Acids and Lewis Bases

A Lewis acid (electrophile) shares an electron pair furnished by a Lewis base (nucleophile) to form a covalent (coordinate) bond. The Lewis concept is especially useful in explaining the acidity of an aprotic acid (no available proton), such as BFj. 

Problem 3.33 Explain the observation that neither pure H,SO, nor pure HCIO4 conducts an electric current but a mixture of the two does.  

Neither pure acid is ionized. In the mixture the stronger acid, HCIO, donates a proton to H2SO4, which acts as a base because of the unshared electron pairs on the O s. 

Problem 3.34 Given the following Lewis acid-base reactions  

Problem 3.35 Which of the following reactions can take place with carbocations Give examples when reactions do occur, (a) acts as an acid (b) reacts as an electrophile (c) reacts as a nucleophile (d) undergoes rearrangement.  

The general structure of amino acids, the building blocks of proteins, is 

Amine group Carboxyl group (basic) (acidic) 

Amino acids contain a carboxyl group and can therefore serve as acids. They also contain an NH2 group, characteristic of amines (Section 16.7), and thus they can also act as bases. Amino acids, therefore, are amphiprotic. For glycine, we might expect the acid and base reactions with water to be 

Although the form of the amino add on the right in this equation is electrically neutral overall, il has a positively charged end and a negatively charged end. A molecule of this type is called a zwitterion (German for hybrid ion ). 

Do amino acids exhibit any properties indicating that they behave as zwitterions If so, their behavior should be similar to that of ionic substances. (Section 8.2) Crystalline amino adds have relatively high melting points, usually above 200 C, which is characteristic of ionic solids. Amino adds are far more soluble in vrater than in nonpolar solvents. In addition, the dipole moments of amino acids are large, consistent with a laige separation of charge in the molecule. Thus, the ability of amino adds to ad simultaneously as adds and bases has important effeds on their properties. 

In Lewis acid-base theory, proton donors are not the only acids many other species are acids as well. Aluminum chloride, for example, reacts with ammonia in the same way that a proton donor does. Using curved arrows to show the donation of the electron pair of ammonia (the Lewis base), we have the following examples  

Bases are much the same in the Lewis theory and in the Brpnsted-Lowry theory, because in the Br0nsted-Lowry theory a base must donate a pair of electrons in order to accept a proton. 

Any electron-deficient atom can act as a Lewis acid. Many compounds containing group IIIA elements such as boron and aluminum are Lewis acids because group IIIA atoms have only a sextet of electrons in their outer shell. Many other compounds that have atoms with vacant orbitals also act as Lewis acids. Zinc and iron(III) halides (ferric halides) are frequently used as Lewis acids in organic reactions. 

Verify for yourself that you can calculate the formal charges in these structures. 

A zinc ion acts as a Lewis acid in the mechanism of the enzyme carbonic anhydrase (Chapter 24). 

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

Lewis bases are structurally the same as Br0nsted-Lowry bases. Both have an avaQable electron pair—a lone pair or an electron pair in a ji bond. A Br0nsted-Lowry base always donates this electron pair to a proton, but a Lewis base donates this electron pair to anything that is electron deficient. 

A Lewis acid must be able to accept an electron pair, but there are many ways for this to occur. All Brpnsted-Lowry acids are also Lewis acids, but the reverse is not necessarily true. Any 

Common examples of Lewis acids (which are not Br0nsted-Lowry acids) include BFg and AICI3. These compounds contain elements in group 3A of the periodic table that can accept an electron pair because they do not have fiUed valence shells of electrons. 

These compounds are both Bronsted-Lowry acids and Lewis acids. 

