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Fabric Stain Removal Guide

  • Acids and Bases

The use of acid and base as classifications of chemical substances is an ancient one and predates even alchemy. Acids were originally substances that tasted sour and reacted with bases; bases were alkaline salts such as CaO (lime), NaOH, and KOH, while amphoteric substances were substances which could react with either acids or bases. In modern chemistry, these concepts have taken on considerably more precise meaning. The chemistry of acids and bases is the chemistry of reactions in which protons are transferred.

Arrhenius Theory of Acid-Base Behavior

The first quantitative approach to acid-base equilibria was developed around 1884 by the Swedish chemist Svante Arrhenius. According to his Arrhenius theory, an acid is any compound or ion which yields hydrogen ions in water solution and a base is any compound or ion which yields hydroxyl ions in water solution.

The approach taken by Arrhenius with acids and bases was in accordance with his general approach to substances in aqueous solution. Materials which dissolved in whole or in part as ions were, and still are, called electrolytes. A strong electrolyte, like sodium chloride, was a substance which dissociated completely into ions while a weak electrolyte remained partially in molecular form. Thus a strong acid is a strong electrolyte which ionized completely to give hydrogen ions in aqueous solution. All of the common strong acids - sulfuric acid, H2SO4; nitric acid, HNO3; hydrochloric acid, HCl; and perchloric acid, HClO4 - easily fit in this category. Likewise, a strong base is a strong electrolyte which ionized completely to give hydroxyl ions in aqueous solution. The common strong bases, NaOH and KOH, fit in this category. Acid-base reactions are then simply the reaction of the hydrogen ion with the hydroxyl ion as a neutralization reaction, H+(aq) + OH-(aq) --> H2O. The remaining ions simply continue to exist as a solution of a salt. The reaction including these "bystander ions" could be written as

H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) --> Na+(aq) + Cl-(aq) + H2O

The Arrhenius concept was not intended to, and does not, deal with acid-base chemistry in solvents other than water. Even in water, however, it has deficiencies. Ammonia, NH3, clearly behaves in a qualitative manner as a base and reacts to neutralize acids. However, ammonia does not contain a hydroxide ion. This is of itself no obstacle because ammonia might well produce hydroxyl ion by reaction with water:

NH3(aq) + H2O --> NH3.HOH --> NH4OH --> NH4+(aq) + OH-(aq)

However, no experimental evidence for the existence of NH4OH has yet been foumd. For these two reasons, the Arrhenius concept is no longer considered an adequate treatment of acids and bases even in aqueous solution.

Bronsted-Lowry Theory of Acid-Base Behavior

The Arrhenius theory was the only theory used to explain the behaviour of acids and bases for about forty years. In 1923, J. N. Bronsted in Denmark and T. M. Lowry in England independently, and almost simultaneously, proposed the modern "protonic" or Bronsted-Lowry theory of acid-base behaviour. According to the Bronsted-Lowry concept, an acid is any compound or ion which can give up a proton, while a base is any compound or ion which can accept a proton.

Any species which can either accept or give up a proton is said to be amphiprotic. Thus the water molecule is amphiprotic, since it can give up a proton, H2O --> H+ + OH-, to form the hydroxyl ion OH-. Alternatively, water can accept a proton to form the hydronium ion H3O+, according to the equation H+ + H2O --> H3O+. The above two equations can be combined to give the dissociation equation for water: 2H2O --> H3O+ + OH-.

The Bronsted-Lowry concept is an extension of the Arrhenius concept in that all Arrhenius bases, being sources of hydroxide, can accept protons. Ammonia and amines will also accept protons to form the corresponding ammonium ions, so the existence of NH4OH is no longer necessary to explain the action of ammonia as a base. The Bronsted-Lowry concept also is useful in protonic solvents other than water, such as liquid ammonia or glacial acetic acid, where the Arrhenius concept is not useful. We will, however, generally confine our discussion to aqueous solutions because they are so much more important.

Lewis Explanation of Acid-Base Behavior

The basic principles of the Lewis theory of acid-base behaviour were also set down in 1923, by the American physical chemist G. N. Lewis. The Lewis definitions of acids and bases are even more inclusive than the Bronsted definitions. The Lewis definitions are that an acid is an electron-pair acceptor and a base is an electron-pair donor.

Since a base like ammonia, H3N:, has a lone pair of electrons, it can be considered to "donate" them to a proton in forming the conjugate acid NH4+. The Lewis definitions are used to explain the effect of compounds such as AlCl3, which acts as an acid in nonaqueous organic solvents, on organic reactions. In protonic solvents, however, the Lewis definitions are far less useful than are the Bronsted definitions.

 

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