Chapter 14, Problem 58A

Interpretation Introduction

Interpretation: The term London dispersion forces and their appearance in a non-polar molecule is to be explained and compared with dipole-dipole attraction and covalent bonds.

Concept introduction: The existence of weak attractive forces among the non-polar molecule was given by Vander-waal forces.

Answer to Problem 58A

The forces with which the induced momentary dipoles attract each other are called London dispersion forces.

These forces are thought to arise from the motion of the electrons. Vanderwaal forces or London dispersion forces are weaker than the dipole-dipole force of attraction. Covalent bonds are stronger than London dispersion forces because of the intramolecular forces of attraction present between them.

Explanation of Solution

These forces are arisen due to the distortion of electronic cloud of molecule and a momentary dipole is produced. This momentary dipole induces dipole in the neighboring molecules and these are then attracted to each other. These induced momentary dipoles attract each other and produces a force called as London dispersion force.

London dispersion forces are weaker than dipole-dipole interaction because the molecular forces attraction are present between non-polar molecules like $H_2$ , $O_2$ and $N_2$

Intramolecular forces are present within the molecule itself. The polar covalent bond is much stronger than London dispersion forces because of the intra molecular forces of attraction present between them.

The polar molecules have positive and negative charges on their constituent atoms which attract each other with a strong force where London forces are momentary forces. These occur in only non-polar molecules. Therefore, London forces are weaker than dipole-dipole force.

Conclusion

Due to the intramolecular forces of attraction present between the molecules, the extent of overlap is more within the molecule whereas in case of London dispersion forces, the weakest intermolecular forces of attraction are present between them.