This bond is highly oriented (XHY angle, 180° ± 15°), ranges in enthalpy from 12 to 35 kJ mol−1, and involves a combination of dipole–dipole attraction and molecular orbital overlap. Other much weaker hydrogen bonds are known, such as the one formed between an acidic proton and the π system of an aromatic ring, but whether they are important in soil systems in competition with hydrogen bonds from water molecules is unclear.
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There are two kinds of forces, or attractions, that operate in a molecule—intramolecular and intermolecular.
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