Inorganic Acids and Bases - pKa Values

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

pK_a 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 compounds,  pKa 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 pK_as).
NB!  For bases, the dissosiation constant is given in terms of the pK_a value of their conjugated acid, BH⁺. These values are marked with an *.

Inorganic Acids and Bases - pKa Values

Acid name	Formula	pKa1	pKa2	pKa3	pKa4
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, K_a, 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⁺)

K_a = (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           K_a = (A^-)*(H⁺) / (HA)

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

pK_a = -log₁₀ K_a

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

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

Strong acids has pK_a 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. pK_a values for strong acids can, however, be estimated by theoretical means.

After rearranging the expression defining K_a, and putting pH = -log₁₀(H⁺), one obtains

pH = pK_a + (A^-) / (HA)  and further

pH - pK_a = log(A^-) / (HA)

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

Polyprotic acids are acids that can lose more than one proton. Then we have more than one dissiciation constant; K_a1, K_a2, etc..  and similar pK_a1, pK_a2, etc.

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

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

BH⁺  =  B  + H⁺

The  pK_b for a base may be calculated from the pK_a value of its conjugate acid:

pK_w  =  pK_a  + pK_b

At 25°C the pK_w is 14 and

pK_b = 14 - pK_a

The figure below shows the fractions of the different acid forms (H₂A, HA^- and A^2-for twoprotic acids) of di- and triprotic acids in aqueous solutions at varying pH: