Chapter 11Chapter 11

Extension of K-M Theory to liquids & solids:

1. Still composed of particles but they are closer together in liquids & solids

2. Particles are still moving & rate is dependent on T, but not proportional

3. In liquid - particles move over, under & around each other - held by attractions

4. In solid - only vibrations - basically fixed in relative position

5. Liquids still diffuse only slower.

We have previously discussed what I call primary chemical bonds (bonding between atoms within molecules or ions). Now we will discuss what I call secondary chemical bonds (attractions between different molecules and/or ions). These are also called intermolecular attractions or more commonly Van der Waals attractions. There are several types:

1. Between 2 non-polar molecules - called London Forces

2. Between two polar molecules - called dipole-dipole attraction

3. Hydrogen bonding

4. Between a polar and a non-polar molecule

London Forces – Also called Dispersion Forces. Caused by attraction of outer electrons of one molecule for the + charged center of a nearby molecule. This is extremely temporary as the molecules are constantly moving and this attraction is constantly changing. For the instant of attraction, however, the 2 molecules get distorted and thus each temporarily is polarized ( a + and – end for each molecule created by the distortion of the electron cloud). Each of these will then induce distortions in other nearby molecules, by reason of the polarization.

These are generally very weak - about 4 kJ /mole or less

Evidence for the existence of London Forces:

Assume non-polar substance. Should have no attractions (no charged ends). But we know there must be attractions because all such substances do have liquid and solid phases, where intermolecular attractions are significant. Best explanation is these are temporary, but always present somewhere in the sample, called polarizations.

Two Major Effects of London Forces on M.P. and B.P:

Review process of melting & boiling, vis-à-vis overcoming attractive forces

Therefore greater London Forces, higher M.P.’s & B.P.’s, all other factors being equal.

1. All other factors being equal, bigger atoms in a molecule, higher London forces, because bigger atoms are easier to polarize. (Ex. F2 [-188], Cl2[-34.7], Br2[58.0], I2[183]

2. All other factors being equal, more atoms in a molecule, larger London Forces, because molecule is larger and therefore easier to polarize. (Ex. CH4, C2H6, C3H8 and

C4H10 -162, -84.5, -42, 0)

In actuality both of these rules result from same concept - all other factors being equal, the larger the molecule, the greater the London Forces. Be careful to compare similar compounds.

Hydrogen Bond – Essentially this is an extreme case of dipole-dipole attraction. It occurs whenever H is covalently bonded to N, O or F only. Any time that occurs the effects of H- bonding will be seen - Higher B.P.’s and M.P.’s than expected

Illustration - H2Te, H2Se, H2S, and H2O --> B.P.’s are -

2, -41.5, -60.7 and +100. For more examples see Figure 11.7 on page 423.

Hydrogen bonds are so strong that they can sometimes be almost as strong as a primary covalent bond. This is the case in H2O. In the liquid state, H2O exists as a group of

several molecules traveling together and sticking together because of the H-bond. When H2O freezes, each group

gets locked into a position in the new crystal that forms, but since this is an uneven group, this results in empty spaces in the crystal (you can see these empty spaces if you look closely at an ice cube. The end result of all this is that ice is less dense than liquid H2O ( one of the few substances

where the solid is less dense than the liquid). This is very good, because ice floats. If not for this, in winter cold climates, all life in lakes, rivers and oceans would die, because they would freeze from the bottom up, rather than from the top down. They would totally freeze.

True solids exist as crystals, which are regularly shaped solid forms. There are several common shapes. In all cases the particles making up the crystal (atoms, ions or molecules) try to get as close as possible to each other. There are various ways to get this done, but we won’t worry about them.

Amorphous solids don’t form regularly shaped crystals and thus have slightly different properties. The most pronounced is that they don’t have exact melting points, but they gradually get softer and softer until they appear to be liquid. Glass is the most common amorphous solid.