Other Definitions of Acids

When equations of the type (8) and (9) are applied to reactions in nonaqueous solution, qualitative rather than quantitative relationships are usually found. This can be attributed to the difficulties in unequivocally assigning catalytic constants on the basis of the experimental data. Bell has reviewed this problem and the problem of the acid-strength scale to be used (Bell, 43). In most cases the dissociation constants in water are used as the basis of acid strength, and a general discussion of this problem will follow the presentation of the other definitions of acids. 

Almost simultaneously with the papers of Brpnsted (26) and Lowry (25) on proton acids, another viewpoint on acids was formulated by G. N. Lewis (44). Lewis defined a basic molecule as one that has an elec- 

It should be pointed out that this definition includes the proton donor and 

The subsequent dissociation of the complex to hydrogen ions and chloride ions may have been overemphasized in the past. In 1938 Lewis (46) reformulated his ideas on acids and bases in terms of chemical behavior. The important traits were listed as follows  

When an acid or base can combine, the process of combination or neutralization is a rapid one. 

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Other definitions of acids and bases are useful, the most notable being those of Lewis, also proposed in 1923. However, for this chapter, the Brpnsted definition is entirely adequate. 

Other definitions of acid and base are presented later in this section and again in Chapter 18, along with a fuller meaning of neutralization.)... 

Other definitions of acids and bases are discussed in Chapter 16. 

There are other definitions of acids and other measures of acidity. One such definition of acidity is ch,o, which represented a milestone when first proposed by Arrhenius in 1886. However, with the advent of studies in strong acid systems, this definition lost favor because it failed to parallel protonating ability. For example, as the concentration of H2SO4 rises from 0 to 85%, both Chjo" and the ability to protonate bases continue to rise. 

It was G. N. Lewis who extended the definitions of acids and bases still further, the underlying concept being derived from the electronic theory of valence. It provided a much broader definition of acids and bases than that provided by the Lowry-Bronsted concept, as it furnished explanations not in terms of ionic reactions but in terms of bond formation. According to this theory, an acid is any species that is capable of accepting a pair of electrons to establish a coordinate bond, whilst a base is any species capable of donating a pair of electrons to form such a coordinate bond. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. These definitions of acids and bases fit the Lowry-Bronsted and Arrhenius theories, and cover many other substances which could not be classified as acids or bases in terms of proton transfer. 

According to the Arrhenius theory of acids and bases, the acidic species in water is the solvated proton (which we write as H30+). This shows that the acidic species is the cation characteristic of the solvent. In water, the basic species is the anion characteristic of the solvent, OH-. By extending the Arrhenius definitions of acid and base to liquid ammonia, it becomes apparent from Eq. (10.3) that the acidic species is NH4+ and the basic species is Nl I,. It is apparent that any substance that leads to an increase in the concentration of NH4+ is an acid in liquid ammonia. A substance that leads to an increase in concentration of NH2- is a base in liquid ammonia. For other solvents, autoionization (if it occurs) leads to different ions, but in each case presumed ionization leads to a cation and an anion. Generalization of the nature of the acidic and basic species leads to the idea that in a solvent, the cation characteristic of the solvent is the acidic species and the anion characteristic of the solvent is the basic species. This is known as the solvent concept. Neutralization can be considered as the reaction of the cation and anion from the solvent. For example, the cation and anion react to produce unionized solvent ... 

The first clear definition of acidity can be attributed to Arrhenius, who between 1880 and 1890 elaborated the theory of ionic dissociation in water to explain the variation in strength of different acids.3 Based on electrolytic experiments such as conductance measurements, he defined acids as substances that dissociate in water and yield the hydrogen ion whereas bases dissociate to yield hydroxide ions. In 1923, J. N. Brpnsted generalized this concept to other solvents.4 He defined an acid as a species that can donate a proton and defined a base as a species that can accept it. This... 

These ideas were rather limiting since they only applied to aqueous solutions. There were situations where acid-base reactions were taking place in solvents other than water, or even in no solvent at all. This problem was addressed in 1923 by the Danish chemist Johannes Bronsted (1879-1947) and the English chemist Thomas Lowry (1874-1936) when they independentiy proposed a more general definition of acids and bases, and the study of acids and bases took a great step forward. This theory became known as the Bronsted-Lowry theory of acids and bases. 

