Chemical characteristics of Acid

• Monoprotic acids

Monoprotic acids are those acids that are able to donate one proton per molecule during the process of dissociation (sometimes called ionization) as shown below (symbolized by HA): HA(aq) + H2O(l) is in equilibrium with H3O+(aq) + A−(aq) Ka Common examples of monoprotic acids in mineral acids include hydrochloric acid (HCl) and nitric acid (HNO3). On the other hand, for organic acids the term mainly indicates the presence of one carboxyl group and sometimes these acids are known as monocarboxylic acid. Examples in organic acids include formic acid (HCOOH), acetic acid (CH3COOH) and benzoic acid (C6H5COOH).

• Polyprotic acids

Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic acid (two potential protons to donate) and triprotic acid (three potential protons to donate). A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2. H2A(aq) + H2O(l) is in equilibrium with H3O+(aq) + HA−(aq) Ka1 HA−(aq) + H2O(l) is in equilibrium with H3O+(aq) + A2−(aq) Ka2 The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2. For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion (HSO4−), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO42−), wherein the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2CO3) can lose one proton to form bicarbonate anion (HCO3−) and lose a second to form carbonate anion (CO32−). Both Ka values are small, but Ka1 > Ka2 . A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3. H3A(aq) + H2O(l) is in equilibrium with H3O+(aq) + H2A−(aq) Ka1 H2A−(aq) + H2O(l) is in equilibrium with H3O+(aq) + HA2−(aq) Ka2 HA2−(aq) + H2O(l) is in equilibrium with H3O+(aq) + A3−(aq) Ka3 An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three protons can be successively lost to yield H2PO4−, then HPO42−, and finally PO43−, the orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. Even though the positions of the protons on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

• Neutralization

Hydrochloric acid (in beaker) reacting with ammonia fumes to produce ammonium chloride (white smoke). Neutralization is the reaction between an acid and a base, producing a salt and neutralized base; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) Neutralization is the basis of titration, where a pH indicator shows equivalence point when the equivalent number of moles of a base have been added to an acid. It is often wrongly assumed that neutralization should result in a solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction. Neutralization with a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidic ammonium chloride, which is produced from the strong acid hydrogen chloride and the weak base ammonia. Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt, e.g. sodium fluoride from hydrogen fluoride and sodium hydroxide.

• Weak acid/weak base equilibria

In order to lose a proton, it is necessary that the pH of the system rise above the pKa of the protonated acid. The decreased concentration of H+ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H+ concentration in the solution to cause the acid to remain in its protonated form, or to protonate its conjugate base (the deprotonated form). Solutions of weak acids and salts of their conjugate bases form buffer solutions.