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oxyacid, any oxygen-containing acid. Most covalent nonmetallic oxides react with water to form acidic oxides;
that is, they react with water to form oxyacids that yield hydronium ions (H3O+) in solution.
There are some exceptions, such as carbon monoxide, CO, nitrous oxide, N2O, and nitric oxide, NO.
The strength of an oxyacid is defined by the extent to which it dissociates in water (i.e., its ability to form H+ ions).
In general, the relative strength of oxyacids can be predicted on the basis of
the electronegativity and oxidation number of the central nonmetal atom.
The acid strength increases as the electronegativity of the central atom increases.
For example, because the electronegativity of chlorine (Cl) is greater than that of sulfur (S),
which is in turn greater than that of phosphorus (P), it can be predicted that perchloric acid, HClO4,
is a stronger acid than sulfuric acid, H2SO4, which should be a stronger acid than phosphoric acid, H3PO4.
For a given nonmetal central atom, the acid strength increases as the oxidation number of the central atom increases.
For example, nitric acid, HNO3, in which the nitrogen (N) atom
has an oxidation number of +5, is a stronger acid than nitrous acid, HNO2,
where the nitrogen oxidation state is +3. In the same manner, sulfuric acid, H2SO4,
with sulfur in its +6 oxidation state, is a stronger acid than sulfurous acid, H2SO3, where a +4 oxidation number of sulfur exists.
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Acid Strength and Bond Strength
Binary acids are certain molecular compounds in which hydrogen is combined with a second nonmetallic element; these acids include HF, HCl, HBr, and HI.
HCl, HBr, and HI are all strong acids, whereas HF is a weak acid. The acid strength increases as the experimental pKa values decrease in the following order:
HF (pKa = 3.1) < HCl (pKa = -6.0) < HBr (pKa = -9.0) < HI (pKa = -9.5).
Why is HF a weak acid, when the rest of the hydrohalic acids are strong?
One might correctly assume that fluorine is very electronegative, so the H-F bond is highly polar and we can expect HF to dissociate readily in solution; this reasoning is not wrong, but the electronegativity argument is trumped by considerations of ionic size. Recall the periodic trend that ionic size increases as we move down the periodic table. Because fluorine is at the top of the halogens, the F– ion is the smallest halide; therefore, its electrons are concentrated around its nucleus, and as a result, the H-F bond is relatively short. Shorter bonds are more stable, and thus the H-F bond is more difficult to break.
Once we move down to chlorine, however, the trend changes. Chlorine is larger and has more electrons, and therefore the H-Cl bond is longer and weaker. In the presence of water, the electrostatic attractions between water’s partially negative oxygen and the partially positive hydrogen on H-Cl is strong enough to break the H-Cl bond, and the ions dissociate in solution.
The same reasoning applies for both HBr and HI. These acids are even stronger than HCl because the Br– and I– ions are even larger. As such, the H-Br and H-I bonds are even weaker, and these compounds also readily dissociate in solution.