The Bronsted-Lowry Definition

During the early twentieth century, scientists J.N. Bronsted and T.M. Lowry put forth their theories. According to them, an acid is a proton donor, and a base is a proton acceptor. 

Consider the ionization reaction of hydrogen iodide. Since hydrogen iodide is a strong acid, it completely dissociates. In this ionization reaction, note that the water molecule acts as the base. 

THIS REACTION IS AN EXAMPLE THAT FITS THE BRONSTED- LOWRY DEFINITION 

In this ionization reaction, obviously HI is the acid. HI transfers or donates its proton to water. Thus, water is the base in this reaction because it accepts the proton. The conjugate base of HI is I. 

Although the Arrhenius definition of acids and bases works in many cases, it cannot easily explain why some substances act as bases even though they do not contain OH . The Arrhenius definition also does not apply to nonaqueous solvents. A second definition of acids and bases, called the Bronsted-Lowry definition, introduced in 1923, applies to a wider range of acid-base phenomena. This definition focuses on the transfer of H ions in an acid-base reaction. Since an H ion is a proton—a hydrogen atom with its electron taken away—this definition focuses on the idea of a proton donor and a proton acceptor. 

According to this definition, HCl is a Bronsted-Lowry acid because, in solution, it donates a proton to water. 

Johannes Bronsted, working in Denmark, and Thomas Lowry, working in England, developed the concept of proton transfer in aoid-base behavior independently and simultaneously. 

This definition more clearly accounts for what happens to the H ion from an acid it associates with a water molecule to form (a hydronium ion). The 

Bronsted-Lowry definition also works well with bases (such as NH3) that do not inherently contain OH ions but that still produce OH ions in solution. NH3 is a Bronsted-Lowry base because it accepts a proton from water. 

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A useful definition of acids and bases is that independently introduced by Johannes Bronsted (1879-1947) and Thomas Lowry (1874-1936) in 1923. In the Bronsted-Lowry definition, acids are proton donors, and bases are proton acceptors. Note that these definitions are interrelated. Defining a base as a proton acceptor means an acid must be available to provide the proton. For example, in reaction 6.7 acetic acid, CH3COOH, donates a proton to ammonia, NH3, which serves as the base. 

The Lewis definition of acids and bases is broader and more encompassing than the Bronsted-Lowry definition because it s not limited to substances that donate or accept just protons. A Lewis acid is a substance that accepts an electron pair, and a Lewis base is a substance that donates an electron pair. The donated electron pair is shared between the acid and the base in a covalent bond. 

The Lewis definition of a base as a compound with a pair of nonbonding electrons that it can use to bond to a Lewis acid is similar to the Bronsted-Lowry definition. Thus, H20, with its two pairs of nonbonding electrons on oxygen, acts as a Lewis base by donating an electron pair to an H+ in forming the hydronium ion, H30+. 

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 ... 

We call such substances Bronsted acids and bases or just plain acids and bases because the Bronsted-Lowry definition is the one commonly accepted today and the one used throughout this text. 

In an acid-base reaction, a proton (H ) is transferred from one chemicai species to another. A species that donates a proton is an acid, and a species that accepts a proton is a base. This identification of acids and bases is the Bronsted-Lowry definition of acid-base reactions. From this perspective, every acid-base reaction has two reactants, an acid and a base. Every acid-base reaction aiso forms two products ... 

Water as the solvent is essential for the acid-base setting reaction to occur. Indeed, as was shown in Chapter 2, our very understanding of the terms acid and base at least as established by the Bronsted-Lowry definition, requires that water be the medium of reaction. Water is needed so that the acids may dissociate, in principle to yield protons, thereby enabling the property of acidity to be manifested. The polarity of water enables the various metal ions to enter the liquid phase and thus react. The solubility and extent of hydration of the various species change as the reaction proceeds, and these changes contribute to the setting of the cement. 

The Bronsted-Lowry definition of an acid is essentially the same as Arrhenius idea An acid is any substance that releases a hydrogen ion. Their idea has come to be known as the Bronsted-Lowry theory of acids and bases. 

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,... 

The chlorine ion can now accept a proton (and become hydrochloric acid again). If the chlorine can accept a proton, according to the Bronsted-Lowry definition, it is a base. Chemists actually call this chlorine ion the conjugate base of hydrochloric acid. Any time an acid gives up its proton, the substance that is left over can act as a base. So every acid has a conjugate base. 

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

Lewis defined a base as an electron pair donor and an acid as an electron pair acceptor. Lewis electron pair donor was the same as Bronsted-Lowry s proton acceptor, and therefore, was an equivalent way of defining a base. Lewis acids were defined as a substance with an empty valence shell that could accommodate a pair of electrons. This definition broadened the Bronsted-Lowry definition of an acid. The three definitions of acids and bases are summarized in Table 13.3. 

Under the Bronsted-Lowry definition, an acid is a substance that donates a hydrogen ion (H+) in an acid-base reaction, while a base is a substance that accepts that hydrogen ion from the acid. When ionized to form a hydrogen cation, hydrogen loses its one and only electron and is left with only a single proton. For this reason, Bronsted-Lowry acids are often called proton donors, and Bronsted-Lowry bases are called proton acceptors. 