Lewis Definition of Acids and Bases An even more generalised theory of acids and bases was put forward by Lewis. According to Lewis definition an acid is a substance that can accept an electron pair to form a covalent bond and base is a substance that can furnish an electron pair to form a covalent bond. In other words, an acid is an electron pair Inorganic Chemistry... 

The Br0nsted-Lowry definition of acids and bases depends on the transfer of a proton from the acid to the base. The base uses a pair of nonbonding electrons to form a bond to the proton. G. N. Lewis reasoned that this kind of reaction does not need a proton. Instead, a base could use its lone pair of electrons to bond to some other electron-deficient atom. In effect, we can look at an acid-base reaction from the viewpoint of the bonds that are formed and broken rather than a proton that is transferred. The following reaction shows the proton transfer, with emphasis on the bonds being broken and formed. Organic chemists routinely use curved arrows to show the movement of the participating electrons. 

Seeing that the Brpnsted definitions of acids and bases are not related to a specific solvent, this theory can readily explain the reaction shown in Eq. (5.4). In that case, HCI donates a proton to NH3, resulting in the formation of the ionic salt NH4C1. Therefore, HCI is an acid. Because NH3 accepts a proton, it is acting as a base. Likewise, the Brpnsted theory is applicable to many reactions in which there is a solvent other than water, which makes the Brpnsted theory much more generally applicable than the Arrhenius theory. 

The Brpnsted theory expands the definitions of acids and bases to explain the acidity or basicity of solutions previously regarded as salts chemskiii Builder 18.1 and to explain reactions in other than aqueous solutions. 

This tendency of water molecules to break apart other molecules is part of the first definition of acids and bases, known as the Arrhenius theory of acids and hases and named after (surprise ) a chemist named Svante Arrhenius. According to Arrhenius, an acid is a substance that increases the concentration of H-i- ions in an aqueous (water) solution. So, hydrogen chloride (HCl) dissolved in water is an acid because the water breaks this molecule up into ions and CT ions. Actually, this statement is a bit of a lie because ions (which are simply protons—take away an electron from hydrogen, which is com-... 

Ironically, until 1953, Nazarov incorrectly described the mechanism of the general transformation which now bears his name. In 1952, Braude and Coles were the first to suggest the intermediacy of car-bocations and demonstrated that the formation of 2-cyclopentenones actually proceeds via the a,a -divi-nyl ketones (equation 1). This fact together with further mechanistic clarification, has led to the specific definition of the Nazarov cyclization as the acid-catalyzed closure of divinyl ketones to 2-cyclopentenones. This process was already documented in 1903 by Vorliinder who isolated a ketol of unknown structure by treatment of dibenzylideneacetone with concentrated sulfuric acid and acetic acid followed by mild alkaline hydrolysis (equation 2). The correct structure of Vorliinder s ketol, finally proposed in 1955, ° arises from Nazarov cyclization followed by oxidation and isomerization. Other examples of acid-catalyzed cyclizations of divinyl ketones exist in the early literature. ... 

In 1923 G. N. Lewis" proposed a definition of acid-base behavior in terms of electron-pair donation and acceptance. The Lewis definition is perhaps the most widely used of all because of its simplicity and wide applicability, espeaally in the field of organic reactions. Lewis defined a base as cn electron-pair donor and an acid as an electron-pair acceptor. In addition to all of llie reactions discussed above, the Lewis definition includes reactions in which no ions are formed and no hydrogen ions or other ions are transferred 2... 

The fact that a Lewis acid must be able to accept an electron pair means that it must have either a vacant, low-energy orbital or a polar bond to hydrogen so it can donate H" which has an empty Is orbital). Thus, the Lewis definition of acidity is much broader than the Bronsted-Lowry definition and includes many other species in addition to H. For example, various metal cations such as are Lewis acids because they accept a pair of electrons when they form a bond to a base. In the same way, compounds of group 3A elements such as BF3 and AlCln are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases, as shown in Figure 2.5. Similarly, many transition-metal compounds, such as TiCU, FeCla, ZnCl, and SnCl4, are Lewis acids. 

