Lewis definition

The Lewis definitions of acids and bases provide for a more general view of acid-base reactions than either the Arrhenius or Br0nsted-Lowry pic ture A Lewis acid is an electron pair acceptor A Lewis base is an electron pair donor The Lewis approach incorporates the Br0nsted-Lowry approach as a subcategory m which the atom that accepts the electron pair m the Lewis acid is a proton... 

Achiral (Section 7 1) Opposite of chiral An achiral object is supenmposable on its mirror image Acid According to the Arrhenius definition (Section 1 12) a substance that ionizes in water to produce protons Accord mg to the Br0nsted-Lowry definition (Section 1 13) a sub stance that donates a proton to some other substance According to the Lewis definition (Section 1 17) an electron pair acceptor... 

A further important concept related to electronegativity and polarity is that of acidity and basicity. We ll see, in fact, that much of the chemistry of organic molecules can be explained by their acid-base behavior. You may recall from a course in general chemistry that there are two frequently used definitions of acidity the Brtfnsted-Lowry definition and the Lewis definition. We ll look at the... 

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 fact that a Lewis acid is able to accept an electron pair means that it must have either a vacant, low-energy orbital or a polar bond to hydrogen so that it can donate H+ (which has an empty7 Is orbital). Thus, the Lewis definition of acidity includes many species in addition to H+. For example, various metal cations, such as Mg2+, are Lewis acids because they accept a pair of electrons when they form a bond to a base. We ll also see in later chapters that certain metabolic reactions begin with an acid-base reaction between Mg2+ as a Lewis acid and an organic diphosphate or triphosphate ion as the Lewis base. 

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+. 

In 1923. Lewis published a classic book (later reprinted by Dover Publications) titled Valence and the Structure of Atoms and Molecules. Here, in Lewis s characteristically lucid style, we find many of the basic principles of covalent bonding discussed in this chapter. Included are electron-dot structures, the octet rule, and the concept of electronegativity. Here too is the Lewis definition of acids and bases (Chapter 15). That same year, Lewis published with Merle Randall a text called Thermodynamics and the Free Energy of Chemical Substances. Today, a revised edition of that text is still used in graduate courses in chemistry. 

Another definition of acids and bases is due to G. N. Lewis (1938). From the experimental point of view Lewis regarded all substances which exhibit typical acid-base properties (neutralisation, replacement, effect on indicators, catalysis), irrespective of their chemical nature and mode of action, as acids or bases. He related the properties of acids to the acceptance of electron pairs, and bases as donors of electron pairs, to form covalent bonds regardless of whether protons are involved. On the experimental side Lewis definition brings together a wide range of qualitative phenomena, e.g. solutions of BF3, BC13,... 

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. 

The Lewis definition of a base is broader than the Bronsted definition. That is, although every Bronsted base is a Lewis base, not every Lewis base is a Bronsted base. For instance, carbon monoxide is an important Lewis base in its reactions with metals, but it is not a Bronsted base because it does not accept protons. 

In Chapter H, we introduce a second definition of acids and bases, the Lewis definition, which focuses attention on electron movement rather than proton movement Until then, acid-base always means proton transfer."... 

In both cases the nitrogen atom uses its pair of nonbonding electrons to make a new covalent bond. This similarity led G. N. Lewis to classify ammonia as a base in its reaction with B (CH3)3 as well as in its reaction with H3 O . Whereas the Br< )nsted definition focuses on proton transfer, the Lewis definition of acids and bases focuses on electron pairs. 

Thus an acid-base reaction involves the transfer of an oxide ion (compared with the transfer of a proton in the Bronsted theory) and the theory is particularly applicable in considering acid-base relationships in oxide, silicate and aluminosilicate glasses. However, we shall find that it is subsumed within the Lewis definition. 

Although Lewis and Bronsted bases comprise the same species, the same is not true of their acids. Lewis acids include bare metal cations, while Bronsted-Lowry acids do not. Also, Bell (1973) and Day Selbin (1969) have pointed out that Bronsted or protonic acids fit awkwardly into the Lewis definition. Protonic acids cannot accept an electron pair as is required in the Lewis definition, and a typical Lewis protonic add appears to be an adduct between a base and the add (Luder, 1940 Kolthoff, 1944). Thus, a protonic acid can only be regarded as a Lewis add in the sense that its reaction with a base involves the transient formation of an unstable hydrogen bond adduct. For this reason, advocates of the Lewis theory have sometimes termed protonic adds secondary acids (Bell, 1973). This is an unfortunate term for the traditional adds. 

The Lewis definition covers all AB cements, including the metal oxide/metal oxysalt systems, because the theory recognizes bare cations as aprotic acids. It is also particularly appropriate to the chelate cements, where it is more natural to regard the product of the reaction as a coordination complex rather than a salt. Its disadvantages are that the definition is really too broad and that despite this it accommodates protonic acids only with difficulty. 

It is better than the Lewis theory for describing acid-base cements, for it avoids the awkwardness that the Lewis definition has with protonic acids. However, as Day Selbin (1969) have observed, the generality of the theory is such that it includes nearly all chemical reactions, so that acid-base reactions could simply be termed chemical reactions . 

The various acid-base definitions are summarized in the Venn diagram (Fig. 2.1). From this it can be seen that the Usanovich definition subsumes the Lewis definition, which in turn subsumes all other definitions (i.e. Arrhenius, Bronsted-Lowry, Germann-Cady-Elsey, Lux-Flood). 

B The Lewis Definition of Acids and Bases 1. Lewis acid-base theory in 1923 proposed by G. N. Lewis (1875-1946 Ph. D. Harvard, 1899 professor, Massachusetts Institute of Technology, 1905-1912 professor, University of California, Berkeley, 1912-1946). 

In practice, it s much simpler to use the Arrhenius or Bronsted-Lowry definition of acid and base, but you ll need to use the Lewis definition when hydrogen ions aren t being exchanged. You can pick and choose among the definitions when you re asked to identify the acid and base in a reaction. 

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. 

Bronstctl-Lowry and Lewis definitions of acids and bases were introduced and discussed. 

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

A water molecule, HzO, becomes strongly held to a dissolved aluminum ion, Al3+. According to the Lewis definition of acids and bases, which is behaving... 

The hydride ion is a hydrogen atom with two electrons, which pair together as a lone pair. According to the Lewis definition, the hydride ion behaves as a base. In fact, the hydride ion is a very strong base. It will react with water, pulling a positive hydrogen ion off of the water to create H2 plus hydroxide ions, OH-, as follows ... 

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