This is a simple double replacement reaction (see Chapter 8 for an introduction to these types of reactions). A hydrogen ion from water switches places with the sodium of sodium carbonate to form the products carbonic acid and sodium hydroxide. By the Bronsted-Lowry definition, water is the acid because it donates its hydrogen to Na2COj. This makes Na2C03 the base because it accepts the hydrogen from H2O. 

All Br0nsted-Lowry acids are Lewis acids, but in practice, the term Lewis acid is generally reserved for Lewis acids that don t also fit the Bronsted-Lowry definition. The best way to spot a Lewis acid-base pair is to draw a Lewis dot structure of the reacting substances, noting the presence of lone pairs of electrons. (We introduce Lewis structures in Chapter 5.) For example, consider the reaction between ammonia (NH3) and boron trifluoride (BFj) ... 

Drugs cross biological membranes most readily in the unionised state. The unionised drug is 1000-10000 times more lipid-soluble than the ionised form and thus is able to penetrate the cell membrane more easily. Chemical compounds in solution are acids, bases or neutral. The Bronsted-Lowry definition of an acid is a species that donates protons (H+ ions) while bases are proton acceptors. Strong acids and bases in solution dissociate almost completely into their conjugate base and H+. Weak acids and weak bases do not completely dissociate in solution, and exist in both ionised and unionised states. Most drugs are either weak acids or weak bases. For an acid, dissociation in solution is represented by ... 

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

Which definition of acids and bases is more universal the Bronsted-Lowry definition or the Lewis definition ... 

Whenever we refer simply to acids and bases in this text, we use the Bronsted-Lowry definitions. 

The proton is fundamental to both the Arrhenius and the Bronsted-Lowry definitions of an acid. Dissociation of an Arrhenius acid HA gives an aqueous hydrogen ion, or hydrated proton, written as H + (aq) ... 

Since all proton acceptors have an unshared pair of electrons, and since all electron-pair donors can accept a proton, the Lewis and the Bronsted-Lowry definitions of a base are simply different ways of looking at the same property. All Lewis bases are Bronsted-Lowry bases, and all Bronsted-Lowry bases are Lewis bases. The Lewis definition of an acid, however, is considerably more general than the Bronsted-Lowry definition. Lewis acids include not only H+ but also other cations and neutral molecules having vacant valence orbitals that can accept a share in a pair of electrons donated by a Lewis base. 

For most aqueous acid-base chemistry, the Lewis definitions are too general and lack the symmetry of the acid-conjugate base relationship. We will mostly use the Bronsted-Lowry definitions. 

The Bronsted-Lowry definition of an acid takes into account the nature of the solvent. Although water does not ionize well, it does ionize to a small extent. The result is the appearance of H+ and OH- ions in an equilibrium equation, which is... 

According to the Bronsted-Lowry definitions, and as implied in the previous example, an acid is any moiety that will donate a proton while a base is one that will accept a proton from another... 

Acid-base equilibrium — Using the Bronsted-Lowry definition (see -> acid-base theories), an acid-base reaction involves a -> proton transfer from an acid to a base. Removal of a proton from an acid forms its conjugate base, while addition of a proton to a base forms its conjugate acid. Acid-base equilibrium is achieved when the -> activity (or -> concentration) of each conjugate... 

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. 

E While the Arrhenius definition of an acid says that an acid yields hydro-nium ions as the only positive ions in solution, the Bronsted-Lowry definition says that acids are proton donors. 

C The Lewis definition of acids states that acids are electron pair acceptors while bases are electron pair donors. Choices A, D, and E show the Arrhenius definition whereas choice B shows the Bronsted-Lowry definition. 

According to the Bronsted-Lowry definition, acids and bases are proton donators and -acceptors, respectively, as expressed in the following equilibrium... 

The first person to recognize the essential nature of acids and bases was Svante Arrhenius. Based on his experiments with electrolytes, Arrhenius postulated that acids produce hydrogen ions in aqueous solution, and bases produce hydroxide ions. At the time of its discovery the Arrhenius concept of acids and bases was a major step forward in quantifying acid—base chemistry, but this concept is limited because it applies only to aqueous solutions and allows for only one kind of base—the hydroxide ion. A more general definition of acids and bases was suggested independently by the Danish chemist Johannes N. Bronsted (1879-1947) and the English chemist Thomas M. Lowry (1874-1936) in 1923. In terms of the Bronsted—Lowry definition, an acid is a proton (H+) donor, and a base is a proton acceptor. For example, when gaseous HCl dissolves in water, each HCl molecule donates a proton to a water molecule, and so HCl qualifies as a Bronsted-Lowry acid. The molecule that accepts the proton—water in this case—is a Bronsted-Lowry base. 

The Bronsted-Lowry definition is not limited to aqueous solutions it can be extended to reactions in the gas phase. For example, consider the reaction between gaseous hydrogen chloride and ammonia that we discussed when we studied diffusion (Chapter 5) ... 

According to the Arrhenius concept, a base is a substance that produces OH-ions in aqueous solution. According to the Bronsted-Lowry definition, a base is a proton acceptor. The bases sodium hydroxide (NaOH) and potassium hydroxide (KOH) fulfill both criteria. They contain OH- ions in the solid lattice and behave as strong electrolytes, dissociating completely when dissolving in water ... 

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

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. 

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. 

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