The first point to be made concerning acids and bases is that so-called acid-base theories are in reality definitions of what an acid or base is they are not theories in the sense of valence bond theory or molecular orbital theory. In a very real sense, we can make an acid be anything we wish the differences between the various acid-base concepts are not concerned with which is right but which is most convenient to use in a particular situation. All of the current definitions of acid-base behavior are compatible with each other. In fact, one of the objects in the following presentation of many different definitions is to emphasize their basic parallelism and hence to direct the students toward a cosmopolitan attitude toward acids and bases which will stand them in good stead in dealing with various chemical situations, whether they be in aqueous solutions of ions, organic reactions, nonaqueotis titrations, or other situations. 

It may surprise you to find that, despite the reason that acids and bases are found in so many of the materials around us, there are several different theories defining what acids and bases are. The real difference between the theories has to do with how broad or limited our definitions of acids and bases should be. For example, should only compounds that ionize to release hydroxide ions (OH ) be considered bases, or should other compounds that neutralize acids without releasing hydroxide ions be included in our definition Examine the following theories. 

The final acid-base theory that we shall consider was proposed by chemist Gilbert Lewis in the early 1920s. The Lewis Theory is the most general, including more substances under its definitions than the other theories of acids and bases. A Lewis acid is a substance that accepts a pair of electrons to form a covalent bond. A Lewis base is a substance that provides a pair of electrons to form a covalent bond. In order for a substance to act as a Lewis base, it must have a pair of unshared electrons in its valence shell. An example of this is seen when a hydrogen ion attaches to the unpaired electrons of oxygen in a water molecule, as shown here ... 

Acids are compounds that ionise to release hydrogen ions, or protons, to their surroundings. Bases are compounds that can accept hydrogen ions. This is called the Bransted-Lowry definition of acids and bases (named after yet another Scandinavian chemist, Johannes Nicolaus Bronsted, and Thomas Martin Lowry, who was British). There are other ways of explaining acidity and basicity, but the Bransted-Lowry theory works most of the time, and will be used throughout this book. 

See pH. When dealing with chemical reactions in solvents other than water, it is sometimes convenient to define an acid as a substance that ionizes to give the positive ion of the solvent. The common definitions of acid have been extended as more detailed studies of chemical reactions have been made. The Lowry-Brpnsted definition of an acid as a substance that can give up a proton is more useful in connection with an understanding of bases (see base). Perhaps the most significant contribution to the theory of acids was the electron-pair concept introduced by G. N. Lewis around 1915. 

Because both dihydrogen phosphate and hydrogen carbonate (and other substances like them) can be either Bronsted-Lowry acids or bases, they cannot be described as a Bronsted-Lowry acid or base except with reference to a specific acid-base reaction. For this reason, the Arrhenius definitions of acids and bases are the ones used to categorize isolated substances on the stockroom shelf A substance generates either hydronium ions, hydroxide ions, or neither when added to water, so it is always either an acid, a base, or neutral in the Arrhenius sense. Hydrogen carbonate is an Arrhenius base because it yields hydroxide ions when added to water. Dihydrogen phosphate is an Arrhenius acid because it generates hydronium ions when added to water. 

In each case, the atom to which the proton becomes attached possesses at least one unshared pair of electrons. This characteristic property of OH, NH3, and other Brpnsted bases suggests a more general definition of acids and bases. 

The significance of the Lewis concept is that it is much more general than other definitions. Lewis acid-base reactions include many reactions that do not involve Brpnsted acids. Consider, for example, the reaction between boron trifluoride (BF3) and ammonia to form an adduct compound (Figure 15.11) ... 

Usually, definitions of acids and bases are necessary for the classification of different kinds of chemical reactions, i.e. for dividing them into acid-base and other ones. For example, the Brpnsted-Lowry definition divides reactions into acid-base, which are characterized by redistribution of protons, and other ones. The wider Lewis definition makes the division of reactions into acid-base and redox ones, meaning that in the former case there is redistribution of electron density on account of electron pairs, and the latter case concerns reactions with the transfer of single electrons. Since chemistry concerns just the redistribution of electrons of external shells, we can classify all chemical reactions as acid-base ones by the Usanovitch definition. 

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