Acids and Bases

 

Acids and bases can be distinguished by means of a variety of properties (Table 1). Some properties of acids and bases are more useful than others to a chemist, especially those that can be used as diagnostic tests, such as the litmus test.

 

 

Table 2 shows that solutions and melts of ionic compounds and bases conduct electricity.

 

 

Table 2 also shows that melts of molecular substances, including acids, do not conduct

electricity in their pure states. However, acids (unlike other molecular substances)

become conductors when dissolved in water.  Can we conclude, therefore, that acids in their pure forms contain neutral molecules, but that acids in aqueous solution contain ions?

 

The Arrhenius Theory of Acids and Bases

 

In 1887 Svante Arrhenius created a theory of ions to explain the electrical conductivity

of solutions suggesting that acids are hydrogen compounds that ionize in aqueous solution to increase the hydrogen ion concentration of the solution.

Arrhenius assumed that the water somehow causes the acid molecules to ionize, but he didn’t propose an explanation for this. (We now believe that water molecules help to pull the molecules apart—to ionize the acid.) A typical example of an acid is hydrogen chloride gas dissolving in water to form hydrochloric acid. We can describe this process with an ionization equation.

 

So, although HCl is a molecular compound, it appears to behave in aqueous solution as if it were ionic. It ionizes into ions, which are capable of conducting electricity in solution. 

 

 

 

 

Arrhenius was also able to explain in a theoretical way why bases have their characteristic properties. He suggested that bases are ionic hydroxides that dissolve in water to increase the hydroxide ion concentration of the solution.  The dissociation of bases is similar to that of any other ionic compound.

 

This theory also explained the process of neutralization by assuming that H+(aq) and OH (aq) ions combine to form H2O(l).  The various strengths of acids were explained in terms of the degree (percentage) of ionization.

 

 

Limitations of Arrhenius’ Definitions

 

1- Arrhenius’ definitions cannot always predict whether a substance is an acid or a base. According to Arrhenius, only substances containing hydrogen could be classified as acids, and only those containing hydroxide ions could be classified as bases.  Accordingly, salts should theoretically be neutral.  But this is not always the case.  The same could be said of oxides of metals and non-metals.

 

2- The role of water as a solvent and the nature of the hydrogen ion.

It is highly unlikely that the particle we call an aqueous hydrogen ion, H+(aq), actually exists in an acidic solution. If such a particle were to come near polar water molecules, it would bond strongly to one or more of the molecules (Figure 3), that is, it would be hydrated.

Figure 3

The Lewis (electron dot) diagram for a hydrogen ion has no electrons. A water molecule is believed to have two lone pairs of electrons, as shown in its Lewis diagram. The hydrogen ion (proton) is believed to bond to one of these lone pairs of electrons to produce the H 3 O + ion.

 

        

 

 

 There is no evidence for the existence of unhydrated hydrogen ions in aqueous solution. However, the Canadian scientist Paul Giguère has done experiments that provide clear evidence for the existence of hydrated protons. The simplest representation of a hydrated proton is H3O+(aq), commonly called the hydronium ion.

Acid and base definitions, revised to include the ideas of the hydronium ion and reaction with water, are more effective in describing, explaining, and predicting the behaviour of acids and bases than are Arrhenius’ original definitions.  However, chemical research has shown that even these revised definitions are still too restrictive.  A broader concept is needed to describe, explain, and predict these properties of acids and bases. 

 

 

 

The Brønsted–Lowry Theory of Acids and Bases

 

In 1923 the Danish and British chemists independently developed a more general theory of acids and bases that was able to explain the role of the solvent as well as the existence of acidic and basic salt solutions. 

They defined acids as substances that donate protons (H+) and bases as substances that accept protons in a chemical reaction. 

Therefore, an acid is a hydrogen -ion donor, and a base is a hydrogen-ion acceptor (any molecule or ion that has a lone pair of electrons available for bonding with a hydrogen-ion).  In other words, acid/base reactions are proton transfer reactions.

According to the Brønsted-Lowry concept, hydrogen chloride donates a proton (H+) to a water molecule,

We can describe this reaction as NH3 molecules removing protons from H3O- ions. Hydronium ions act as the acid, and ammonia molecules act as the base. Water is present as the solvent, but not as a primary reactant. In fact, water does not even have to be present, as evidenced by the reaction of hydrogen chloride and ammonia gases.

 

A substance can be classified as a Brønsted-Lowry acid or base only for a specific reaction. It is not a general property of a substance. This point is important —a substance may gain protons in one reaction, but lose them in another reaction with another substance. (For example, in the reaction of HCl with water shown above, water acts as the Brønsted-Lowry base; whereas, in the reaction of NH3 with water, water acts as the Brønsted-Lowry acid.) A substance that appears to act as a Brønsted-Lowry acid in some reactions and as a Brønsted-Lowry base in other reactions is called amphiprotic (or amphoteric). The hydrogen carbonate ion (HCO-3(aq)) in baking soda is amphiprotic, like every other hydrogen polyatomic ion. Hydrogen polyatomic ions, as their name suggests, are polyatomic ions containing hydrogen. Examples of amphiprotic substances include:

 

Note that amphiprotic entities can either gain or lose a proton, as shown by the following reactions. First let’s see what happens when the bicarbonate ion is added to the solution of a strong acid, which will contain hydronium ions.

Now let’s look at the reaction of the bicarbonate ion with the solution of a strong base, which will contain hydroxide ions.

 

Definitions of Acids and Bases

Arrhenius Definitions

• An acid ionizes in water to increase the hydrogen ion concentration.

• A base dissociates in water to increase the hydroxide ion concentration.

• A neutralization reaction involves the reaction of a hydrogen ion with a

hydroxide ion.

 
Revised Arrhenius Definitions

• An acid reacts with water to increase the hydronium ion concentration.

• A base reacts with water to increase the hydroxide ion concentration.

• A neutralization reaction involves the reaction of a hydronium ion with a

hydroxide ion.

 

Brønsted–Lowry Definitions

• An acid is a proton donor.

• A base is a proton acceptor.

• An acid–base neutralization reaction involves the transfer of one proton

from the strongest acid present to the strongest base present.

• An amphiprotic substance is one that appears to act as a Brønsted-Lowry

acid in some reactions and as a Brønsted-Lowry base in other reactions.

• A conjugate acid–base pair consists of two substances that differ only by

one proton.