James Richard Fromm
The chemical activities aor their approximations, concentrations c, of the species which appear in the ionization constants of acids and bases are measurable and as a consequence numeric values of these constants can be determined experimentally. Tables of a few selected values of molar ionization constants of aqueous acids are included below for convenience in the discussion of this section. These are small subsets of the larger tables of acid ionization constants for monoprotic acids and polyprotic acids given elsewhere. These measured values can be calculated from tabulated thermodynamic values (which are themselves derived from experimental measurements), and vice versa, as is shown in other sections
|Acetic Acid||CH3COOH||1.75 x 10-5||4.756|
|Ammonium Ion||NH4+||5.60 x 10-10||9.252|
|Chloroacetic Acid||CH2ClCOOH||1.36 x 10-3||2.866|
|Hydrochloric Acid||HCl||Greater than 1||Negative|
|Hydroxylammonium Ion||HONH3+||1.12 x 10-6||5.951|
|Hydrogen Sulfide||H2S||K1: 1.02 x 10-7||6.992|
|Hydrogen Sulfide Ion||HS-||K2: 1.22 x 10-13||12.915|
The inherent or intrinsic strength of an aqueous acid (or base) is its ability to remove a proton from (or donate a proton to) the solvent water or other ions and molecules in aqueous solutions. For quantitative comparisons between different aqueous acids or bases, this ability is compared with the ability of the solvent water itself. The reaction used is the reaction which corresponds to the ionization equilibrium whose equilibrium constant is called the ionization constant. In other words, strengths of acids and bases are expressed quantitatively in terms of the values of their ionization constants.
Aqueous ionization constants are quantitative measures of the tendency of the acid or base to either donate a proton, written as the hydronium ion H3O+ or often simply as H+, or accept a proton from water. The greater the value of the equilibrium constant, the greater the percentage of the acid or base that will be in ionized form. As a generalization, we can use the value of the ionization constant equal to 0.1 as the point of distinction between a strong acid and a weak acid. Thus a strong acid is one for which the value of the acid ionization constant Ka is large (greater than 0.1) and a weak acid is one for which the value of the acid ionization constant Ka is small (much less than 0.1). Likewise, a strong base is one for which the value of the base ionization constant Kb is large (greater than 0.1) and a weak base is one for which the value of the base ionization constant Kb is small (much less than 0.1).
There are only a few common strong acids: hydrochloric acid, HCl; nitric acid, HNO3; perchloric acid, HClO4; and sulfuric acid, H2SO4. In the case of sulfuric acid, only the ionization of the proton from H2SO4 to give the hydrogen sulfate ion HSO4- is strong; the ionization of HSO4- to give the sulfate ion SO42- is not strong, but weak.
Common strong bases include NaOH and KOH. On reaction with water, CaO gives the strong base Ca(OH)2 and for that reason CaO is considered a strong base also, as are the oxides of sodium and potassium.
Using the values of the ionization constant as quantitative measures of acid strength is equivalent to the qualitative statement that a strong acid is an acid for which loss of the proton to water is essentially complete, while a weak acid is an acid for which loss of the proton is noticeably incomplete. Likewise, a strong base is a base for which acquisition of a proton from water is essentially complete while a weak base is a base for which acquisition of a proton is noticeably incomplete.
These distinctions can be used to illustrate some of the terminology of acid-base reactions. The ionization of hydrogen cyanide as an acid in water is the reaction
HCN(aq) + H2O H3O+ + CN-
ACID 1 + BASE 2 ACID 2 + BASE 1
The acid-base pair marked 1 is referred to as a conjugate pair, that is, an acid-base pair differing by one proton. Hydrogen cyanide is said to be the conjugate acid of cyanide ion, and cyanide ion is said to be the conjugate base of hydrogen cyanide. Likewise, water is the conjugate base of the hydrogen ion H3O+, and the hydrogen ion is the conjugate acid of water. The same terminology is illustrated by ammonia:
NH3(aq) + H2O NH4+ + OH-
is also a reaction of the form
BASE 1 + ACID 2 ACID 1 + BASE 2.
In this reaction ammonia and ammonium ion form a conjugate pair, as do water and hydronium ion.
Water can act either as an acid or a base. The ionization of pure water can therefore be treated as the acid-base equilibrium
H2O + H2O H3O+(aq) + OH-(aq),
for which the equilibrium constant may be written:
Ka(H2O) = a(H3O+)a(OH-)/a2(H2O) = Kb(H2O) = Kw
When the activity of water is taken as one because a dilute aqueous solution is essentially pure water, and activities are approximated by concentrations, then Kw = [H3O+][OH-].
The constant Kw, the autoprotolysis constant or the autoionization constant of water, has the value 1.014 x 10-14 at 25oC when molar units are used. It usually suffices to use the approximate numerical value of 1.0 x 10-14 mol2/L2.
The autoprotolysis of water sets limits on the acid strength and base strength which can exist in aqueous solution. Any base significantly stronger than hydroxide ion, such as methoxide ion, will react quantitatively to deprotonate water producing hydroxide ion. Likewise, any acid significantly stronger than hydronium ion, such as hydrochloric acid, will react quantitatively to protonate water, producing hydronium ion. For this reason the range of molar ionization constants observed in aqueous solutions is from about one to about 1 x 10-14 in molar units.
In any aqueous solution, the value of Ka and of Kb for an acid-base conjugate pair are related through Kwof the solvent, water, because water participates in the ionization reactions. For any acid-base conjugate pair it is true that KaKb = Kw. For hydrocyanic acid, or hydrogen cyanide, as an example:
[H3O+][CN-]/[HCN] x [OH-][HCN]/[CN-] = [H3O+][OH-]
Ka(HCN) x Kb(CN-) = Kw
As a consequence, values of the aqueous ionization constants of bases can always be calculated whenever desired, simply by dividing the aqueous ionization constant of their conjugate acid by the value of Kw. Separate tables of ionization constants of aqueous bases are rarely used for this reason.
A polyprotic acid is an acid which can lose more than one proton. The protons can be, and usually are, considered as ionizing separately and each can be treated using a separate acid ionization constant. These successive acid ionization constants are designated as Ka1 and Ka2, or more often simply as K1 and K2, for the first and second proton respectively.
The successive ionization constants for most polyprotic acids decrease by several orders of magnitude because the loss of one proton leaves the acid more negatively charged. Loss of the next positively charged proton is then more difficult due to the increased electrostatic attraction. For molecules such as EDTA (ethylenediaminetetraacetic acid) and succinic acid, the values of K1 and K2 are similar because these acid ionizations occur at opposite ends of large molecules and do not influence each other greatly. In EDTA, however, the third and fourth acid ionizations must occur near the previous acid ionizations and the decrease is observed for the values of K3 and K4.