What is an Acid?

The Arrhenius Definition

Svant Arrhenius

According to Arrhenius (pictured left) an acid disassociates in an aqueous solution to give H₃O⁺ ions (Hydronium ion, which is a protonated water molecule as a bare proton H+ cannot exist in aqueous solution as a free species).

An Arrhenius base on the other hand disassociates in aqueous solution to give OH- ions

According to this definition:

Acids                                    Bases

H₂SO₄ ,                               NaOH

HCl ,

HNO_{3 etc}                                      Ca(OH)_{2 etc}

Credits: Z(J)ubin

What The Fact... Most students, thus tend to believe that bases contain an OH group, but they fail to realize that some acids too have an OH group attached to them, but still disassociate to give H+ (or H₃O⁺ ions).  eg H₂SO₄

Well this anomaly of sorts can be explained on the basis of difference in electronegativity.

In case of the H₂SO₄   molecule, the difference in electronegativity between Oxygen and Hydrogen is greater than the difference in electronegativtiy between Oxygen and Sulfur. Therefore the bond between Oxygen and Hydrogen polarizes giving HSO₄^-  and H⁺   Refer to structure pictured below.

In case of a base (NaOH) the difference in electronegativity between Sodium and Oxygen is greater than the difference in electronegativity between Hydrogen and Oxygen, thus the bond between Oxygen and Sodium polarizes giving Na⁺   and OH^-  Refer to structure pictured below.

<Tip: Think of a bond as a match of tugowar in which the more electronegative element will pull the electrons towards itself>

Drawbacks: This theory fails to account for the acidic behavior of compounds like BCl3 and basic nature of compounds like NH3.

Brønsted–Lowry Definition

Brønsted and Lowry defined an acid as a proton donor and a base as a proton acceptor.

The removal of a proton (hydrogen ion) from an acid gives its conjugate base, and the addition of a proton to a base produces its conjugate acid. 

For instance:

• HCl + H₂O   Cl^- + H₃O⁺            

Conjugate acid:  H₃O⁺

Conjugate base: Cl^- 

• NH₃ + H₂O    OH^- + NH₄⁺

Conjugate acid:  OH^-

Conjugate base: NH₄⁺

Note: Weak acids give strong conjugate bases and weak bases give strong conjugate acids. Whereas strong acids give weak conjugate bases and strong bases give weak conjugate acids.

Water being amphoteric acts as both acid and a base. For instance in the reaction with Hydrochloric acid, water acts as a base but in the reaction with Ammonia water acts as an acid.

H₂O + H₂O  H₃O⁺ + OH⁻

Lewis Definition

Gilbert N Lewis removed the hydrogen (or proton requirement) of the Arrhenius concept and Bronsted-Lowry Definition and instead based his definition on electron pairs.

Lewis acids thus are compounds that are electron deficient and thus can accept a lone pair of electrons 

eg: BCl₃ 

In BCl₃  Boron is sp2 hybridised and it is made apparent from the diagram pictured below it has a vacant p-orbital and is thus electron deficient.

Relation between alkalis and bases

And Lewis base is a compound that can donate a lone pair of electron 

eg. NH₃

Note: On an unrelated note, a lot of people tend to use the terms alkali and base interchangeably, this however is incorrect. All alkalis are bases, but only water soluble bases are alkalis. 

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That's all folks.

The Passive Observer Out!