Acids and Bases The Br0nsted-Lowry Definition

Still another important concept related to electronegativity and polarity is that of acidity and basicity. We ll soon see that the acid-base behavior of organic molecules helps explain much of their chemistry. You may recall from a course in general chemistry that there are two frequently used definitions of acidity, the Br0nsted-Lowry definition and the Lewis definition. 

We ll look at the Br0nsted Lowry definition in this and the next three sections, and then discuss the Lewis definition in Section 2.11. 

A Br0n ted-Lowry acid is a substance that donates a hydrogen ion (H ), and a Br0nsted-Lowry base is a substance that accepts (The name proton is often used as a synonym for because loss of the valence 

Hydronium ion, the product that results when the base H2O gains a proton, is called the conjugate acid of the base chloride ion, the product that results when the acid HCl loses a proton, is called the conjugate base of the acid. Other common mineral acids such as H2SO4 and HNO3 behave similarly, as do organic acids such as acetic acid, CH3COOH. 

Note that water can act either as an acid or as a base, depending on the circumstances. In its reaction with HCl, water is a base that accepts a 

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We will look at two definitions for the terms acid and base, the Br0nsted-Lowry definitions and the Lewis definitions. 

According to the Br0nsted-Lowry definitions, any species that contains hydrogen can potentially act as an acid, and any compound that contains a lone pair of electrons can act as a base. Therefore, neutral molecules can also act as bases if they contain an oxygen, nitrogen or sulphur atom. Both an acid and a base must be present in a proton transfer reaction, because an acid cannot donate a proton unless a base is present to accept it. Thus, proton-transfer reactions are often called acid-base reactions. 

This is a very useful relationship. You should practice writing equations according to the Br0nsted-Lowry definitions of acids and bases and familiarize yourself with Table 1.7 which gives the pXa s of various Brpn-sted acids. 

When a base accepts a proton, it becomes an acid capable of returning that proton. When an acid donates its proton, it becomes a base capable of accepting that proton back. One of the most important principles of the Br0nsted-Lowry definition is this concept of conjugate acids and bases. For example, NH4 and NH3 are a conjugate acid-base pair. NH3 is the base when it accepts a proton, it is transformed into its conjugate acid, NH4. Many compounds (water, for instance) can react either as an acid or as a base. Here are some additional examples of conjugate acid-base pairs. 

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. 

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. 

The Lewis definition of acids and bases is more general than the Br0nsted-Lowry definition. 

The Br0nsted-Lowry definition of acidity discussed in the previous four sections encompasses all compounds containing hydrogen. Of even more use, however, is the Lewis definition of acids and bases, which is not limited to compounds that gain or lose protons. A Lewis acid is a substance that accepts an electron pair, and a Lewis base is a substance that donates an... 

In contrast to the Br0nsted-Lowry theory, which emphasizes the proton as the principal species in acid-ba.se reactions, the definition proposed by Lux and extended by Floods describes acid-base behavior in terms of the oxide ion. This acid-base concept was advanced to treat nonprotonic systems which were not amenable to the Br0nsted-Lowry definition. For example, in high-temperature inorganic melts, reactions such as the following take place ... 

Before continuing on to the last definition of acids and bases, it will be helpful to consider the definitions for strong and weak acids within the context of the Br0nsted-Lowry model of acids and bases. The definitions are really an extension of the Arrhenius ideas. In the Arrhenius definitions, strong acids and bases were those that ionize completely. Most Brpnsted-Lo wry acids and bases do not completely ionize in solution, so the strengths are determined based on the degree of ionization in solution. For example, acetic acid, found in vinegar, is a weak acid that is only about 1 % ionized in solution. That means that when acetic acid, 11C, 11,(),. is placed in water, the reaction looks like... 

The Br0nsted-Lowry definitions of acids and bases are widely used in organic chemistry. As noted in the preceding equation, the conjugate acid of a substance is formed when it accepts a proton from a suitable donor. Conversely, the proton donor is converted to its conjugate base. A conjugate acid-base pair always differ by a single proton. 

