An Introduction to Acids and Bases
Chemistry is the study of matter and the changes it undergoes. One of the challenges that the chemist faces in trying to understand matter is to divide it into categories in much the same way that a biologist assigns plants and animals to categories called phyla (divisions), classes, etc.. Although the early chemists known as alchemists were really not concerned about understanding the fundamental nature of matter, they used types of matter that they referred to as acids and bases in their quest to change ordinary metals into gold. Acids have a sour taste (we now know that chemicals should never be tasted because many are toxic) and change the color of litmus (a plant dye) red, while bases have a slippery feel and turn litmus blue. Probably the most important property of acids and bases observed by the alchemists was the often violent reaction when they were mixed with one another. This reaction of an acid with a base is called neutralization because it results in a neutralization of their properties.
In the 1700's, scientists such as the Frenchman Lavoisier became interested in composition, that is, what elements are present in certain types of matter. This was the beginning of laboratory research and of what we now call the “scientific method.” Lavoisier, who discovered the element oxygen (along with Priestley), did a variety of experiments on oxygen and was convinced that oxygen was present in every acid. A student of Lavoisier, Davy, finally convinced the scientific community, but not Lavoisier, that hydrogen, not oxygen, was actually the element responsible for the properties of acids.
In the late 1800s, the Swedish chemist Arrhenius argued that acids could be defined as substances that produce the hydrogen ion (H+) in water and that bases produce the hydroxide ion (OH-) in water. These definitions were particularly important because most acid-base reactions were then, and are now, carried out in that amazing solvent water. According to this definition, acetic acid, the ingredient in vinegar that gives it the sour taste, is an acid because it reacts as follows:
CH3COOH --> H+ + CH3COO-
A neutralization reaction involves the reaction of the hydrogen ion with the hydroxide ion to form water:
H+ + OH- --> H2O
Although acid-base and many other reactions were (and still are) carried out in water, Arrhenius was not concerned about the role of water in the reaction.
In the early 1900s scientist began to explore reactions in greater detail, including the role of the solvent in which they took place and it was realized that a hydrogen ion could not exist as an individual entity in water. The hydrogen atom consists after all of just a proton and an electron, and when the electron is stripped off only the positively charged ion H+ remains. This proton is about one ten thousandths the size of the neutral hydrogen atom. Not only is it very small, but its ratio of charge to volume, called the charge density, is very high. Many aspects of chemical reactions depend on charge and size and the proton has one of the largest charge densities of any species. For example, the charge density of the sodium ion, Na+, is lower than the charge density of the plus three aluminum ion (Al3+), both of which are considerably smaller than the charge density of H+. Species with such high charge density are attracted very strongly to the electron density around the water molecule (which is a result of its unbonded or lone pair electrons).
In thinking about the lone pair of electrons on the water molecule it is important to realize that in the early 1900s there was no good model to describe the location or nature of the electrons in molecules. By 1920 G. N. Lewis had proposed his electron dot model, and the water molecule was pictured as shown below, with two O-H covalent bonds and two nonbonded pairs of electrons
on the oxygen atom. These nonbonded pairs of electrons are the source of electron density to which the hydrogen ion is attracted. The new species formed as a result of that attraction is called the hydronium ion:
This knowledge of the nature of the hydrogen ion in water now allows us to write the neutralization reaction as
H3O+ + OH- --> 2 H2O
The Lewis model of electronic structure was one of the factors that prompted two chemists, Bronsted and Lowry, in 1923 to propose a new definition of acids and bases. According to their independently announced definitions an acid can be thought of as a species that can donate a proton and a base is a species that can accept a proton. This definition is more general than the Arrhenius definition; it is not restricted to reactions in water and it expands bases beyond the hydroxide ion to any species that can form a bond to a proton. Because the Lewis model uses a shared pair of electrons to explain a chemical bond, any species that has a nonbonded (lone pair) of electrons can function as a Bronsted-Lowry base. Thus, the reaction of acetic acid with ammonia in the solvent hexane can be written as
This reaction does not involve the formation of the hydroxide ion but does involve the transfer of a proton from acetic acid to the lone pair of electrons on ammonia, thereby forming the ammonium ion, NH4+. Indeed, Bronsted-Lowry reactions are sometimes referred to as proton-transfer reactions to emphasize this aspect of the definition. Notice that Bronsted-Lowry acids are the same as Arrhenius acids because all Arrhenius acids must also be able to form a hydrogen ion.
In the same year that Bronsted and Lowry proposed their definition of acids and bases, Lewis, building upon his new understanding of the nature of the chemical bond, proposed the definition that an acid is an electron pair acceptor while a base is an electron pair donor. This definition is more general than the Bronsted-Lowry definition and builds upon the electron pair as the major instrument of bond formation. A simple reaction that illustrates this new definition is the reaction of a proton with water.
Here, the proton, due to its charge and the absence of electrons around the proton, is capable of sharing a pair of electron. The water has two pairs of electrons to share and the hydronium ion--the product of the reaction--contains a new covalent bond. The product of this sharing of a pair of electrons is called an adduct (other terms such as "complex" are also used for this product). Because the sharing of electrons are involved in all reactions that involve covalent molecules or ions, it should come as no surprise to find that this new definition is one of the most important in chemistry. We also use the term electron-sharing reactions for Lewis acid-base reaction.
Let’s take another look at the reaction of acetic acid with water. Arrhenius thought of acetic acid as a source of the hydrogen ion; Bronsted and Lowry thought of acetic acid as a species capable of donating a proton to a proton acceptor, in this case water. Lewis thought of acetic acid as a species that could accept a pair of electrons from water to form a new covalent bond that appears in the hydronium ion. Later we will talk about how to visualize this interaction of an electron pair at the acetic acid, but for the moment we are mostly interested in the definitions. We must also be aware that the definitions are progressively more general. Lewis acids include the hydrogen ion, but they also include many more species of many different types. Lewis bases are the same as Bronsted-Lowry bases, because both must contain a non-bonded pair of electrons. Bronsted-Lowry bases include, but are not limited to the hydroxide ion, as is the case in the Arrhenius definition. This change in definitions is illustrated in the Table below:
Definition Acid Base
Arrhenius species that produces H+ species that produces OH-
Bronsted-Lowry species that donates H+ species that accepts H+
Lewis species that accepts lone pair species that donates lone pair
About the same time that Lewis proposed his acid-base definition, a Russian chemist Usanovich proposed an even more general definition of acids and bases. According to him, an acid donates a positive species or accepts a negative one, while a base does the opposite. This definition is applicable even to oxidation-reduction reactions, but is consequently so general that most chemists do not find it useful.