Bronsted Lewis definition

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 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. 

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 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). 

KolthoflF, I. M. (1944). The Lewis and Bronsted-Lowry definitions of acids and bases. Journal of Physical Chemistry, 48, 51-7. 

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. 

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) ... 

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. 

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

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. 

Bronsted-Lowry and Lewis definitions of acids and bases are reviewed in Sec. 7.6. Alcohols are comparable in acidity to water, but phenols are much more acidic. This increased acidity is due to charge delocalization (resonance) in phenoxide ions. Electron-withdrawing groups, such as -F and -N02, increase acidity, through either an inductive or a resonance effect, or both. 

However, many species which are acids under Lewis definition cannot be termed so according to Bronsted definition. A few examples are sulphur trioxide and halides of boron, aluminium, iron (ferric) and zinc. The central atom in each is able to accept a pair of electrons to complete is octet. 

Similarly, in a Bronsted acid like HC1, there is no vacant orbital to accept an electron pair from a base. There is, therefore, no obvious reason to label it as an acid according to Lewis definition. 

According to the Lewis definition, an acid is an electron pair acceptor and a base is an electron pair donor. All Bronsted-Lowry bases are also Lewis bases. However, Lewis acids include many species that are not proton acids instead of H+, they have some other electron-deficient species that acts as the electron pair acceptor. An example of a Lewis acid-base reaction is provided by the following equation. In this reaction the boron of BF3 is electron/deficient (it has only six electrons in its valence shell). The oxygen of the ether is a Lewis base and uses a pair of electrons to form a bond to the boron, thus completing boron s octet. 

BRONSTED-LOWRY AND LEWIS DEFINITIONS OF ACIDS AND BASES... 

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. 

A furtlier important conccpi related to electronegativity and polarity is that of acUity nd biisicity. We ll see, in fact, that much of the chemistry of organic molecules can be explained by their acid-base behavior. may recall from a course in general chemistry that there are two frequently used definitions of acidity the Bivusted-Lowry definition ajid the Lewis definition. We ll look at the Bronsted-l.owry definition in this and tlie next three sections and then discuss the Lewis definition in Section 2.11. 

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. 

The Lewis bonding model with its electron pairs can be used to define a more general kind of acid-base behavior of which the Arrhenius and Bronsted-Lowry definitions are special cases. A Lewis base is any species that donates lone-pair electrons, and a Lewis acid is any species that accepts such electron pairs. The Arrhenius acids and bases considered so far fit this description (with the Lewis acid, H, acting as an acceptor toward various Lewis bases such as NH3 and OH , the electron pair donors). Other reactions that do not involve hydrogen ions can still be considered Lewis acid-base reactions. An example is the reaction between electron-deficient BF3 and electron-rich NH3 ... 

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