Three Theories of Acids and Alkalis.

ARRHENIUS DEFINITION  (1890)

This considered that acids were substances that released H⁺ ions in solution and that bases were substances that produced OH^- ions in solution.  It explained adequately how acids and bases neutralized each other

i.e.                                                        H⁺_(aq)     +     OH^‑_(aq)    ®      H₂O_(l)

                                                           acid                   base                    neutral       

However, as time passed the Arrhenius definition began to seem less satisfactory. For one thing, there seemed to be two kinds of base.  Metal hydroxides such as NaOH produce OH^- ions in water by ionic dissociation.  Bases, such as ammonia produce OH^- ions in aqueous solution by undergoing a reaction with water.

NaOH_(s)                        ®      Na⁺_(aq)     +     OH^-_(aq)

NH_3(g)     +     H₂O_(l)      ®     NH⁺_4(aq)     +     OH^-_(aq)

But more important, the Arrhenius definition of acids and bases was narrow.

It defined acids and bases in aqueous solution only.

THE LOWRY-BRONSTED DEFINITION    (1923)

Lowry-Bronsted produced  more general definitions of acids and bases.

These were in terms of H⁺ ions (called a protons, since a H⁺ ion has neither electrons nor neutrons).

            According to Bronsted - an acid is a substance that can donate a proton

                                                            a base is a substance that can accept a proton

 This definition can be represented by the general chemical reaction

A     D     B    +    H⁺

which  does not attempt to show electrical charge balance.

In this equation -

·      A is the acid.,

·      B is the base and

·      H⁺ (a hydrogen atom without an electron) is a proton. 

Together A and B are called a CONJUGATE ACID AND BASE PAIR.

We can call B the conjugate base of A and call A the conjugate acid of B. This Lowry-Bronsted definition is more general than the Arrhenius definition.  It does not refer to any specific solvent. We can simply say that any substance that can lose a proton is an acid.  The following are examples of Bronsted acids -

Similarly the Bronsted bases are substances which can accept protons and include the following examples :-

e.g.  Mg(OH)⁺     +     H⁺     =     Mg²⁺     +     H₂O

While the relationship   A ® B + H⁺    is a very good general definition of an acid and a base, there is a problem in applying this definition to an acid/base system in solution.  The problem arises because the free protons (H⁺) cannot exist in solution to any great extent.  In many cases the protons interact with the solvent

e.g.       H₂O_(l)     +     H⁺_(aq)      ®     H₃O⁺_(aq)  (a hydroxonium ion)

The solvent is then acting as a Bronsted base in accepting a proton.

The same reasoning can be applied to a base in solution.  The base must obtain a proton from a proton source.  Very often the proton source is the solvent.

                                                   H₂O

                                       NH_3(g)                 NH⁺_4(aq)

The solvent is then acting as a Bronsted acid by providing a proton. 

We can see that acid base reactions in solution are not simply processes in which an acid loses a proton or a base gains a proton.  We must regard all acid/base reactions in solution as PROTON TRANSFER REACTIONS.  Any such reaction includes TWO ACID - BASE CONJUGATE PAIRS.  The original acid base pair, plus another conjugate pair, to accept the proton from the acid or donate the proton to the base.  As we have said, this second acid-base pair is often derived from the solvent.

Each substance in the acid column is converted to its partner in the base column by the removal of a proton. Each substance in the base column is converted to its partner in the acid column by the addition of a proton.

ANY TWO CONJUGATE ACID/BASE PAIRS CAN BE C0MBINED IN AN ACID/BASE REACTION

e.g.                   HCl     +     OH^-     D     H₂O     +     Cl^-

HCl and Cl^- are one conjugate acid-base pair and H₂O  and OH^-  are another pair.  The proton donated by the acid HCl is accepted by the base  OH^- in the forward reaction and a proton accepted by the Cl^- is donated by the H₂O in the reverse reaction.

The Bronsted definition can be summarized  in a few sentences. 

An acid is -

·      a proton donor

·      a base a proton acceptor.

