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Inorganic Acids and Bases - pKa Values

Values for the negative logarithm of the acid dissociation constant, pKa, of inorganic acids and bases, as well as hydrated metal ions.

Definitions of the acid dissociation constant and pKa are given below the table.

pKa values given in the table are measuered at 25°C, unless other temperature(°C) is indicated with superscript at the pKa value. 

See also Acid-base properties of aqueous solutions of salts with ions from both acids and bases, Buffer solutions,  pKa of amines, diamines and cyclic organic nitrogen compoundspKa of phenols, alcohols and carboxylic acids and acid and base pH indicators

Below the table, figures showing the fractions of the different acid forms in aqueous solutions at varying pH are given for some common acids (values calculated from the tabulated pKas).
NB!  For bases, the dissosiation constant is given in terms of the pKa value of their conjugated acid, BH+. These values are marked with an *.

Inorganic Acids and Bases - pKa Values
Acid nameFormulapKa1pKa2pKa3pKa4
Monoprotic acids
Ammonia NH3 9.24*
Arsenious acid H3AsO3 9.29
Boric acid H3BO3 9.24
Chlorous acid HClO2 1.94
Cyanic acid HOCN 3.46
Hydrazoic acid HN3 4.72
Hydrazine N2H4 8.1*
Hydrobromic acid HBr -9
Hydrochloric acid HCl -7
Hydrocyanic acid HCN 9.21
Hydrofluoric acid HF 3.17
Hydroiodic acid HI -10
Hydrogen peroxide H2O2 11.65
Hydroxylamine H2NOH 5.94
Hypobromous acid HOBr 8.62
Hypochlorous acid HOCl 7.54
Hypoiodous acid HOI 10.64
Hypophosphorous acid H3PO2 1.23
Iodic acid HIO3 0.78
Nitrous acid HNO2 3.25
Nitric acid HNO3 -1.37
Perchloric acid HClO4 -1.620
Periodic acid HIO4 1.64
Sulfamic acid H2NSO3H 1.05
Tetrafluoroboric acid HBF4 0.5
Thiocyanic acid HSCN -1.8
Water H2O 13.995
Diprotic acids
Boric acid H3BO3 9.2720 1420
Carbonic acid H2CO3 6.35 10.33
Chromic acid H2CrO4 -0.98 6.49
Germanic acid H2GeO3 9.01 12.3
Hydrogen selenide H2Se 3.89 11.0
Hydrogen sulfide H2S 7.02 19
Hydrogen telluride H2Te 2.618 11
Phosphorous acid H3PO3 1.320 6.720
Selenous acid H2SeO3 2.62 8.32
Selenic acid H2SeO4 <0 1.7
Silicic acid H2SiO3 9.91 11.81
Sulfurous acid H2SO3 1.85 7.20
Sulfuric acid H2SO4 -3 1.99
Tellur(VI)ic acid H6TeO6 7.6818 1118
Tellurous acid H2TeO3 6.27 8.43
Triprotic acids
Arsenic acid H3AsO4 2.26 6.76 11.29
Phosphoric acid H3PO4 2.16 7.21 12.32
Tetraprotic acids
Pyrophosphoric acid H4P2O7 0.91 2.10 6.70 9.32
Orthosilicic acid H4SiO4 9.930 11.830 1230 1230
Hydrated metal ions
Aluminium (III) ion Al3+ 4.85
Barium (II) ion Ba2+ 13.4
Beryllium(II) ion Be2+ 5.7
Calcium(II) ion Ca2+ 12.6
Chromium (III) ion Cr3+ 3.95
Copper (II) ion Ba2+ 7.34
Iron (III) ion Fe3+ 2.17
Lead (II) ion Ba2+ 7.8
Lithium (I) ion Li+ 13.8
Magnesium (II) ion Mg2+ 11.4
Manganese (II) ion Mn2+ 10.59
Nickel (II) ion Ni2+ 9.86
Sodium (I) ion Na2+ 14.8
Scandium (III) ion Sc3+ 4.61
Strontium (II) ion Sr2+ 13.2
Uranium (IV) ion U4+ 0.68
Vanadium (III) ion V3+ 2.9
Zinc (II) ion Zn2+ 8.96

An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation of acid–base reactions. In aqueous solution, the equilibrium of acid dissociation can be written symbolically as:

HA + H2O   =  A- +H3O+

where HA is an acid that dissociates into A-, (known as the conjugate base of the acid) and a hydrogen ion which combines with a water molecule to make a hydronium ion.

The chemical species HA, A- and H3O+ are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A-] and [H3O+]

Ka = [A-]*[H3O+] / [HA]*[H2O]

In all, but the most concentrated, aqueous solutions of an acid the concentration of water can be taken as constant and can be ignored. The definition can then be written more simply

HA = A- + H+        and           Ka = [A-]*[H+] / [HA]

This is the definition in common usage. For many practical purposes it is more convenient to discuss the logarithmic constant, pKa

pKa = -log10 Ka

The larger the value of pKa, the smaller the extent of dissociation at any given pH - that is, the weaker the acid.

A weak acid has a pKa value in the approximate range -2 to 12 in water.

Strong acids has pKa values of less than about -2; the dissociation of a strong acid is effectively complete such that concentration of the undissociated acid is too small to be measured. pKa values for strong acids can, however, be estimated by theoretical means.

After rearranging the expression defining Ka, and putting pH = -log10[H+], one obtains

pH = pKa + [A-] / [HA]  and further

pH - pKa = log [A-] / [HA]

Then, a solution with 50% dissociation has pH equal to the pKa of the acid.

Polyprotic acids are acids that can lose more than one proton. Then we have more than one dissiciation constant; Ka1, Ka2, etc..  and similar pKa1, pKa2, etc.

All data given in the figures apply to dilute aqueous solutions at ambient temperature.

For bases, the pka value is given for the conjugate bases BH+ and BH22+.

BH+  =  B  + H+

The  pKb for a base may be calculated from the pKa value of its conjugate acid:

pKw  =  pKa  + pKb

At 25°C the pKw is 14 and

pKb = 14 - pKa

The figure below shows the fractions of the different acid forms (H2A, HA- and A2- for twoprotic acids) of di- and triprotic acids in aqueous solutions at varying pH:

Fractions of acid ions as function of pH

Related Topics

  • Material Properties

    Properties of gases, fluids and solids. Densities, specific heats, viscosities and more.

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