Atoms and compounds participate in weak electrostatic interactions.
The strength of these interactions impacts physical properties like melting and boiling points.
Types of Intermolecular Interactions
Dispersion Forces (London Forces):
The weakest intermolecular interactions.
Dipole-Dipole Interactions:
Intermediate strength.
Hydrogen Bonds:
The strongest type of interaction (but still weak compared to covalent bonds).
A misnomer: there is no actual sharing or transfer of electrons.
Only about 10% of the strength of a covalent bond.
Can be overcome with small or moderate amounts of energy.
London Dispersion Forces
In nonpolar covalent bonds, electrons appear to be shared equally, but at any given moment, electron density may be unequally distributed.
This results in rapid polarization and counter-polarization of the electron cloud and the formation of short-lived dipole moments.
These dipoles interact with the electron clouds of neighboring compounds, inducing the formation of more dipoles.
The momentarily negative end of one molecule causes the closest region in a neighboring molecule to become temporarily positive.
The attractive or repulsive interactions of these short-lived and rapidly shifted dipoles are known as London dispersion forces, a type of Van der Waals force.
Dispersion forces are the weakest of all intermolecular interactions because they are the result of induced dipoles that change and shift moment to moment.
They do not extend over long distances and are significant only when molecules are in close proximity.
The strength of the London force depends on the degree and ease by which the molecules can be polarized (how easily the electrons can be shifted around).
Large molecules are more easily polarizable than comparable small molecules and thus possess greater dispersion forces.
Without dispersion forces, noble gases would not liquefy at any temperature.
The low temperatures at which noble gases liquefy are indicative of the very small magnitude of the dispersion forces between the atoms.
Dipole-Dipole Interactions
Polar molecules tend to orient themselves such that oppositely charged ends of respective molecular dipoles are closest to each other.
The positive region of one molecule is close to the negative region of another molecule.
This arrangement is energetically favorable because an attractive electrostatic force is formed between the two molecules.
This attractive force is denoted by dashed lines in most molecular notations and indicates a temporary bond and interaction.
Dipole-dipole interactions are present in the solid and liquid phases but become negligible in the gas phase because of the significantly increased distance between gas particles.
Polar species tend to have higher melting and boiling points than nonpolar species of comparable molecular weight due to these interactions.
London forces and dipole-dipole interactions are different not in kind, but in duration.
Both are electrostatic forces between opposite partial charges.
The difference is only in the transients or permanence of the molecular dipole.
Hydrogen Bonds
A specific, unusually strong form of dipole-dipole interaction that may be intra- or intermolecular.
Not actually bonds; there is no sharing or transferring of electrons between two atoms.
When hydrogen is bonded to one of three highly electronegative atoms (fluorine, oxygen, or nitrogen), the hydrogen atom carries only a small amount of the electron density in the covalent bond.
The hydrogen atom essentially acts as a naked proton.
The positively charged hydrogen atom interacts with the partial negative charge of fluorine, oxygen, or nitrogen on nearby molecules.
Substances that display hydrogen bonding tend to have unusually high boiling points compared to compounds of similar molecular weights that do not exhibit hydrogen bonding because of the energy required to break the hydrogen bonds.
Hydrogen bonding is particularly important in the behavior of water, alcohols, amines, and carboxylic acids.
Many biochemical molecules such as nucleotides have different regions that are stabilized by hydrogen bonding.
If water could not form hydrogen bonds and exist in the liquid state at room temperature, life as we know it would not exist.
Conclusion
Atoms partner together to form compounds by exchanging electrons to form ions (held together by electrostatic attractions between opposite charges) or by sharing electrons to form covalent bonds.
Covalent bonds have relative lengths, energies, and polarities.
Lewis dot structures and VSEPR theory are used for predicting likely bond arrangements, resonance structures, and molecular geometries.
Even the strongest intermolecular electrostatic interactions (hydrogen bonds) are much weaker than actual covalent bonds.