The Lewis definition of acids and bases is broader than the Bronsted-Lowry definition. According to the Lewis definition, acidity and basicity are described in terms of electrons, rather than protons. A Lewis acid is defined as an electron acceptor, while a Lewis base is defined as an electron donor. As an illustration, consider the following Bronsted-Lowry acid-base reaction  

HCl is an acid according to either definition. It is a Lewis acid because it serves as an electron acceptor, and it is a Bronsted-Lowry acid because it serves as a proton donor. But the Lewis definition is an expanded definition of acids and bases, because it includes reagents that would otherwise not be classified as acids or bases. For example, consider the following reaction  

According to the Bronsted-Lowty definition, BF3 is not considered an acid because it is has no protons and cannot serve as a proton donor. However, according to the Lewis definition, BF3 can serve as an electron acceptor, and it is therefore a Lewis acid. In the reaetion above, H2O is a Lewis base because it serves as an electron donor. 

Take special notice of the curved-arrow notation. There is only one curved arrow in the reaction above, not two. 

Chapter 6 will introduce the skills necessary to analyze reactions, and in Section 6.7 we will revisit the topic of Lewis acids and bases. In fact, we will see that most of the reactions in this textbook occur as the result of the reaction between a Lewis acid and a Lewis base. For now, lets get some practice identifying Lewis acids and Lewis bases. 

A Lewis acid is any substance that accepts an electron pair in forming a coordinate bond (Section 1.1). Examples include H , BF3, AICI3, TiCLt, ZnCl2 and SnCLt. They have unfilled valence shells and so can accept electron pairs. A Lewis base is any substance that donates an electron pair in forming a coordinate bond. Examples include H2O, ROH, RCHO, R2C=0, R3N and R2S. They all have a lone pair(s) of electrons on the heteroatom (O, N or S). 

Another acid-base theory is the Lewis acid-base theory. According to this theory, a Lewis acid will accept a pair of electrons and a Lewis base will donate a pair of electrons. In order to make it easier to see which species is donating electrons, it is helpful to use Lewis structures for the reactants and if possible for the products. 

The following is an example of a Lewis acid-base reaction. 

The hydrogen ion accepts the lone pair of electrons from the ammonia to form the ammonium ion. The hydrogen ion, because it accepts a pair of electrons, is the Lewis acid. The ammonia, because it donates a pair of electrons, is the Lewis base. This reaction is also a Brpnsted-Lowry acid-base reaction. This illustrates that a substance may be an acid or a base by more than one definition. All Brpnsted-Lowry acids are Lewis acids, and all Brpnsted-Lowry bases are Lewis bases. However, the reverse is not necessarily true. 

Write equations for the following add-base reactions. Label the conjugate acids and bases, and show any resonance stabilization. Predict whether the equilibrium favors the reactants or products. 

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. 

Lewis bases are species with nonbonding electrons that can be donated to form new bonds. Lewis acids are species that can accept these electron pairs to form new bonds. Since a Lewis acid accepts a pair of electrons, it is called an 

A nucleophile donates electrons An electrophile accepts electrons. Acidic protons may serve as electron acceptors. 

The Lewis acid-base definitions include reactions having nothing to do with protons. Following are some examples of Lewis acid-base reactions. Notice that the common Br0nsted-Lowry acids and bases also fall under the Lewis definition, with a proton serving as the electrophile. Curved arrows (red) are used to show the movement of electrons, generally from the nucleophile to the electrophile. 

Lewis bases are species with available electrons that can be donated to form new bonds. Lewis acids are species that can accept these electron pairs to form new bonds. Since a Lewis acid accepts a pair of electrons, it is called an electrophile, from the Greek words meaning lover of electrons. A Lewis base is called a nucleophile, or lover of nuclei, because it donates electrons to a nucleus with an empty (or easily vacated) orbital. In this book, we sometimes use colored type for emphasis blue for nucleophiles, green for electrophiles, and occasionally red for acidic protons. 

Some of the terms associated with acids and bases have evolved specific meanings in organic chemistry. When organic chemists use the term base, they usually mean a proton acceptor (a Brpnsted-Lowry base). Similarly, the term acid usually means a proton donor (a Brpnsted-Lowry acid). When the acid-base reaction involves formation of a bond to some other element (especially carbon), organic chemists refer to the electron donor as a nucleophile (Lewis base) and the electron acceptor as an electrophile (Lewis acid). 

