Common Acid-Base Theories

• Arrhenius Theory

    Augustus Svante Arrhenius proposed in his study of          electrolytic dissociation that in an aqueous solution, a strong electrolyte only exists in the form of ions, whereas a weak electrolyte exists partly as ions and partly as molecules. Years later, specifically three, Arrhenius extended this theory by suggesting that acids are neutral compounds that ionize when dissolved in water to give hydrogen ions and a corresponding negative ion. Bases are neutral compounds that either dissociate or ionize in water to give hydroxide ions and a positive ion.

    An Arrhenius Acid, therefore, is any substance that ionizes when it dissolves in water to give hydrogen ion(s). For example,

       HCl    yields to   H⁺ _(aq) + Cl^-_(aq)

     _{An Arrhenius Base is any substance that gives hydroxide ions when it dissolves in water. For example,}

             NaOH    yields to     Na⁺_(aq) + OH^-_(aq)   

    The Arrhenius Theory has several disadvantages. It is only applicable for aqueous solutions. The theory also does not explain why not all compounds containing a proton give an acidic solution and why substances not containing hydroxides are able to produce basic solutions.   

• Bronsted-Lowry Theory

    One change suggested by Johannes Bronsted and Thomas Lowry was the existence of a hydronium ion. They proposed that hydrogen ions cannot exist in water but rather covalently bond with water. They based their definitions of acids and bases on the behavior of hydrogen ions. Acids are substances from which a proton can be removed. bases are substances that binds protons. It can be simplified as Bronsted-Lowry acids being proton donors and Bronsted-Lowry bases as proton acceptors. Take note that protons are not actually being given but rather the bond between the proton and another element is being broken to form a new bond to the base. This new definition broadens the range of acids and bases but still requires a protic acid. The concept of conjugates is an additional concept introduced by this theory. A conjugate base is the species remaining when a proton is removed from an acid. A conjugate acid is the species formed when a proton is transferred to a base. For example,

• Lewis Theory

    The theory proposed by Lewis further widens the classification of acids and bases to non-aqueous solutions. He proposed that the electrons and not the protons are the ones transferring. The acids and bases can be charged or unionized. A Lewis acid is a species that can accept a pair of electrons while a Lewis base is a species that can donate a pair of electrons. Note that these species can be ions or molecules. For example,

•  Solvent System Theory

     One of the limitations of Arrhenius definition was its reliance on water solutions. E. C. Franklin studied the acid-base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius theory, and Albert F. O. Germann, working with liquid COCl₂, generalized Arrhenius definition to cover aprotic solvents and formulated the solvent system theory in 1925.

     

     Germann pointed out that in many solvents there is a certain concentration of a positive species, solvoniumsolvate (earlier lyate) anions, in equilibrium with the neutral solvent molecules. For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and ammonium and amide, respectively: (earlier lyonium) cations and negative species,

2 H₂O H₃O⁺ + OH⁻

2 NH₃ NH+ 4 + NH− 2

     Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride.

N₂O₄ NO⁺ + NO− 3

2 SbCl₃ SbCl+ 2 + SbCl− 4

COCl₂ COCl⁺ + Cl⁻

     A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is defined as an acid and one causing the reverse is defined as a base. Thus, in liquid ammonia, KNH₂NH−
2) is a strong base, and NH₄NO₃ (supplying NH+
4) is a strong acid. In liquid sulfur dioxide (SO₂), thionyl compounds (supplying SO²⁺) behave as acids, and sulfites (supplying SO2−
3) behave as bases. (supplying (

The non-aqueous acid-base reactions in liquid ammonia are similar to the reactions in water:

2 NaNH₂ (base) + Zn(NH₂)₂ (amphiphilic amide) → Na₂[Zn(NH₂)₄]

2 NH₄I (acid) + Zn(NH₂)₂ (amphiphilic amide) → [Zn(NH₃)₄)]I₂

     Nitric acid can be a base in liquid sulfuric acid:

HNO₃ (base) + 2 H₂SO₄ → NO+ 2 + H₃O⁺ + 2 HSO− 4

     The unique strength of this definition shows in describing the reactions in aprotic solvents, for example in liquid N₂O₄:

AgNO₃ (base) + NOCl (acid) → N₂O₄ (solvent) + AgCl (salt)

      Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, HClO₄ is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid. This was seen as both a strength and a weakness, since some substances, such as SO₃ and NH₃ were felt to be acidic or basic on their own right. On the other hand, solvent system theory was criticized as too general to be useful; it was felt that there was something intrinsically acidic about hydrogen compounds, not shared by non-hydrogenic solvonium salts.