The Br0nsted-Lowry definition provides a new way to look at acid-base reactions because it focuses on the reactants and the products. For example, let s examine the reaction between hydrogen sulfide and ammonia ... 

The definitions of acid and base that we use now were provided by Brpnsted and Lowry in 1923. In the Br0nsted-Lowry definitions, an acid is a species that donates a proton, and a base is a species that accepts a proton. (Remember that positively charged hydrogen ions are also called protons.) In the following reaction, hydrogen chloride (HCl) meets the Brpnsted-Lowry definition of an acid because it donates a proton to water. Water meets the definition of a base because it accepts a proton from HCl. Water can accept a proton because it has two lone pairs. Either lone pair can form a covalent bond with a proton. In the reverse reaction, H3O is an acid because it donates a proton to CF, and CF is a base because it accepts a proton from H30. ... 

Lewis defined a base as an electron-pair donor and an acid as an electron-pair acceptor. Modem inorganic chemistry extensively uses the Lewis definition, which encompasses the Br0nsted-Lowry definition, since accepts an electron pair from a Brpnsted base during protonation. The Lewis definition dramatically expands the acid list to include metal ions and main group compounds, and provides a framework for nonaqueous reactions. The Lewis definition includes reactions such as... 

Section 1.14 According to the Br0nsted-Lowry definitions, an acid is a proton donor and a base is a proton acceptor. 

Understand the Br0nsted-Lowry definitions of an add and a base discuss how water can act as a base or as an acid and how an add-base reaction is a proton-transfer process involving two conjugate add-base pairs, with the stronger acid and base forming the weaker base and add ( 18.3) (SPs 18.4,18.5) (EPs 18.24-18.39)... 

This analysis relies on the Br0nsted-Lowry definitions of an acid and a base as used in general chemistry, as well as an extension of the fundamental ionization equations to solvents other than water. What is the pointi Water is not used for many reactions found in organic chemistry, and it is important to change the focus from acids that ionize in water to the concept that acids react with bases in order to ionize. This concept is the basis for many of the reactions that will be introduced in succeeding chapters. For the moment, it is important to remember that both the acid and the base must be identified in a reaction based on their reactivity. 

A Br0nsted-Lowry acid is defined as a proton donor, and a Br0nsted-Lowry base is a proton acceptor. A Lewis acid accepts an electron pair from a Lewis base, which is an electron pair donor. There may be confusion about how these definitions are related rather than about just how they differ. Part of the problem may lie with removing the base (water) from acid-base reactions done in water. When HCl reacts with water, the Br0nsted-Lowry definition states that the proton (H+)... 

The Br0nsted-Lowry definition states that an acid is a proton donor (Chapter 2, Section 2.1), but the proton does not fly off. In fact, the acidic proton in A H is pulled off by the base, and this reaction leads to cleavage of the covalent bond between A and H with transfer of those two electrons from the A-H bond to A. This reaction generates the electron-rich A (the conjugate base). Therefore, the acid-base reaction shown is simply a chemical reaction in which the electron-rich base donates two electrons to the electron-poor proton, forming a new covalent bond and breaking the covalent bond between A-H with transfer of those two electrons to A. 

The Lewis definition, like the Br0nsted-Lowry definition, requires that a base have an electron pair to donate, so it does not expand the classes of bases. However, this definition greatly expands the classes of acids. Many species, such as CO2 and Cu, that do not contain H in their formula (and thus cannot be Brpnsted-Lowry acids) are Lewis acids because they accept an electron pair in reactions. Thus, the proton itself is a Lewis acid because it accepts the electron pair donated by a base ... 

The Br0nsted-Lowry definition of acids and bases broadened the definition of acids and bases. Acids are proton donors and bases are proton acceptors. This is in agreement with Arrhenius theory but includes reactions such as that between ammonia and hydrogen chloride gases ... 

To avoid confusion between the Lewis and the Br0nsted-Lowry definitions of acids and bases, Lewis bases are sometimes called nucleophiles, and Lewis acids are called electrophiles. In the example above, water acts as a nucleophile (donates electrons), and the carbocation acts as an electrophile (receives electrons). 

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