The following illustration shows electrostatic potential maps for the reaction of NH3 (the nucleophile/electron donor) with BF3 (the electrophile/electron acceptor). The 

In the previous section, we presented ideas about the molecular structures of acids and bases. In 1923, G. N. Lewis proposed an acid-base theory closely related to bonding and structure. The Lewis acid-base theory is not limited to reactions involving H and OH It extends acid-base concepts to reactions in gases and in solids. It is especially important in describing certain reactions between organic molecules. 

A Lewis acid is a species (an atom, ion, or molecule) that is an electron-pair acceptor, and a Lewis base is a species that is an electron-pair donor. A reaction between a Lewis acid (A) and a Lewis base (B=) results in the formation of a covalent bond between them. The product of a Lewis acid-base reaction is called an adduct (or addition compound). The reaction can be represented as 

By these definitions, OH , a Bronsted-Lowry base, is also a Lewis base because lone-pair electrons are present on the O atom. So too is NH3 a Lewis base. HCl, conversely, is not a Lewis acid It is not an electron-pair acceptor. We can think of HCl as producing H, however, and H is a Lewis acid. H forms a coordinate covalent bond with an available electron pair. 

Species with an incomplete valence shell are Lewis acids. When the Lewis acid forms a coordinate covalent bond with a Lewis base, the octet is completed. A good example of octet completion is the reaction of BF3 and NH3. 

The reaction of lime (CaO) with sulfur dioxide is an important reaction for reducing SO2 emissions from coal-fired power plants. This reaction between a solid and a gas underscores that Lewis acid-base reactions can occur in all states of matter. The smaller curved red arrow in reaction (16.24) suggests that an electron pair in the Lewis structure is rearranged. 

For a substance to be a proton acceptor (a Br0nsted-Lowry btise), it must have an unshared pair of electrons for binding the proton, as, for example, in NH3. Using Lewis structures, we can write the reaction between and NH3 as 

Every base that we have discussed thus far—whether OH , H2O, an amine, or an anion—is an electron-pair donor. Everything that is a base in the Bronsted-Lowry sense (a proton acceptor) is also a base in the Lewis sense (an electron-pair donor). In the Lewis theory, however, a base can donate its electron pair to something other than H. The Lewis definition therefore greatly increases the number of species that can be considered acids in other words, is a Lewis acid but not the only one. For example, the reaction between NH3 and BF3 occurs because BF3 has a vacant orbital in its valence shell, ooo (Section 8.7) It therefore acts as an electron-pair acceptor (a Lewis acid) toward NH3, which donates the electron pair  

What feature must any molecule or ion have in order to act as a Lewis acid  

Our emphasis throughout this chapter has been on water as the solvent and on the proton as the source of acidic properties. In such cases we find the Bronsted-Lowry definition of acids and bases to be the most useful. In fact, when we speak of a substance as being acidic or basic, we are usually thinking of aqueous solutions and using these terms in the Arrhenius or Bronsted-Lowry sense. The advantage of the Lewis definitions of acid and base is that they allow us to treat a wider variety of reactions, including 

So far we have discussed acid-base properties in terms of the Bronsted theory. For example, a Bronsted base is a substance that must be able to accept protons. By this definition, both the hydroxide ion and ammonia are bases  

In each case, the atom to which the proton becomes attached possesses at least one unshared pair of electrons. This characteristic property of OH , NH3, and other Bronsted bases suggests a more general definition of acids and bases. 

In 1932 G. N. Lewis defined what we now call a Lems base as a substance that can donate a pair of electrons. A Lewis acid is a substarKe that can accept a pair of electrons. In the protonation of ammonia, for example, NH3 acts as a Lewis base because it donates a pair of electrons to the proton H, which acts as a Lewis acid by accepting the pair of electrons. A Lewis acid-base reaction, therefore, is one that involves the donation of a pair of electrons from one species to another. 

