Molecular Structure and the Strength of Acids
The strength of an acid depends on a number of factors, such as the properties of the solvent, the temperature, and the molecular structure of the acid. When we compare the strengths of two acids, we can eliminate some variables by considering their properties in the same solvent and at the same temperature and temperature. That way we can focus on the structure of the acid.
Binary Hydrides
Lets consider a certain acid HA. The strength of the acid is measured by its tendency to ionize:
HA ---> H+ + A-
Two factors influence the extent to which the acid undergoes ionization. One is the strength of the H-A bond, the stronger the bond, the more difficult it is for the HA molecule to break up and hence the weaker the acid. The other factor is the polarity of the H-A bond. The difference in electronegativities between H and A results in a polar bond. If the bond is highly polarized, that is if there is a large accumulation of positive and negative charges on the H and A atoms, HA will tend to break up into H+ and A- ions. So a high degree of polarity characterizes a stronger acid.
The halogens form a series of binary acids called the hydrohalic acids. The strengths of the hydrohalic acids increase in the following order:
HF<<HCl<HBr<HI
HF has the highest bond dissociation energy of the four hydrogen halides. Since it takes 568.2 kJ to break the H-F bond, HF is a weak acid. At the other extreme in the series, HI has the lowest bond energy (298.3 kJ), so HI is the strongest acid of the group. In this series of acids the polarity of the bond actually decreases from HF to HI. This property should enhance the acidity of HF relative to the other acids in the series, but its magnitude is not great enough to overcome the trend in bond dissociation energies.
In any vertical column (Group) of nonmetallic elements, there is a tendency toward increasing acidity of the hydride with increasing atomic number (as you go down the group). For example, among the group VIA elements the acid strength increases in the order:
H2O< H2S<H2Se<H2Te
This order arises primarily because the bond energies steadily decrease in this series as the central atom grows larger and the overlaps of atomic orbitals grow smaller, just as in the case of the hydrides of the halogens above.
Oxyacids
Many common acids contain one or more O-H bonds. Sulfuric acid, for example, contains two such bonds:
Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom are called oxyacids. The OH group is also present in bases. So what factors determine whether an OH group will behave as a base or as an acid? Let's consider an OH group bound to some atom X, which might in turn have other groups attached to it:
(Z)3X -O-H
At one extreme X might be a metal, such as Na, K, or Ca. Because of their low electronegativities, the pair of electrons shared between X and O is completely transferred to oxygen, and an ionic compound involving OH- is formed. Such compounds are therefore sources of OH- ions and behave as bases.
When X is a nonmetal, the bond to O is covalent in character, and the substance does not readily lose OH- . Instead, these compounds are either acidic or neutral. As a general rule, as the electronegativity of X increases, so will the acidity of the substance. This happens for two reasons: First, the O-H becomes more polar, thereby favoring loss of H+. Second, because the conjugate base is usually an anion, its stability generally increases as the electronegativity of X increases.
Many oxyacids, H2SO4 for example, contain additional oxygen atoms bonded to the central atom X. The additional electronegative oxygen atoms pull electron density from the O-H bond, further increasing its polarity. Increasing the number of oxygen atoms also helps stabilize the conjugate base by increasing its ability to spread out its negative charge.
We can summarize these ideas as two simple rules that relate the acid strength of oxyacids to the electronegativity of X and to the number of groups attached to X: