The Bronsted-Lowry Definition of Acids and Bases

The definition of acids and bases we will use in this book is called the Bronsted-Lowry definition, named for Johannes Bronsted (1879-1947), a Danish chemist, and Thomas Lowry (1874-1936), an English chemist, who put forth definitions of acids and bases independently of each other in 1923. 

In their definition of an acid, the important chemical species is the hydrogen ion (H ), often written as the hydronium ion (H3O ). A hydronium ion is a water molecule, HjO, with an H+ ion attached to the oxygen atom. The reason we often write the H ion as H3O is to acknowledge the fact that most neutral hydrogen atoms consist only of a proton and an electron. To form the H+ ion, the electron is removed, leaving only the proton. Protons are incredibly tiny centers of positive charge that are not going to exist in water as independent chemical species. Rather, they will always attach themselves to water molecules. With bases, the important chemical species is the hydroxide ion (OH ). 

According to Bronsted and Lowry, a molecule or ion is an acid if, when the species is added to water, the concentration of H+ (or equivalently, H3O ) increases. Consider, for example, what happens when hydrochloric acid is added to water  

The formation of H3O+ ions makes the solution acidic. Another example is the reaction of acetic acid (HC HjO ) with water  

Notice a subtle difference between the two equations. In the equation with HCl, the arrow only points to the right. This is because HCl is what we call a strong acid, meaning that when it dissolves in water, the result is a relatively high concentration of H3O+ ions. In fact, essentially 100 percent of the HCl molecules will react, leaving no neutral HCl molecules left in the solution. 

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The problem with the Arrhenius definitions is that they are specific to one particular solvent, water. When chemists studied nonaqueous solvents, such as liquid ammonia, they found that a number of substances showed the same pattern of acid-base behavior, but plainly the Arrhenius definitions could not be used. A major advance in our understanding of what it means to be an acid or a base came in 1923, when two chemists working independently, Thomas Lowry in England and Johannes Bronsted in Denmark, came up with the same idea. Their insight was to realize that the key process responsible for the properties of acids and bases was the transfer of a proton (a hydrogen ion) from one substance to another. The Bronsted-Lowry definition of acids and bases is as follows ... 

According to the Bronsted-Lowry definition of acids and bases, an acid is a proton donor. The particle that is left over after an acid donates its proton, however, can now accept a proton and,... 

Thus, a reducing agent donates electrons, while an oxidizing agent receives them. The Bronsted-Lowry definitions of acid and base specify that... 

What are the Bronsted-Lowry definitions of acid and base ... 

The Br0nsted-Lowry definition of acids and bases does not replace the Arrhenius definition, but extends it. The Bronsted-Lowry definition of acids and bases requires you to take a closer look at the reactants and products of an acid-base reaction. In this case, acids and bases are not easily defined as having hydronium and hydroxide ions. Instead, you are asked to look and see which substance has lost a proton and which has gained the very same proton that was lost. 

Compare the Bronsted-Lowry definitions of acids and bases with the Arrhenius definitions of acids and bases. 

Chapters 10 and 11 describe the special properties of liquid water. Because of its substantial dipole moment, water is especially effective as a solvent, stabilizing both polar and ionic solutes. Water is not only the solvent, but also participates in acid-base reactions as a reactant. Water plays an integral role in virtually all biochemical reactions essential to the survival of living organisms these reactions involve acids, bases, and ionic species. In view of the wide-ranging importance of these reactions, we devote the remainder of this chapter to acid-base behavior and related ionic reactions in aqueous solution. The Bronsted-Lowry definition of acids and bases is especially well suited to describe these reactions. 

According to the Bronsted-Lowry definition of acids and bases, what is the conjugate acid of ammonia ... 

You can use the ability to exchange a hydrogen ion as the basis of a broader definition of an acid or a base. In this definition, called the Bronsted-Lowry definition of acids and bases, an acid is defined as a substance that donates, or gives up, a hydrogen ion in a chemical reaction. A base, not surprisingly, is just the opposite. A base is a substance that accepts a hydrogen ion in a chemical reaction. 

There are various definitions of acids and bases. The one used here is attributed to a theory developed in 1923 independently by Johannes Bronsted (1879-1947), a Danish chemist, and Thomas Lowry (1874-1936), a British chemist. Recall that an atom of ordinary hydrogen has only a proton and an electron, and no neutrons. Therefore, a cation of ordinary hydrogen (H+) is just a proton. In the Bronsted-Lowry definition of acids and bases, an acid is a proton donor, that is, it can react with other compounds or ions by transferring one or more ions to the other compounds or ions. A base is a proton acceptor It can react with the H+ ions of compounds or ions that are acids. Some chemical species, such as H O, are said to be amphiprotic, that is, they are both donors and acceptors of protons they are both an acid and a base. These definitions are illustrated in the following examples ( aq means the reaction is taking place in aqueous solution) ... 

