Last lecture we saw how to figure out the overall polarity of a molecule from its geometry and the polarity of the bonds. We indicated that this information was important in understanding the intermolecular forces between molecules, the attractive forces that hold the particles together in liquids and solids.

We reviewed our idea of gases according to the kinetic molecular theory, and pointed out that it requires energy to convert a solid into a liquid and a liquid into a gas, energy needed to overcome these attractive forces.

There are several kinds of forces acting between particles, and they vary in their strength. We will consider the following:

· ionic forces (coulombic attraction between ions)

· ion-dipole forces (solutions of ions in polar liquids)

· dipole-dipole forces (between polar molecules)

    o hydrogen bonds (special case of dipole forces)

· dispersion forces (between neutral molecules)

Ionic forces are the strongest. The force of attraction depends on the size of the charge on the ion and the distance between them. These are called coulombic forces. Ionic compounds, such as sodium chloride, are high melting (812 ^oC), and very high temperatures are necessary to vaporize the ions (1413 ^oC). Recall that in the solid, each ion is surrounded by several ions of opposite charge.

These are the next strongest forces. You will only have these interactions in solutions, most commonly solutions of ionic substances in water. The interactions of the dipole of water with the ions must be strong enough to help break apart the lattice arrangement of the ions in solution.

We refer to this interaction as the solvation of ions by the solvent water. Notice each positive ion is surrounded by the negative end of the polar molecule, and each negative ion is surrounded by the positive end of the polar molecule.

Polar molecules can behave like little magnets, with the positive end of one molecule attracted to the negative end of the other. A variety of such arrangements are possible:

When hydrogen atoms are attached to very electronegative, very small atoms (specifically F, O, and N), the polarity of the bond is very great, and the positive end of the bond at the hydrogen atom is capable of getting very close to and interacting with the lone pairs on other F, O, or N atoms.

These bonds are considered special because the strength of attraction is quite a bit more than the attraction between dipoles. Lets look at the trend in boiling points for the hydrogen compounds of the halogens and the oxygen family:

Notice that the trend is increasing boiling point as the molecular weight of the compound gets larger, except for HF and H₂O. These both have unusually high boiling points, attributed to the hydrogen bonding that can occur between these molecules. (Remember, only hydrogens attached to F, O, or N can participate in hydrogen bonding).

In ice (solid water), the hydrogen bonding between molecules forms a tetrahedral bonding structure. This bonding gives ice a very open structure, as illustrated by this chime figure.  Melting of ice breaks some of the multiple hydrogen bonds, allowing the molecules to move relative to one another, but still many of the bonds remain. However, as they can move, they can on average get closer together, so that on melting, the liquid water becomes more dense than the solid. Hence ice floats. (If it didn't, the conditions for developing life on earth might be quite different).

So the hydrogen bonding of water explains these unusual properties of water:

· Very high melting and boiling point

· Very high specific heat

· High heat of fusion

· High heat of vaporization

· Ice less dense than liquid water

What holds non-polar molecules together? Even very non polar atoms like helium can be liquefied at a low enough temperature. The explanation for these forces is a bit more complicated.

Let's first look at helium atoms (2 electrons). Consider the particle nature of electrons

· The average distribution of electrons around each nucleus is spherically symmetrical

· The atoms are non-polar and posses no dipole moment

· The distribution of electrons around an individual atom, at a given instant in time, may not be perfectly symmetrical

· Both electrons may be on one side of the nucleus

· The atom would have an apparent dipole moment at that instant in time (i.e. a transient dipole)

· A close neighboring atom would be influenced by this apparent dipole - the electrons of the  neighboring atom would move away from the negative region of the dipole

· Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom

· This will cause the neighboring atoms to be attracted to one another

These are referred to as London dispersion forces.  They are effective only when the particles are very close to one another.  They are much weaker than the other forces, but they increase as the size of the atom increases, and as the total number of electrons in the atom or molecule increase. Therefore dispersion forces get stronger as molecular weight increases, and they are what is responsible for boiling points increasing with molecular weight. (Refer back to the comparison of HCl, HBr, and HI, which decrease in polarity as you go to higher molecular weight, but increase in boiling point).

Whether substances can mix to form solutions depends on these properties. In general, substances mix best when they have similar bonding properties. (Like dissolves like). Consider gasoline, which consists of a mixture of hydrocarbons such as octane:

The C-H bonds are not very polar, and the tetrahedral arrangement tends to cancel out what little polarity exists, so hydrocarbons are not polar. They are held together by dispersion forces. Hydrocarbons will dissolve other non-polar substances, such as grease, but will not mix with water. Whereas water can solvate polar molecules and ions, they interact with themselves more strongly than they can interact with a non-polar molecule.

Molecules with OH bonds, like ethyl alcohol and the sugar glucose, are not only polar, but they also can participate in hydrogen bonding with water, so they are soluble in water.

Any questions before we move on to the next chapter?

In lecture 2, you were asked to learn the names and symbols for all the elements listed in Table 1.1. Go back and review those, because I'm now going to ask you to expand your list of symbols and names to that of naming compounds.

Ionic compounds are named by simply naming the ions that make up the compound. So you need to learn the names of the ions. Most are pretty simple.

Cations (positive ions) are simply given the name of the element. For alkali metal, alkaline earth metal, and aluminum, that is all you need to use, because they form only one kind of ion. We pointed out that the transition metals can sometimes form more than one ion, and in that case you need to designate which one you are talking about. There are two ways of doing the, the modern Stock nomenclature, which simply puts the charge in parenthesis after the ion, but in Roman numerals. An older method uses the suffixes ous and ic to differentiate between the lower and higher charge when two are possible. So, for example:

Note that you just have to learn which charge ion each forms, and which are the ous and ic forms. Note also that the ous and ic nomenclature is used with the Latin names of the elements.

Anions (negative ions) add the ending ide to the stem of the element name. Chlorine becomes chloride, bromine becomes bromide, etc. Sometimes the stem is changed slightly—oxygen becomes oxide.

The important monatomic ions for you to remember are given in Table 6.1, page 142. Learn them.

Many compounds contain both ionic and covalent bonds. Collections of atoms held together with covalent bonds, yet carrying a net charge, are called polyatomic ions.

Polyatomic cations are very few. You need concern yourself only with ammonium (NH₄⁺) and hydronium (H₃O⁺).

There are three polyatomic anions with common names that you should learn: hydroxide (OH^-), cyanide (CN^-) and acetate (CH₃CO₂^-).

Note you can draw Lewis dot structures for these ions just like for covalent compounds. If the ion is positive, you need to deduct one or more electrons in calculating the number of valence electrons in the structure. If the ion is negative, you need to add one or more electrons to the number for the structure.

The rest of the polyatomic anions that concern us are oxyanions, that is they contain one or more oxygens surrounding a central non-metal atom.   We will discuss these in the next lecture.

Special Note:

Hour Test 2 will cover Chapters 4 and 5 and pages 141-148 of Chapter 6 (through nomenclature).