We emphasize the relationship by designating conjugate acid base pairs.

e.g.       HCl/Cl^-

Acid base reactions in solution require two conjugate acid/base pairs because free protons (H⁺) do not exist in solution

THE LEWIS DEFINITION  (1938)

The Lowry-Bronsted theory has its short-comings.  Since it defines an acid as a proton donor,  it excludes from the category of acids, substances that have no protons to donate.  This limitation was overcome by a more general definition of acids and bases, BASED ON ELECTRONIC STRUCTURE,  which was proposed by G N Lewis.

By the Lowry-Bronsted definition, a substance must accept a proton to be classified as a base.  In other words the base must form a chemical bond with the proton (H⁺ ion).  Since the proton has no electrons the base must have an electron pair available to form a bond.

In the Lewis definition:-

A BASE IS A SUBSTANCE THAT HAS A NON-BONDING VALENCE ELECTRON PAIR THAT CAN BE USED TO FORM A CHEMICAL BOND.

more simply - a Lewis base is an electron - pair donor.

The Lewis definition does not greatly expand our ideas about bases, but it does significantly broaden the category of substances that can be classified as acids.  When a base accepts a proton it donates electrons to form the bond with the proton.  Therefore the proton accepts the electrons.

A LEWIS ACID IS A SUBSTANCE THAT IS AN ELECTRON-PAIR ACCEPTOR

A proton (H⁺) is the simplest Lewis acid but many other substances fit the definition,  including a number of cations,  e.g Zn²⁺

Zn²⁺     +     2OH^-       D     Zn(OH)₂

Here the Zn²⁺ is the Lewis acid and OH^-  the Lewis base.

In organic chemistry it is especially common to find cations that behave as Lewis acids., accepting electrons.  The reaction :-

C₂H₄     +     Br₂         C₂H₄Br₂

can be considered to proceed in 3 steps.

1.     Br₂ forms Br⁺ and Br    ions.  The Br⁺  is the Lewis acid.

2.     This joins to C₂H₄  (a Lewis base).

3.     In the third step Br^- , a Lewis base, donates an electron pair and joins to the cation formed in the first step.  The cation is a Lewis acid

              H          H                                                                    H         H

Br           C          C     +          +          Br                       Br          C          C          Br

              H         H                                                                    H          H

A great many reactions can be understood as the combination of a Lewis acid and a Lewis base.  In fact a common method of classifying chemical reagents uses the Lewis acid and Lewis base definition.

A LEWIS ACID, WHICH IS A SUBSTANCE SEEKING ELECTRONS, IS CALLED AN ELECTROPHILE.

A LEWIS BASE, WHICH IS AN ELECTRON DONOR (OR A SUBSTANCE THAT SEEKS SUBSTANCES THAT ARE ELECTRON DEFICIENT), CAN BE CALLED A NUCLEOPHILE.

In the above reaction between C₂H₄ and Br₂, Br⁺ is an electrophile and Br^- is a nucleophile.

Many Lewis acids are substances that are chemically neutral.  These substances can accept electrons and form additional bonds, either because they have incomplete octets or octet expansion is possible.

Many compounds of the elements of Group III are Lewis acids, the best known being the halides such as BF₃, BCl₃, AlCl₃ and AlBr₃

Question:  Use the Lewis theory to explain what is happening when AlCl₃ molecules dimerise

                                    Cl                     Cl                     Cl

                                                Al                     Al

                                    Cl                     Cl                     Cl

Co-ordinate (dative covalent) bond in Al₂Cl₆

OTHER EXAMPLES INCLUDE -

            BF₃     +     NH₃              F₃B          NH₃

   Lewis Acid     Lewis Base

            AlCl₃     +     Cl^-                    AlCl₄

   Lewis Acid     Lewis Base

Conclusion.

The Lowry-Bronsted definition is quite adequate for the chemistry of acids and bases in aqueous solution.  The Lewis definition is useful because it allows us to classify so many substances as acids or bases, and because it provides an understanding of the details of many chemical reactions. 

  Chemistry Department: Loreto College, Coleraine.