# Lecture 21: Intermolecular Forces

• Intermolecular Forces - Electrostatic interactions between molecules
• Atoms within a molecule make covalent and ionic bonds within each other, but molecules also have interactions with other molecules
1. Ion ﻿$\rightarrow$﻿ Ion Interactions
2. Strongest intermolecular forces because they involve formal charges
3. Ion ﻿$\rightarrow$﻿ Dipole Interactions
4. Dipole - A side of an electron with excess and a side of deficiency
5. The partially negative side is attracted to positive ions and vice versa
6. Each ion can make several of these interactions which store energy
7. Dipole ﻿$\rightarrow$﻿ Dipole Interactions
8. Hydrogen bonds - OH, FH, or NH bonds - interact with each other
9. Very strong dipole - Dipole bond because they are the most electronegative elements ﻿$\rightarrow$﻿ Most polarized bonds ﻿$\rightarrow$﻿ Strong Dipole ﻿$\rightarrow$﻿ Strong Dipole (Dipole Interactions)
10. The greater the partial charge - the greater the interaction, but never as strong as formerly charged particles
11. Van-der Waals/London Dispersion Force
12. Any substance can have this bond
13. Ion - Ion > Ion - Dipole > Dipole - Dipole > Van der Waals
14. The stronger the forces between the molecules, the more heat energy is needed to provide to melt and boil the sample
15. Non-Polar Covalent = Van der Waals ﻿$(CH_4)$﻿
16. Polar Covalent = *check geometry* ﻿$(CO_2,H_2O)$﻿
17. (﻿$Bf_3$﻿ (non-polar - van der waals), ﻿$NH_3$﻿ (polar-dipole-dipole))
18. (﻿$CF_4$﻿ (non-polar - van der waals), ﻿$CH_3F$﻿ (polar-dipole-dipole))
19. Ionic = Ion-Ion ﻿$(NaCl)$﻿
• Polar Molecules
• Non-Polar Molecules
• H-H
• In a moment the electrons can be close to one H (temporary) and gives one H
• Temporary Dipole ﻿$\rightarrow$﻿ Has a force of attraction with other molecules ﻿$\rightarrow$﻿ London Dispersion Force (weak bands)
• Ideal Gas - ﻿$PV = nRT$﻿, no attraction between molecules and there is infinite amount of space between them
• Real Gas - Attraction between molecules

#### dispersion forces

• Fluctuations in the electron distribution in atoms and molecules result in a temporary dipole
• Region with excess electron density has a partial (-) charge
• Region with depleted electron density has partial (+) charge
• The attraction forces caused by these temporary dipoles are called dispersion forces = London forces
• Magnitude depends on
• Polarity of the electron cloud
• Large molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions
• Shape of the molecule
• More surface - surface contact = larger induced dipole = stronger attraction
• Larger molecules have more electrons ﻿$\rightarrow$﻿ Increased polarizability

#### effect of molecular size on size of dispersion force

• Noble gases are all non-polar atomic elements
• As the molar mass increases, the number of electron increases
• Strength of dispersion forces increases
• The stronger the attractive forces between molecules the higher the boiling point will be
• If intermolecular forces are weaker, vapour pressure is higher
• Boiling point is lower
• Vapour Pressure ﻿$\rightarrow$﻿
• Intermolecular Force ﻿$\rightarrow$﻿
• Boiling Point ﻿$\rightarrow$﻿
• Colligative properties depend only on the number of dissolved particles in solution and not on their identity
• Non-colligative properties depend on the identity of the dissolved species and the solvent
• Vapour pressure of solutions
• The vapour pressure of a solvent above solution is lower than the vapour pressure of the pure solvent
• The solute particles replace some of the solvent molecules at the surface

#### raoult's law

• ﻿$P$﻿ solvent solution = ﻿$X$﻿ solvent
• ﻿$P$﻿ solution is lower than ﻿$P$﻿ solvent because the mole fraction is always less than 1
• ﻿$P$﻿ : Lowering of vapour pressure
• The vapour pressure of the solution is directly proportional to the amount of the solvent in the solution
• Boiling point
• Temperature at which the vapour pressure is equal to atmospheric pressure
• Vapour pressure lowering occurs at all temperatures
• Results in the temperature required to boil the solution being higher than the boiling point of the pure solvent
• Also results in the temperature required to freeze the solution lower than the freezing point of the pure solvent
• Vant Hoff Factors
• ﻿$i$﻿ is the ratio of moles of solute particles to moles of formula units dissolved
• The measured Vant Hoff factors are generally less than the theoretical due to ion pairing in solution
• Therefore, the measured Vant Hoff factor often causes the ﻿$T$﻿ to be lower than expected
• Vapour pressure of the solution is lower than that of the solvent
• Boiling point of a solution is higher than that of a solvent