# In this example, which species acts as an acid and why? H_2SO_4^- +H_2O -> SO_4^(2-) + H_3O^+

Jul 1, 2016

The sulfuric acid and bisulfate ion are the acids. Water molecules are the base.

#### Explanation:

An acid is a proton donor; a base is a proton acceptor. Acids and bases are substances that respectively increase concentrations of the characteristic cation and the characteristic anion of the solvent, and here the solvent is water.

Now sulfuric acid is diprotic in water. We can separate out 2 successive acid base reactions:

${H}_{2} S {O}_{4} + {H}_{2} O \rightarrow {H}_{3} {O}^{+} + H S {O}_{4}^{-}$

$H S {O}_{4}^{-} + {H}_{2} O \rightarrow {H}_{3} {O}^{+} + S {O}_{4}^{2 -}$

In each reaction, the base, the proton acceptor, is water.

You will note that acid/base reactions usually designate water as the solvent. Chemists can use other solvent systems, but at this level (?) stick to the water-based (hydroxide and hydronium) ion definition.

As far as anyone knows, the hydronium or acidium species, ${H}_{3} {O}^{+}$, is a cluster of water molecules, 3 or 4, with an extra ${H}^{+}$ to give ${H}_{7} {O}_{3}^{+}$ or ${H}_{9} {O}_{4}^{+}$; this cluster can easily exchange the ${H}^{+}$ with other water clusters; think of a rugby maul. We write ${H}_{3} {O}^{+}$ or even ${H}^{+}$ for convenience, and we can certainly do calculations on this basis.