Brønsted Acids and Bases
|Brønsted Acid-Base Theory||Brønsted Acids and Bases in Nonaqueous Solutions|
|Typical Brønsted Acids and Their Conjugate Bases|
For more than 300 years, substances that behaved like vinegar have been classified as acids, while those that have properties like the ash from a wood fire have been called alkalies or bases. The name "acid" comes from the Latin acidus, which means "sour," and refers to the sharp odor and sour taste of many acids. Vinegar tastes sour because it is a dilute solution of acetic acid in water; lemon juice is sour because it contains citric acid; milk turns sour when it spoils because of the formation of lactic acid; and the sour odor of rotten meat can be attributed to carboxylic acids such as butyric acid formed when fat spoils.
Today, when chemists use the words "acid" or "base" they refer to a model developed independently by Brønsted, Lowry, and Bjerrum. Since the most explicit statement of this theory was contained in the writings of Brønsted, it is most commonly known as the "Brønsted acid-base" theory.
Brønsted Acid-Base Theory
Brønsted argued that all acid-base reactions involve the transfer of an H+ ion, or proton. Water reacts with itself, for example, by transferring an H+ ion from one molecule to another to form an H3O+ ion and an OH- ion.
According to this theory, an acid is a "proton donor" and a base is a "proton acceptor."
Acids are often divided into categories such as "strong" and "weak." One measure of the strength of an acid is the acid-dissociation equilibrium constant, Ka, for that acid.
When Ka is relatively large, we have a strong acid.
HCl: Ka = 1 x 103
When it is small, we have a weak acid.
CH3CO2H: Ka = 1.8 x 10-5
When it is very small, we have a very weak acid.
H2O: Ka = 1.8 x 10-16
In 1909, S. P. L. Sørenson suggested that the enormous range of concentrations of the H3O+ and OH- ions in aqueous solutions could be compressed into a more manageable set of data by taking advantage of logarithmic mathematics and calculating the pH or pOH of the solution.
pH = - log [H3O+]
pOH = - log [OH-]
The "p" in pH and pOH is an operator that indicates that the negative of the logarithm should be calculated for any quantity to which it is attached. Thus, pKa is the negative of the logarithm of the acid-dissociation equilibrium constant.
pKa = - log Ka
The only disadvantage of using pKa as a measure of the relative strengths of acids is the fact that large numbers now describe weak acids, and small (negative) numbers describe strong acids.
|HCl:||pKa = -3|
|CH3CO2H:||pKa = 4.7|
|H2O:||pKa = 15.7|
An important features of the Brønsted theory is the relationship it creates between acids and bases. Every Brønsted acid has a conjugate base, and vice versa.
Just as the magnitude of Ka is a measure of the strength of an acid, the value of Kb reflects the strength of its conjugate base. Consider what happens when we multiply the Ka expression for a generic acid (HA) by the Kb expression for its conjugate base (A-).
If we now replace each term in this equation by the appropriate equilibrium constant, we get the following equation.
KaKb = Kw = 1 x 10-14
Because the product of Ka times Kb is a relatively small number, either the acid or its conjugate base can be "strong." But if one is strong, the other must be weak. Thus, a strong acid must have a weak conjugate base.
A strong base, on the other hand, must have a weak conjugate acid.
Brønsted Acids and Bases in Nonaqueous Solutions
Water has a limiting effect on the strength of acids and bases. All strong acids behave the same in water -- 1 M solutions of the strong acids all behave as 1 M solutions of the H3O+ ion -- and very weak acids cannot act as acids in water. Acid-base reactions don't have to occur in water, however. When other solvents are used, the full range of acid-base strength shown in the following table can be observed.
Typical Brønsted Acids and Their Conjugate Bases
|HI||3 x 109||-9.5||I-||3 x 10-24||23.5|
|HCl||1 x 106||-6||Cl-||1 x 10-20||20|
|H2SO4||1 x 103||-3||HSO4-||1 x 10-17||17|
|H3O+||55||-1.7||H2O||1.8 x 10-16||15.7|
|HNO3||28||-1.4||NO3-||3.6 x 10-16||15.4|
|H3PO4||7.1 x 10-3||2.1||H2PO4-||1.4 x 10-12||11.9|
|CH3CO2H||1.8 x 10-5||4.7||CH3CO2-||5.6 x 10-10||9.3|
|H2S||1.0 x 10-7||7.0||HS-||1 x 10-7||7.0|
|H2O||1.8 x 10-16||15.7||OH-||55||-1.7|
|CH3OH||1 x 10-18||18||CH3O-||1 x 104||-4|
|HCCH||1 x 10-25||25||HCC-||1 x 1011||-11|
|NH3||1 x 10-33||33||NH2-||1 x 1019||-19|
|H2||1 x 10-35||35||H-||1 x 1021||-21|
|CH2=CH2||1 x 10-44||44||CH2=CH-||1 x 1030||-30|
|CH4||1 x 10-49||49||CH3-||1 x 1035||-35|
The strongest acids are in the upper-left corner of this table; the strongest bases in the bottom-right corner. Each base is strong enough to deprotonate the acid in any line above it. The hydride ion (H-), for example, can convert an alcohol into its conjugate base
and the amide (NH2-) ion can deprotonate an alkyne.