intermolecular forces

intermolecular forces

forces that attract one molecule to another

• induced dipole-dipole interactions (London forces)

• permanent dipole-dipole interaction

• hydrogen bonding

London forces

• features and what

induced dipole-dipole interactions

electrons are always moving even in nonpolar bonds - on average they are shared evenly

→ at any one instant in time they may be closer to on atom than the other

= temporary δ- or δ+

movement produces an instantaneous dipole (separation of charge across a bond)

position of dipole is constantly changing - on average electrons are evenly shared

can induce a dipole in a neighbouring molecule

• exist between all non-polar and all polar molecules - EVERYTHING has London forces

weakest intermolecular force but can be stronger than permanent dipole dipole if lots of electrons

factors affecting the strengths of London forces

• shape of molecule

• number of electrons in the molecule

as the atom size incr.= the electron no. incr.= London forces strength incr.

= more energy to overcome london forces = boiling point incr.

permanent dipole-dipole interactions

polar molecules have permanent dipoles

one bonded atom = more electronegative than the other = dipole

occur in addition to London forces

HCl molecules are attracted together by London forces and permanent dipole dipole interactions

= stronger intermolecular forces

bond enthalpy

how strong the bonds are

hydrogen bond

an intermolecular bond between molecules containing O,N and F and the H atom of-OH,-NH or -FH

H must be bonded to a very electronegative atom (N, O, or F) to have sufficient positive

where is lone pair that H bonds to found

attracted to the lone pair - found in n=2 orbital → small orbitals

= don’t occupy much space

=high density of negative charge

strongest intermolecular force

hydrogen bonds → relatively strong

hydrogen bonds impact on solubility

increase in solubility if the molecule can form H bonds with the water (contains O,N or F)

Methanol vs Methanal (2 per)

Methanol -

form H bonds with itself = high BP

can form H bonds with water = solublility incr

Methanal -

no delta pos H = cant form H bonds with itself

can form H bonds with water = incr solubility

anomalous properties of water (2)

• relatively high MP and BP

• solid ice is less dense than liquid water

water has a relatively high MP and BP , why ? (3)

water molecules → relatively strong H bonds

also London forces and p.d.d interactions

H bonds require a significant amount of energy to overcome

ice is less dense than water (2)

Held apart by hydrogen bonds

= open lattice

MP and BP of simple molecular substances (4)

• atoms within the molecule are held together by strong covalent bonds

• molecules are held together by weak intermolecular forces

to melt or boil, IMFs need to be overcome

not much energy

(large molecules = solid at RT, relatively strong London forces)

solubility of simple molecular substances (2)

• like dissolves like

polar molecules tend to dissolve in polar solvents (water)

non-polar molecules tend to dissolve in non-polar solvents (hexane)

• solubility massively improved if molecule can form H bonds with the solvent

electrical conductivity of simple molecular substances

do not have mobile, charged particles

must have charged particles that are free to move and carry current

do not conduct