MORISSON   Organic Chemistry

MORISSON Organic Chemistry

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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, boils 
100 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: 
H H 
I i 
H-C-H H-0-H H-C-0-H 
I I 
H H 
Methane Water Methanol 
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. 
. . 
W e m 1.8 Which of the following organic compounds would you predict to 
be associated liquids? Draw structures to show the hydrogen bonding you would expect. 
(a) CH30CH3; (b) CH3F; (c) CH3CI; (d) CH3NH,; (e) (CH3),NH; (f) (CH3),N. 
1.21 Solubility 
When a solid or liquid didsolves, the structural units-ions or molecules- 
become 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 solvents 
are 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 ion- 
dipole 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 (see Fig. 1.22). 
Figure 1.22 Ion-dipole interactions: a solvated cation and anion. 
To dissolve ionic compounds asolvent must also have a high dielectric constant, 
thaf is, h b e high insulating properties to lower the attraction between oppositely 
charged ions once they are solvated. 
Water owes its shperiority as a solvent for ionic substances not only to its 
polarity and its high dielectric constant but to another factor as well: it contains 
the -OH group and thus can form hydrogen bonds. Water solvates both cations 
and 'ixhions: cations, at its negative pole (its unshared electrons, essentially); 
anions, through hydrogen bonding. 
Now let us turn to the dissolution of non-ionic solutes. 
The solubility characteristics of non-ionic compounds are determined chiefly 
by their polarity. Non-polar or weakly polar compounds dissolve in non-polar or 
weakly polar solvents; highly polar compounds dissolve in highly polar solvents. 
"Like dissolves like" is an extremeiy useful rule of thumb. Methane dissolves in 
carbon tetrachloride because the forces holding methane molecules to each other 
and carbon tetrachloride molecules to each other-van der Waals interactions- 
are replaced by very similar forces holding methane molecules to carbon tetra- 
chloride molecules. 
Neither methane nor carbon tetrachloride is readily soluble in water. The 
highly polar water molecules are held to each other by very strong dipole-dipole 
interactions-hydrogen bonds; there could be only very weak attractive forces 
between water molecules on the one hand and the non-polar methane or carbon 
tetrachloride molecules on the other. 
In contrast, the highly polar organic compound methanol, CH,OH, is quite . 
soluble in water. Hydrogen bonds between water and methanol molecules can 
readily replace the very similar hydrogen bonds between different methanol 
molecules and different water molecules. 
An understanding of the nature of solutions is fundamental to an understan- 
ding of organic chemistry. Most organic reactions are camed out in solution and, 
ir is W a g increasingly clear, the solvent does much more than simply bring 
different molecules together so that they can react with each other. The solvent is 
iradFcd m the reactions that take place in it: just how much it is involved, and in 
what ways, is only now being realized. In Chapter 7, when we know a little more 
about organic reactions and how they take place, we shall return to this subject- 
which we have barely touched upon here-and examine in detail the role played 
by the solvent. 
1.22 Acids and bases 
Turning from physical to chemical properties, let us review briefly one familiar 
topic that is fundamental to the understanding of organic chemistry: acidity and 
The terms acid and base have been defined in a number of ways, each definition 
corresponding to a particular way of looking at the properties of acidity and 
basicity. We shall find it useful to look at acids and bases from two of these 
viewpoints; the one we select will depend upon the problem at hand. 
According to the Lowry-Brensted definition, an acid is a substance t b t gives 
up a proton, and a base is a substance that accepts a proton. When sulfuric acid 
dissolves in water, the acid H2S04 gives up a proton (hydrogen nucleus) to the 
base H 2 0 to form the new acid H,O+ and the new base HS04-. When hydrogen 
chloride reacts with ammonia, the acid HCl gives up a proton to the base NH, to 
form the new acid NH4+ and the new base C1- 
H2S04 + H20 H30+ + HS04- 
Stronger Stronger Weaker Weaker 
acid base acid 5ase 
HCI + NH3 S NH4+ + C1- 
Stronger Stronger Weaker Weaker 
acid base acid base 
According to the Lowry-Brmsted definition, the strength of an acid depends 
upon its tendency to give up a proton, and the strength of a base depends upon its 
tendency to accept a proton. Sulfuric acid and hydrogen chloride are strong acids 
since they tend to give up a proton very readily; conversely, bisulfate ion, 
HS04-, and chloride ion must necessarily be weak bases since they have little 
tendency to hold on to protons. In each of the reactions just described, the 
equilibrium favors the formation of the weaker acid and the weaker base. 
If aqueous H2S04 is mixed with