The significance of the Lewis concept is that it is more general than other definitions. Lewis acid-base reactions include many reactions that do not involve Bronsted acids. Consider, for example, the reaction between boron trifluoride (BF3) and ammonia to form an adduct compound  

The B atom inBFs is 5p -hybridized [M Section 9.4]. The vacant, unhybridized Ip orbital accepts the pair of electrons from NH3. Thus, BF3 functions as an acid according to the Lewis definition, even though it does not contain an ionizable proton. A coordinate covalent bond [M Section 8.8] is formed between the B and N atoms. In fact, every Lewis acid-base reaction results in the formation of a coordinate covalent bond. 

Another Lewis acid containing boron is boric acid (H3BO3). Boric acid (a weak acid used in eyewash) is an oxoacid with the following stractnre  

Note that boric acid does not ionize in water to produce an tion with water is 

The Arrhenius definition of an acid and a base attributed acidity to the presence of H (aq), and alkalinity to OH (aq). Br0nsted-Lowry theory generalizes the acid-base concept by focusing on proton transfer, rather than on particular aqueous ions. Here, we discuss an attempt to generalize it further by focusing on the changes in electronic structure that occur when acid-base reactions take place, ideas introduced by G. N. Lewis. 

When ammonia, a base, accepts a proton from an acid, the process can be written  

The boron atom acquires these two electrons if BF3 and NH3 are mixed  

A solid compound, H3NBF3, is formed as a white smoke (H). In this compound the coordination around both the boron and the nitrogen atoms is tetrahedral. 

The nitrogen-boron bond in the solid compound is a dative bond because both shared electrons come from the nitrogen atom of the ammonia. The bond can therefore be written H3N Bp3. An alternative, but precisely equivalent way of writing it is H3 N— BF3. The latter provides a more accurate description. Consider Equation 4.1 in which NH3 forms a dative bond with H+. The first representation gives Structure 4.3, which suggests that one of the H atoms is different from the other three. However, in the new representation (Structure 4.4), the four N—H bonds are equivalent, consistent with the fact that they are of equal length. 

We began our definitions of acids and bases with the Arrhenius model. We then saw how the Br0nsted-Lowry model, by introducing the concept of a proton donor and proton acceptor, expanded the range of substances that we consider acids and bases. We now introduce a third model, which further broadens the range of substances that we can consider acids. This third model is the Lewis model, named after G. N. Lewis, the American chemist who devised the electron-dot representation of chemical bonding (Section 9.1). While the Br0nsted-Lowry model focuses on the transfer of a proton, the Lewis model focuses on the transfer of an electron pair. Consider the simple acid-base reaction between the ion and NH3, shown here with Lewis structures  

Br0nsted-Lowry model Lewis model focuses 

According to the Br0nsted-Lowry model, the ammonia accepts a proton, thus acting as a base. According to the Lewis model, the ammonia acts as a base by donating an electron pair. The general definitions of acids and bases according to the Lewis model focus on the electron pair. 

Lewis acid electron pair acceptor Lewis base electron pair donor 

According to the Lewis definition, H+ in the reaction just shown is acting as an acid because it is accepting an electron pair from NH3. NH3 is acting as a Lewis base because it is donating an electron pair to 

The carbonyl bases constitute another important class of weak bases that present interesting possibilities for investigation of structural effects. In solution, experiments with these compounds are subject to severe difficulties. The result is a serious lack of agreement among different investigators about pAa. Arnett and co-workers point out that pKa values reported for acetophenone cover a range of over four units ( — 3.65 to —7.99), while those for acetone span seven units (-0.2 to -7.2).102 In view of these uncertainties, it is impossible to say whether aldehydes, ketones, or carboxylic acids are the most basic in solution. Gas-phase data are available for some of these substances. 

