1. Chapter 11 Intermolecular Forces, Liquids, and Solids

2. 11.1 Comparing Gases, Liquids, and Solids
The fundamental difference between states of matter is the distance between particles.

3. Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

4. The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:
The kinetic energy of the particles
The strength of the attractions between the particles

5. 11.2 Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.
They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.
These intermolecular forces as a group are referred to as van der Waals forces.

6. van der Waals Forces Dipole-dipole interactions Hydrogen bonding
London dispersion forces

7. 1. Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
These forces are only important when the molecules are close to each other.

8. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.

9. 2. London Dispersion Forces
There can be no dipole-dipole forces between nonpolar atoms and molecules.
But, the motion of electrons in an atom or molecule can create an instantaneous, or momentary, dipole moment.
In He atoms, for example, the average distribution of electrons around a nucleus is spherically symmetrical. The atom has no dipole moment.

10. However, an instantaneous distribution of the He 1s orbital electrons may find both electrons on one side of the nucleus.
In that instant, the atom would have an instantaneous dipole moment.
The helium atom is now polar, with an excess of electrons on one side as compared to the other.

11. The temporary dipole moment on one atom can induce a similar temporary dipole – an induced dipole -- on an adjacent atom, causing the atoms to be attracted to each other.
London dispersion forces (or dispersion forces) are attractions between an instantaneous dipole and an induced dipole.
These forces, like dipole-dipole forces, is significant only when molecules are close together.

12. Factors Affecting Dispersion Forces
The strength of the dispersion force depends on the ease with which the charge distribution in a molecule can be distorted to induce a momentary dipole.
The tendency of an electron cloud to distort in this way is called polarizability.
The more polarizable a molecule is, the greater its dispersion forces.
Dispersion forces tend to increase in strength with increasing molecular weight.
Larger molecules tend to be more polarizable because of a greater number of electrons located further from the nucleus.

13. The effect of molecular weight and size on boiling points is shown below:
Dispersion forces operate between all molecules, whether they are polar or not.
Polar molecules can experience dipole-dipole and dispersion forces at the same time.
Dispersion forces between polar molecules commonly contribute more to intermolecular attractions than do dipole-dipole forces.

14. The shape of the molecule also affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
This is due to the increased surface area in n-pentane.

15. Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces?
If two molecules of a substance are of comparable size and shape, dipole-dipole interactions will likely be the dominating force.
If one molecule of a substance is much larger than another, dispersion forces will likely determine its physical properties.

16. 3. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
These interactions are called hydrogen bonds.
Due to hydrogen bonds, there are very strong intermolecular forces between the H and O atoms on adjacent water molecules.
Hydrogen bonding is the cause of water’s
high boiling point
high melting point
high specific heat
high heat of vaporization

17. Boiling point as a function of molecular weight.
The nonpolar series (SnH4 to CH4) follow the expected trend.
The polar series follows the trend from H2Te through H2S.
Water, despite its relatively low molecular weight, is an anomaly due to strong hydrogen bonding attractions.

18. Hydrogen bonds are unique dipole-dipole attractions.
Bonds between H and O, F or N are very electronegative with H at the + end of the dipole.
The hydrogen proton is nearly exposed.
The small size of the hydrogen atom allows it to approach an electronegative atom very closely, allowing it to strongly interact with it.
The energies of hydrogen bonds vary from about 4 kJ/mol to about 25 kJ/mol.
Hydrogen bonds are much weaker than ordinary chemical bonds, but are generally stronger than other dipole-dipole force or dispersion forces.

19. Hydrogen bonds and water/ice densities.
The lower density of ice compared to liquid water can be understood in terms of hydrogen-bonding interactions.
In ice, the lower temperature slows down the motion of water molecules.
The water molecules assume an orderly, open arrangement with each water molecule forming hydrogen bonds to four other water molecules.
These hydrogen bonds create open cavities in the ice structure. The water molecules are kept further apart than in liquid water and, hence, the density of ice is less than in liquid water. Ice floats!

20. As the temperature increases, molecular motions increase and the cavities of the ice structure collapse.
The hydrogen bonding in the liquid water is more random than in ice, but it is strong enough to hold water molecules closer together giving it a more dense structure than in ice.

21. Ion-Dipole Interactions
A fourth type of force, ion-dipole interactions are an important force in solutions of ions.
The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

22. Summarizing Intermolecular Forces