Based on the Bronsted-Lowry definition of acids and bases, an acid is a proton donor and base is a proton acceptor. A species (charged or uncharged) that can gain or lose a proton is called amphoteric or amphiprotic species. For example, (bicarbonate ion) HCOj" can donate a proton acting as an acid. The same species can accept a proton acting as a base. Hence, amphoteric species can act as an acid or a base depending on the surrounding conditions. 

Our emphasis throughout this chapter has been on water as the solvent and on the proton as the source of acidic properties. In such cases we find the Bronsted—Lowry definition of acids and bases to be the most useful. In fact, when we speak of a substance as being acidic or basic, we are usually thinking of aqueous solutions and using these terms in the Arrhenius or Bronsted—Lowry sense. The advantage of the Lewis definitions of acid and base is that they allow us to treat a wider variety of reactions, including those that do not involve proton transfer, as acid—base reactions. To avoid confusion, a substance such as BF3 is rarely called an acid unless it is clear from the context that we are using the term in the sense of the Lewis definition. Instead, substances that function as electron-pair acceptors are referred to explicitly as Lewis acids. ... 

Notice that the reaction in Equation 16.3 involves a proton donor (HCl) and proton acceptor (H2O). The notion of transfer from a proton donor to a proton acceptor is the key idea in the Bronsted-Lowry definition of acids and bases ... 

You have reviewed the Bronsted-Lowry definition of acids and bases and the meanings of pH and pTQ. You have learned to identify the most acidic hydrogen atoms in a molecule based on a comparison of pIQ values. You will see in many cases that Brensted—Lowry acid-base reactions either initiate or complete an organic reaction, or prepare an organic molecule for further reaction. The Lewis definition of acids and bases may have been new to you. However, you will see over and over again that Lewis acid—base reactions which involve either the donation of an electron pair to form a new covalent bond or the departure of an electron pair to break a covalent bond are central steps in many organic reactions. The vast majority of organic reactions you will study are either Bronsted-Lowry or Lewis acid—base reactions. 

Apart from this modification, the Arrhenius definitions of acid and base are still valid and useful today, as long as we are talking about aqueous solutions. However, the Arrhenius concept of acids and bases is so intimately tied to reactions that take place in water that it has no good way to deal with acid-base reactions in nonaqueous solutions. For this reason, we concentrate in this chapter on the Bronsted-Lowry definitions of acids and bases, which are more useful to us in our discussion of reactions of organic compounds. 

In 1923, two chemists, Johannes Bronsted and Thomas Lowry, described acids and bases in the scientific literature. They were studying how the transfer of hydrogen ions (protons) took place between reacting molecules. A Bronsted-Lowry acid donates a proton in a reaction, while a base is on the receiving end of the proton transfer. In the Bronsted-Lowry definition of acids and bases, ions as well as larger more complex molecules are included. The pairs of (H2O, OH ) and (NH3, NH4) are called Bronsted-Lowry conjugate acid-base pairs. 

Let s consider another example that compares the relationship between the Arrhenius definitions and the Bronsted-Lowry definitions of acids and bases— an aqueous solution of ammonia, in which fhe following equilibrium occurs ... 

The Bronsted-Lowry definitions of acids and bases provide a basis for studying proton-transfer reactions. Suppose that a Bronsted-Lowry acid gives up a proton the remaining ion or molecule can re-accept that proton and can act as a base. Such a base is known as a conjugate base. Thus, the species that remains after a Brensted-Lowry acid has given up a proton is the conjugate base of that acid. For example, the fluoride ion is the conjugate base of hydrofluoric acid. 

Before leaving a discussion of the properties of water and aqueous solutions, we should discuss its self-ionization. Recall that the Bronsted-Lowry definition of acids and bases is in terms of proton donors and acceptors, respectively. Water is an amphoteric substance that is, it can act as either an acid or a base as shown in Equation (11.7) ... 

Johannes Nicolas Bronsted, 1879-1947, was a Danish physical chemist who made various contributions, Including the Bronsted-Lowry definition of acids and bases. He was elected to the Danish parliament In 1947, but died before taking office. 

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