In 1923 G. N. Lewis proposed a definition of acids and bases somewhat different from that of Bronsted 106 

Lewis acids are thus electron-deficient molecules or ions such as BF-, or carbo-cations, whereasTewis bases are molecules or ions containing available electrons. such as amines, ethers, alkoxide ions, and so forth.,A Lewis acid-base reaction is the combination of an acid and a base to form a complex, or adduct. The stabilities of these adducts depend on the structures of the constituent acid and base and vary over a wide range. Some examples of Lewis acid—base reactions are given in Table 3.19. Lewis acid-base reactions abound in organic chemistry  

105 A general discussion of Lewis acids and bases is given by R. J. Gillespie in Friedel-Crqfts and Related Reactions, Vol. 1, G. A. Olah, Ed., Wiley-Interscience, New York, 1963, p. 169. 

Valence and the Structure of Atoms and Molecules, American Chemical Society Monograph, The Chemical Catalog Co., New York, 1923. Lewis also gave a definition equivalent to that of Bronsted at this time, but he considered the electron-pair definition to be more general. 

It is interesting that beryllium hydroxide, like aluminum hydroxide, exhibits amphoterism  

This is another example of the diagonal relationship between berylhum and aluminum (see p. 344). 

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We have seen that a base can be defined as combining with a proton and, therefore, requires at least one lone pair of electrons. A more general definition of acids and bases, due to G. N. Lewis, describes a base as any species (atom, ion or molecule) which can donate an electron pair, and an acid as any species which can accept an electron pair— more simply, a base is an electron-pair donor, an acid an electron-pair acceptor. Some examples of Lewis acids and bases are ... 

Analogous to the classification of Lewis acids and bases in hard and soft species, Ahrland et al. have su ested a division of donors and acceptors into classes a and 6. See Ahrland, S. Chatt, J. Davies, N.R. Quart. Rev. 1958, 77, 265... 

In Chapter 6 we survey what has been accomplished and indicate directions for future research. Furthermore, we critically review the influence of water on Lewis acid - Lewis base interactions. This influence has severe implications for catalysis, in particular when hard Lewis acids and bases are involved. We conclude that claims of Lewis-acid catalysis should be accompanied by evidence for a direct interaction between catalyst and substrate. 

The strength of the complexation is a function of both the donor atom and the metal ion. The solvent medium is also an important factor because solvent molecules that are potential electron donors can compete for the Lewis acid. Qualitative predictions about the strength of donor-acceptor complexation can be made on the basis of the hard-soft-acid-base concept (see Section 1.2.3). The better matched the donor and acceptor, the stronger is the complexation. Scheme 4.3 gives an ordering of hardness and softness for some neutral and ionic Lewis acids and bases. 

Continuum effects indicated by hard and soft acid-base (Lewis acids) and bases. C. K. Jorgensen, Top. Curr. Chem., 1975,56,1-66 (210). 

Identify Bronsted and Lewis acids and bases in a chemical reaction (Self-Test 10.2). 

What Do We Need to Know Already This chapter develops the ideas in Chapters 9 and 10 and applies them to equilibria involving ions in aqueous solution. To prepare for the sections on titrations, review Section L. For the discussion of solubility equilibria, review Section I. The discussion of Lewis acids and bases in Section 11.13 is based on Section 10.2. 

What Do We Need to Know Already This chapter draws on many of the principles introduced in the preceding chapters. In particular, it makes use of the electron configurations of atoms and ions (Sections 1.13 and 2.1) and the classification of species as Lewis acids and bases (Section 10.2). Molecular orbital theory (Sections 3.8 through 3.12) plays an important role in Section 16.12. 

Lewis Acids and Bases Hard and Soft Acids and Bases... 

For a monograph on Lewis acid-base theory, see Jensen, W.B. The Lewis Acid-Base Concept Wiley NY, 1980. For a discussion of the definitions of Lewis acid and base, see Jensen, W.B. Chem. Rev, 1978, 78, 1. 

