1661 - Robert Boyle

Characterized acids and alkalies (bases) as the following:

Acids: Sour taste Corrosive Change litmus (dye extracted from lichens) from blue to red Become less acidic when combined with alkalies. Alkalies (Bases): Feel slippery Change litmus from red to blue Become less alkaline when combined with acids.

Antoine Lavoisier

Believed that all acids contained oxygen after studying several acids e.g.  H2SO4 - sulfuric acid, HNO3 - nitric acid

1811 - Humphry Davy

Questioned Lavoisier's theory, noting that hydrochloric acid (HCl) did not contain oxygen yet is an acid. Soon thereafter, several more acids without oxygen were found. e.g.  HBr, HF, HI

1838 - Justig Liebig

• Suggested that acids contain one or more hydrogen atoms which can be replaced by metal atoms to produce salts.

  e.g.  HSCN is an acid because the H atom can be replace by a metal to form a salt, such as NaSCN.

1884-1887 - Svante Arrhenius

In 1884, Arrhenius proposed that salts dissociate when they dissolve in water to give charged particles which he called ions. In 1887, Arrhenius extended this his idea by defining acids and bases as the following: Arrhenius acid - Any substance that ionizes when it dissolves in water to give the H+ ion. e.g.   Arrhenius base - Any substance that ionizes when it dissolves in water to give the OH- ion. e.g.

Arrhenius's theory helped to explain why acids have similar characteristics, since they all give H⁺ ions when they dissolve in water.  It also explained why acids are neutralized by bases and why bases are neutralized by acids;  the H⁺ ions from acids combine with the OH^- ions from bases to form water:

Though the Arrhenius theory helped to explain more about acids and bases, there were still several drawbacks to this theory.

1. The theory can only classify substances when they are dissolved in water since the definitions are based upon the dissociation of compounds in water.
2. It does not explain why some compounds containing hydrogen such as HCl dissolve in water to give acidic solutions and why others such as CH₄ do not.
3. The theory can only classify substances as bases if they contain the OH^- ion and cannot explain why some compounds that don't contain the OH^- such as Na₂CO₃ have base-like characteristics.

To extend the Arrhenius theory a little further, consider the formation of water from the combination of an H⁺ ion and an OH^- ion:

This reaction is actually reversible, represented by the forward/backward arrow in the following reaction:

Based on the fact that the above reaction is reversible, we can conclude the following operational definitions for acids and bases.

• Acid - Any substance that increases the concentration of the H⁺ ion when it dissolves in water.
• Base - Any substance that increase the concentration of the OH^- ion when it dissolves in water.

NOTE: A common shorthand notation for the concentration of a substance is placing it in brackets:

[H⁺] = concentration of the H⁺ ion

Now, substances not containing H⁺ or OH^- ions can be classified as acids or bases if they alter the [H⁺] or [OH^-] when they dissolve in water.

e.g.  CO₂ cannot dissociate to give H⁺ but it does increases [H⁺] when it is dissolves in water.

     e.g.  CaO cannot dissociate to give OH^- but it does increase [OH^-] when it dissolves in water.

1923 - Johannes Brønsted and Thomas Lowry

Johannes Brønsted	In 1923, Johannes Brønsted and Thomas Lowry separately proposed a new set of defintions for acids and bases which are known as either Brønsted acids and bases or Brønsted-Lowry acids and bases. Brønsted Acid - Any substance that can donate a proton, H+ ion to a base. aka:  hydrogen-ion donors or proton donors Brønsted Base - Any substance that can accept a proton, H+ ion from an acid. aka:  hydrogen-ion acceptor or proton acceptor In the above reaction, the H from HCl is donated to H2O which accepts the H to form H3O+, leaving a Cl- ion.  The dissociation of water can be represented as follows:
Thomas Lowry

The Brønsted-Lowry model of acids and bases brings rise to the concept of conjugate acid-base pairs.  The part of the acid remaining when an acid donates a H⁺ ion is called the conjugate base.  The acid formed when a base accepts a H⁺ ion is called the conjugate acid.  For the generic acid HA:

For the generic base A^-:

More examples of conjugate acid-base pairs:

In the following reactions, it is shown how H₂PO₄^- and H₂O can act as both acids and bases.  Such compounds are said to be amphoteric.

• Strong acids have weak conjugate bases.
• Strong bases have weak conjugate acids.

Water has the tendency to equalize the strengths of all strong acids and strong bases, regardless of the strength of the acid itself.  This is known as the leveling effect.  Acids are limited to the strength of the H₃O⁺ ions that they form when they lose H⁺ ions when they dissolve in water.  Likewise, bases are limited to the strength of the OH^- ions that they form when they gain H⁺ ions when they dissolve in water.

The relative strength of an acid is described as an acid-dissociation equilibrium constant.

• Acid-dissociation equilibrium constant (K_a) - For the generic acid reaction with water:

The acid-dissociation equilibrium constant is the mathematical product of the equilibrium concentrations of the products of this reaction divided by the equilibrium concentration of the original acid:

Strong acids dissociate almost completely in water and therefore have relatively large K_a values.  Weak acids, however, dissociate only slightly in water and therefore have relatively small K_a values.  To distinguish between strong and weak acids, the following guidelines are used:

The relative strengths of acids and bases is discussed further in the Equilibrium notes.

Advantages to the Brønsted-Lowry model of acids and bases

• Acids and bases can now be ions or neutral molecules.
• Acids and bases can now be any molecule with at least one pair of nonbonding electrons.
• It explains the role of water in acid-base reactions; H₂O accepts H⁺ ions from acids to form H₃O⁺ ions.
• It can be applied to solutions with solvents other than water and even in reactions that occur in the gas or solid phases.
• It relates acids and bases to each other with conjugate acid-base pairs and can explain their relative strengths.
• It can explain the relative strengths of pairs of acids or pairs of bases.
• It can explain the leveling effect of water.

1923 - G.N. Lewis

Proposed another method of defining acids and bases. Lewis acid - Any substance that can accept a pair of nonbonding electrons. aka:  electron-pair acceptor Lewis base - Any substance that can donate a pair of nonbonding electrons. aka:  electron-pair donor

In the following example, the Al³⁺ ion acts as an acid, accepting electron pairs from water which acts as a base, an electron-pair donor.  The two combine to form Al(H₂O)₆³⁺, an acid-base complex or a complex ion.

Next:  "Typical Acids and Bases"