Exp. Methyl orange (very little) is added to 10 c.c. of a 0.1 molar solution of phosphoric acid and 10 c.c. of 0.1 molar sodium hydroxide solution is added to the mixture; the color will be found to change from the acid to the neutral tint just as the last drop or two of the alkali are added. Phenolphthaleïn is then added to the mixture and 10 c.c. more of the 0.1 molar sodium hydroxide solution are required to change the color of the mixture to a pronounced pink (alkaline) tint.
Even sulphuric acid, although its two hydrogen atoms are ionized very easily, making sulphuric acid a strong acid, shows a difference in the ease of ionization of the two hydrogen atoms. Since ionization, in general, is favored by dilution, we find that in the case of such a strong acid the difference is most marked in more concentrated solutions, the smaller amount of water starting the ionization in the more favored direction and producing first, chiefly, hydrogen-sulphate ions, HSO4−. When the solution is diluted, the hydrogen-sulphate ions are to a very considerable extent dissociated into sulphate ions and hydrogen ions. The described change in ionization can be roughly followed with the aid of an insoluble sulphate like barium sulphate. Barium sulphate, while very insoluble in water, dissolves in rather strong sulphuric acid to form the acid sulphate, Ba(HSO4)2, the SO42− ion of the sulphate being more or less suppressed by uniting with hydrogen-ion. We have the action
BaSO4 ⇄ Ba2+ + SO42− and
Ba2+ + SO42− + H+ + HSO4− ⇄ Ba2+ + 2 HSO4−.
If the solution of the acid sulphate is poured into a large volume of water, barium sulphate is immediately reprecipitated, the hydrogen-sulphate-ion being dissociated, in the dilute solution, into hydrogen-ion and sulphate-ion, SO42−, whose barium salt is so difficultly soluble:
Ba2+ + 2 HSO4− → Ba2+ + 2 H+ + 2 SO42− → BaSO4 ↓ + 2 H+ + SO42−.
Exp. Finely divided barium sulphate is warmed for a moment with a few cubic centimeters of concentrated sulphuric acid in a test tube, the mixture is allowed to settle, and some of the clear acid is carefully decanted into a large beaker full of water.
It may be added, that while the primary ionization of sulphuric acid does not yield an equilibrium constant for the ratio [H+] × [HSO4] / [H2SO4], even such a strong acid as is sulphuric acid is found to give a fairly good constant[189] for [H+] × [SO42−] / [HSO4−]. The value of this constant[189] is 0.03.
The Ionization Constants[A] of Acids
| Acid. | Equilibrium Ratio. | K. |
|---|---|---|
| Hydrochloric | [H+]×[Cl−]/[HCl] | (1) |
| Hydrobromic | [H+]×[Br−]/[HBr] | (1) |
| Hydroiodic | [H+]×[I−]/[HI] | (1) |
| Nitric | [H+]×[NO3−]/[HNO3] | (1) |
| Chromic[B] | [H+]×[HCrO4−]/[H2CrO4] | (1) |
| [H+]×[CrO42−]/[HCrO4−] | 0.6E−6 | |
| Sulphuric[C],[D] | [H+]×[HSO4−]/[H2SO4] | (1) |
| [H+]×[SO42−]/[HSO4−] | 0.3E−1 | |
| Oxalic[E] | [H+]×[C2O4−]/[H2C2O4] | 3.8E−2 |
| [H+]×[C2O42−]/[HC2O4−] | 0.5E−4 | |
| Phosphoric[F] | [H+]×[H2PO4−]/[H3PO4] | 0.1E−1 |
| [H+]×[HPO42−]/[H2PO4−] | 0.2E−6 | |
| [H+]×[PO43−]/[HPO42−] | 0.4E−12 | |
| Arsenic[D] | [H+]×[H2AsO4−]/[H3AsO4] | 0.5E−2 |
| Nitrous[B] | [H+]×[NO2−]/[HNO2] | 0.5E−3 |
| Acetic[G] | [H+]×[CH3CO2−]/[CH3CO2H] | 1.8E−5 |
| Carbonic[H],[I] | [H+]×[HCO3−]/([H2CO3]+[CO2]) | 0.3E−6 |
| [H+]×[CO32−]/[HCO3−] | 0.7E−10 | |
| Hydrogen Sulphide[J],[K] | [H+]×[SH−]/[H2S] | 0.9E−7 |
| [H+]×[S2−]/[SH−] | 0.1E−14 | |
| Boric[B] | [H+]×[H2BO3−]/[H3BO3] | 0.7E−9 |
| Hydrocyanic | [H+]×[CN−]/[HCN] | 0.7E−9 |
| Arsenious | [H+]×[H2AsO3−]/[H3AsO3] | 0.6E−9 |
| Water[L],[C] | [H+]×[HO−]/[H2O] at 25° | 0.2E−15 |
| at 100° | 0.9E−14 | |
| [H+]×[HO−] at 25° | 1.2E−14 | |
| at 100° | 0.5E−11 |