The scope of this reaction is similar to that of 10-21. Though anhydrides are somewhat less reactive than acyl halides, they are often used to prepare carboxylic esters. Acids, Lewis acids, and bases are often used as catalysts—most often, pyridine. Catalysis by pyridine is of the nucleophilic type (see 10-9). 4-(A,A-Dimethylamino)pyridine is a better catalyst than pyridine and can be used in cases where pyridine fails. " Nonbasic catalysts are cobalt(II) chloride " and TaCls—Si02. " Formic anhydride is not a stable compound but esters of formic acid can be prepared by treating alcohols " or phenols " with acetic-formic anhydride. Cyclic anhydrides give monoesterified dicarboxylic acids, for example,... 

Example provides practice in recognizing Lewis acids and bases. 

Identify the Lewis acids and bases in each of the following reactions and draw structures of the resulting adducts ... 

Lewis acids and bases can be organized according to their polarizability. If polarizability is low, the species is categorized as hard. If polarizability is high, the species is soft. ... 

The terms hard and soft are relative, so there is no sharp dividing line between the two, and many Lewis acids and bases are intermediate between hard and soft. Example shows how to categorize Lewis acids and bases according to their hard-soft properties. 

The classification of Lewis acids and bases relevant to AB cements is shown below. 

Yatsimirskii considered that the hard and soft classification was too general and proposed instead a more detailed approach. He classified Lewis acids and bases into six groups, based on the nature of the adduct bonding. 

III. Comparison of Behavior of Lewis Acids and Bases in Aqueous Solution and in... 

The concept of hard and soft acids and bases ( HSAB ) should also be mentioned here. This is not a new theory of acids and bases but represents a useful classification of Lewis acids and bases from the point of view of their reactivity, as introduced by R. G. Pearson. 

Jolly, W. L. Inorganic Applications of X-Ray Photoelectron Spectroscopy. 71, 149-182 (1977). Jorgensen, C. K. Continuum Effects Indicated by Hard and Soft Antibases (Lewis Acids) and Bases. 56, 1-66 (1975). 

Since the discovery of the first noble gas compound, Xe PtF (Bartlett, 1962), a number of compounds of krypton, xenon, and radon have been prepared. Xenon has been shown to have a very rich chemistry, encompassing simple fluorides, XeF2> XeF, and XeF oxides, XeO and XeO oxyf luorides, XeOF2> XeOF, and Xe02 2 perxenates perchlorates fluorosulfates and many adducts with Lewis acids and bases (Bartlett and Sladky, 1973). Krypton compounds are less stable than xenon compounds, hence only about a dozen have been prepared KrF and derivatives of KrF2> such as KrF+SbF, KrF+VF, and KrF+Ta2F11. The chemistry of radon has been studied by radioactive tracer methods, since there are no stable isotopes of this element, and it has been deduced that radon also forms a difluoride and several complex salts. In this paper, some of the methods of preparation and properties of radon compounds are described. For further information concerning the chemistry, the reader is referred to a recent review (Stein, 1983). 

Early attempts to fathom organic reactions were based on their classification into ionic (heterolytic) or free-radical (homolytic) types.1 These were later subclassified in terms of either electrophilic or nucleophilic reactivity of both ionic and paramagnetic intermediates - but none of these classifications carries with it any quantitative mechanistic information. Alternatively, organic reactions have been described in terms of acids and bases in the restricted Bronsted sense, or more generally in terms of Lewis acids and bases to generate cations and anions. However, organic cations are subject to one-electron reduction (and anions to oxidation) to produce radicals, i.e.,... 

One of the problems encountered when dealing with the interaction of Lewis acids and bases in a quantitative way is in evaluating the role of the solvent. Bond energies in molecules are values based on the molecule in the gas phase. However, it is not possible to study the interaction of many Lewis acids and bases in the gas phase because the adducts formed are not sufficiently stable to exist at the temperature necessary to convert the reactants to gases. For example, the reaction between pyridine and phenol takes place readily in solution as a result of hydrogen bonding ... 

It is the interaction of the Lewis acid and base in the gas phase that gives the enthalpy of the A B bond, AHAB, but what is more often (and conveniently) measured is the enthalpy change when the reaction is carried out in solution, AH AB. 

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