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Boiling

Although the particles in a liquid are arranged less regularly and are freer to

move about than in a crystal, each particle is attracted by a number of other

particles. Boiling involves the breaking away from the liquid of individual molecules

or pairs of oppositely charged ions (see Figs. 1.20 and 1.21). This occurs

when a temperature is reached at which the thermal energy of the particles is

great enough to overcome the cohesive forces that hold them in the liquid. held strongly by a number of oppositely charged ions. Again there is nothing

we could properly call a molecule. A great deal of energy is required for a pair of

oppositely charged ions to break away from the liquid; boiling occurs only at a

very high temperature. The boiling point of sodium chloride, for example, is

141 3 "C. In the gaseousstate we have an ionpair, which can be considered a sodium

chloride molecule.

In the liquid state the unit of a non-ionic compound is again the molecule.

The weak intermolecular forces here-dipole-dipole interactions and van derWaals forces-are more readily overcome than the strong interionic forces of ionic

compounds, and boiling occurs at a very much lower temperature. Non-polar

methane boils at - 161.5 "C, and even polar hydrogen chloride boils at only

- 85 "C.

Liquids whose molecules are held together by hydrogen bonds are called

~cwciatedli quids. Breaking these hydrogen bonds takes considerable energy, and

(;. 3n associated liquid has a boiling point that is abnormally high for a compound

7 9 c molecular weight and dipole moment. Hydrogen fluoride, for example, boils100 degrees higher than the heavier, non-associated hydrogen chloride; water boils

160ldegrees higher than hydrogen sulfide.

There are organic compounds, too, that contain hydrogen bonded to oxygen

or nitrogen, and here, too, hydrogen bonding occurs. Let us take, for example,

methane and replace one of its hydrogens with a hydroxyl group, --OH. The

resulting compound, CH,OH, is methanol, the smallest member of the alcohol

family. Structurally, it resembles not only methane. but also water:

Like water, it is an associated liquid with a boiling point "abnormally" high for a

compound of its size and polarity.

The bigger the molecules, the stronger the van der Waals forces. Other things

being equal-polarity, hydrogen bonding-boiling point rises with increasing

molecular size. Boiling points of organic compounds range upward from that of

tiny, non-polar methane, but we seldom encounter boiling points much above

350 "C;at higher temperatures, covalent bohds within the molecules start to break,

and decomposition competes with boiling. It is to lower the boiling point and thus

minimize decomposition that distillation of organic compounds is often carried

out under reduced pressure.

Solubility

When a solid or liquid didsolves, the structural units-ions or moleculesbecome

separated from each other, and the spaces in between become occupied

by solvent molecules. In dissolution, as in melting and boiling, energy must be

supplied to overcome the interionic or intermolecular forces. Where does the

necessary energy come from? The energy required to break the bonds between

solute particles is supplied by the formation of bonds between the solute particles

and the solvent molecules: the old attractive forces are replaced by new ones.

Now, what are these bonds that are formed between solute and solvent? Let

us consider first the case of ionic solutes.

A great deal of energy is necessary to overcome the powerful electrostatic

forces holding together an ionic lattice. Only water or other highly polar solventsare able to dissolve ionic compounds appreciably. What kinds of bonds are formed

between ions and a polar solvent? By definition, a polar molecule has a positive

end and a xlegative end. Consequently, there is electrostatic attraction between a

positive ion and thenegative end of the solvent molecule, and between a negative

ion and the positive end of the solvent molecule. These attractions are called iondipole

bonds. Each ion-dipole borld is relatively weak, but in the aggregate they

supply enough energy to overcome the interionic forces in the crystal. In solution

each ion is surrounded by a cluster of solvent molecules, and is said to be solvated;

if the solvent happens to be water, the ion is said to be hydrated. In solution, as in

the solid and liquid states, the unit of a substance like sodium chloride is the ion,

although in this case it is a solvated ion