Bransted definition

We will consider more general definitions of acids and bases in Chapter 13 (Bransted-Lowry) and Chapter 15 (Lewis). 

A1C13, or S02 in an inert solvent cause colour changes in indicators similar to those produced by hydrochloric acid, and these changes are reversed by bases so that titrations can be carried out. Compounds of the type of BF3 are usually described as Lewis acids or electron acceptors. The Lewis bases (e.g. ammonia, pyridine) are virtually identical with the Bransted-Lowry bases. The great disadvantage of the Lewis definition of acids is that, unlike proton-transfer reactions, it is incapable of general quantitative treatment. 

Reaction 3.2 is a prime example of the use of the terms Lewis acid and Lewis base. G. N. Lewis suggested the usage such that a donor of an electron pair is a base and the acceptor molecule is an acid. The classical Bransted acid and base pair, H+(aqueous) and OH (aqueous), are encompassed by the Lewis definitions as they combine to give water, the hydroxide ion supplying both electrons. 

A Lewis base transfers an electron pair to a Lewis acid. A Bronsted acid transfers a proton to a Bnansted base. These exist in conjugate pairs at equilibrium. In an Arrhenius base, the proton acceptor (electron pair donor) is OH-. All Arrhenius acids/bases are Bronsted acids/bases and all Bransted acids/bases are Lewis acids/bases. Each definition contains a subset of the one that comes after it. 

While the Arrhenius definitions focus on the H+ and OH- ions, the Bransted-Lowry definitions focus on the behavior of protons—that is, the transfer of a proton from one substance to another. 

Earlier in this chapter we considered Arrhenius s concept of acids and bases An acid is a substance that produces H+ ions when dissolved in water, and a base is a substance that produces OH- ions. Although these ideas are fundamentally correct, it is convenient to have a more general definition of a base, which covers substances that do not produce OH- ions. Such a definition was provided by Bransted and Lowry, who defined acids and bases as follows ... 

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. 

Acidity and basicity are fundamental properties of organic compounds, and acid-base reactions are essential steps in many organic transformations. Although there are several definitions of acidity and basicity, the Bransted theory and the Lewis theory are used most often in organic chemistryIn Lewis theory, an acid is an electron pair acceptor and a base is an electron pair donor, as in the reaction of a trialkylamine as Lewis base with boron trifluoride as Lewis acid (equation 7.1). 

In the Bransted-Lowry add-base definition, a base is any species that... 

We then learn the more general Bransted-Loivry definitions for acid and base. A Bronsted-Lowry acid is a proton donor and a Bronsted—Lowry base is a proton acceptor. 

Arrhenius s definitions of acids and bases are hmited in that they apply only to aqueous solutions. Broader definitions were proposed by the Danish chemist Johannes Bransted in 1932 a Brensted acid is a proton donor, and a Brensted base is a proton acceptor. Note that Bronsted s definitions do not require acids and bases to be in aqueous solution. 

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