topic 7: intermolecular forces

formation of london forces

(instantaneous dipole-induced dipole interactions / dispersion forces / temporary dipole-induced dipole interactions

• [molecule] is non-polar

• Random movement of electrons in [name] molecules result in uneven distribution of electron. Partial positive and negative charges are created in the ends of a non-polar molecule instantaneously.

• Partial positive charges of [name] molecule attracts the electrons of another non-polar molecule and induce the partial charges on another non-polar molecule.

• Partial positive end of a non-polar molecule is attracted by partial negative end of another non-polar molecule.

formation of permanent dipole-permanent dipole interactions

[name] is a polar molecule. Partial positive [atom] atom in [name] molecule is attracted by partial negative [atom] atom in [name] molecule.

formation of permanent dipole-induced dipole interactions between $H_2S$ and $N_2$ molecules

• electrons in $N_2$ are drawn towards the partial positive charge on $H$ atom on $H_2S$

• due to the permanent dipole in $H_2S$, an induced dipole is formed in $N_2$

• there is attraction between positive charge of $H_2S$ and induced negative charge of $N_2$ molecule

determining strength of intermolecular forces

1. no. of Hydrogen bonds per molecule (lone pair on N/O/F + h atom bonded to N/O/F)

2. electronegativity (LF; F > O > N)

3. molecular size (LF ; no. of $e^{-}$ )

4. polarity (pd-pd ; EN)

5. contact surface area (LF ; spherical vs rod-shape)

formation of hydrogen bonds

N/O/F is a highly electronegative atom. H—N/O/F bond is very polar. Highly partial positive charged H atom attract the lone pair electrons on N/O/F atom of another molecule.

soluble requirement

solute-solvent interaction similar to / stronger than solute-solute and solvent-solvent interactions

solubility of alcohol comparison

1. number of hydrogen bonds formed with water

2. size of non-polar part

• hydrophobic alkyl group

• carbon chain length ↑ ; water solubility ↓

• molecular size ↑ ; no. of electron ↑; london force ↑ ; water solubility ↓

branched vs unbranched with hydrogen bonds boiling point and solubility comparison

more branched ⇒ weaker hydrogen bond

• more bulky molecules more steric hindrance on the formation of hydrogen bond

• more difficult to form intermolecular hydrogen bond and strength of hydrogen bond decrease

• lower boiling point, viscosity and solubility

intramolecular H bonds effect on intermolecular H bond

cis vs trans

• cis forms intramolecular hydrogen bond ⇒ form less intermolecular hydrogen bond with other molecule

• while trans will not

• trans forms more extensive intermolecular hydrogen bond

• trans higher melting point

why ice density lower than liquid water

In ice:

• The water molecules form hydrogen bond tetrahedrally and are packed in an open structure

• Empty spaces between molecules prevent the molecules from getting too close to each other

In liquid water:

• Water molecules have relative motion and leads to hydrogen bonds between water molecules break and form continuously

• Open structure collapses and the water molecules become more closely packed

shapes of molecules

non-polar shapes

1. trigonal planar

2. tetrahedral

3. trigonal bipyramidal

4. octahedral

polar shapes

1. v-shaped

2. trigonal pyramidal

3. tetrahedral