Intermolecular Forces
Forces which determine physical and chemical properties of substance categorized into two types.
1. Intramolecular forces
Intramolecular forces are within the molecule and these can be between opposite charges ions, between nucleus and electrons or between metal ions and delocalized electrons. Intramolecular force of attractions are bonding force of attractions. These forces explains chemical properties of substance. Intramolecular forces are relatively strong because the charges are close and magnitude of the charge is more.
– Ionic bonding: Electrostatic forces between cation and anion.– Covalent bonding: Force of attraction between nuclei and shared electron pairs.– Metallic bonding: Force of attraction between metal ions and sea of delocalized electrons.
2. Intermolecular force of attractions
Intramolecular force of attractions are non- bonding in nature and these are present between the molecules with partial charges and between ion and a molecule. Intermolecular forces are very weak when compared with intramolecular forces because charge magnitude is small and far from each other and these forces explains physical properties of substance. Strength of intermolecular forces varies with distance. Intermolecular forces also called as van der Waals forces.
First we need to understand different distances which determines relative strength of forces. Distances in intermolecular forces are: ion- dipole forces, dipole -dipole forces, H – bonding, Dispersion forces (London forces) and the distances in intramolecular forces are: covalent radii, ionic radii, metallic radii.
Ion – Dipole forces
Force between ion and a dipole (of polar molecule) is ion – dipole force. Example: NaCl in water. Na+ ions surrounded by H2O molecules and oxygen with partial negative charge is pointing toward Na+ and hydrogen with partial positive charge pointing towards Cl– ions. Ion-dipole attractions are inversely proportional to the square of the distance between ion and dipole. (F 𝛼 1/r2).
Dipole – Dipole forces
Attraction forces between partial positive end partial negative end of a dipole are dipole-dipole forces. The strength of dipole-dipole forces depends on the magnitude of dipole moment. Dipole-dipole attractions in acetone are stronger than those exist in methanol. Therefore, boiling point of acetone is more than that of methanol. Energy of interaction of dipole-dipole forces is inversely proportional to the cube of the distance (F 𝛼 1/r3). Therefore, dipole-dipole interactions are weaker than ion-dipole forces(which are inversely proportional to the square of the distance).
Induced dipole interactions
If a charge particle (example: cation or anion) is close to the non-polar molecule (example: Ne atom), then charged particle induces charges in the non-polar molecule by distorting electron cloud of the non-polar molecule to one end of the molecule. The energy interaction is inversely proportional to r4 (F 𝛼 1/r4). One dipole can induce charges into another uncharged non-polar molecule and in this case energy of interaction is inversely proportional to r6(F 𝛼 1/r6). These two interactions are very weak. The power of ‘r’ in energy interactions is high (4 and 6) therefore, these energy interactions are effective when the distances are small (very short range and extremely short range interactions).
Dispersion forces (London forces)
Attractive forces between non-polar molecules are Dispersion or London forces. London forces exist between any two particles (atoms, ions and molecules).
In a neutral atom, for example Ne electrons are uniformly distributed and in constant motion. At any moment electron cloud comes one side of the atom and the nucleus other side. Therefore, it forms an instantaneous dipole. The resulting instantaneous dipole induces dipole near by ‘Ne’ atom and both dipoles attract each other.This process continues to all other atoms. These dispersion forces are instantaneous dipole-induces dipole forces.
Dispersion forces are universal and operates between all the molecules. Dispersion forces are dominant even than dipole -dipole forces(Weak dipole dipole forces) but stronger dipole-dipole forces and H-bonding are dominant than dispersion forces. Example: In HCl 85% force of attractions due to dispersion forces and 15% forces due to dipole-dipole attractions and in H2O 75% force of attractions are due to Hydrogen bonding and 25% forces are due to dispersion forces.
Strength of dispersion forces depending on the relative polarizability of the atom. Polarizability is the distortion of electron cloud and electron cloud distorts easily if is it not strongly held by the atom. As the size of atom increases(In a group), atom do not held electron cloud strongly and polarizes easily. Across the period effective nuclear charge increases and electrons cloud strongly held by the nucleus which decreases polarizability.
Polarizability increases down the group and decreases across the period.
Polarizability depends on the number of electrons. Therefore, as the molecular mass increases strength of Dispersion forces increases.
Strength of Dispersion forces Molecular mass. Therefore, boiling point increases.
Examples:
Boiling point of: He > Ne > Ar > Kr > Xe > Rn.
Boiling point of: CH3OH > CH3CH2OH. Because between these two strength of Hydrogen bond is not much different. Ethanol dispersion forces are stronger than methanol.
Hydrogen bond
Hydrogen bond is formed when hydrogen atom with partial positive charge is attracted with the partial negative charge on electronegative atom F, O and N. In hydrogen bond —A—H …B — where A and B are N, O and F. There are two types of hydrogen bonding: Intermolecular hydrogen bonding and intramolecular hydrogen bonding.
Intermolecular hydrogen bonding
Hydrogen bonding between different molecules is the intermolecular H – bonding. Examples include hydrogen bonding in H2O, HF, p -nitrophenol etc. Maximum number of H – bonds formed by the water molecule is 4.
Intramolecular hydrogen bonding
Hydrogen bonding within the. Molecule is Intramolecular hydrogen bonding. Example: Hydrogen bonding in o – nitrophenol.
Important consequences of Hydrogen bonding
1. Boiling point of H2O is more than that of H2S due to the intermolecular H – bonding in water
2. Due to intermolecular hydrogen bonding in ice, it gets structure of hexagons and floated on water due to the lower density than liquid H2O.
3. Molecules which are able to form hydrogen bonding with water are relatively more soluble in water. Example: Glucose is soluble in water where as benzene do not.