Perspective Chapter: The Acidity Concept According to Lewis

Abstract

There are certain essential concepts that need to be considered before moving on to the reaction mechanisms in organic chemistry. Having learned basic concepts appropriately, one can understand reaction mechanisms better. For this reason, we will primarily focus on basic concepts such as electrophile and lewis acids in this part. This chapter presents electrophilic compounds, lewis acids, and bases, which are important in organic chemistry, with understandable examples. The reactivity of polarized bonds to be formed by groups attached to a compound, Lewis acid/base reactions and the reactions of metals with related concepts are clearly explained with examples. In addition, the reaction mechanisms mentioned in this chapter are presented according to the molecular orbital approach and acidity and basicity properties.

1. Introduction: Electrophiles

1.1 Electrophilic compounds

Chemical reactions are reactions based on the exchange of electrons. One compound gives out its electrons, while the other receives electrons. Meanwhile, atoms try to complete the number of electrons in their outer orbits in accordance with the octet rule. For a compound to accept electrons, it must be poor in electrons. Compounds that have fewer electrons and can receive electrons are called electrophilic compounds or electrophiles. Electrophile literally means electron loving. Electrophiles react with electron-rich compounds and try to fill or minimize the electron vacancy in their bodies [1].

It is possible to classify electrophilic compounds under certain groups.

• Cations (compounds with positive (+) charge)

• Lewis acids

• Carbonyl compounds

1.1.1 Compounds containing polarized bonds

Positive (+) charged carbon, nitrogen, halogen compounds, etc. are classified in the cation group. Carbocations are sp² hybridized, positively (+) charged carbon compounds that bind three groups. They have electron vacancies. By reacting with nucleophiles, carbocations try to neutralize the positive (+) charge and to make the number of electrons in their bodies in line with the octet rule (1).

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Lewis acids do not have protogenic properties like classical acids. Compounds that cannot donate protons but have an electron vacancy in their structure are included in this group. Compounds such as AlCl₃, BF₃, and FeCl₃ are neutral and the number of electrons in the outer orbit of the central atom does not comply with the octet rules. The number of electrons in the outer orbit is generally six. These compounds, easily reacting with these electron-rich ones, fill the electron vacancy in their outer orbits (2).

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Carbonyl compounds are also potential electrophilic compounds. The double bonding electrons between carbon and oxygen become polarized because the oxygen attracts the electrons more. Therefore, the carbonyl carbon acquires a positive (+) charge, exhibits electrophilic properties, and reacts easily with nucleophiles. The double bond between the sulfur and oxygen also becomes polarized and the sulfur shows electrophilic properties (3).

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Other than these compounds, all kinds of compounds with bond polarization show electrophilic properties. For example, since halogens are more electronegative than the carbon atom in alkyl halide, they attract the carbon halogen bond electrons towards them. As a result, the carbon atom becomes partially positively charged, while the halogen atom is negatively charged. Therefore, carbons attached to the halogen atom show electrophilic properties and react easily with nucleophiles (4).

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1.1.2 Lewis acid

While the Arrhenius acid-base theory does not take into account the acid-base phenomenon in anhydrous environments, the Bronsted-Lowry acid-base theory completely excludes compounds that do not contain protons, since it is only about proton exchange. According to a theory that the American physical chemist Gilbert Newton Lewis (1875–1946) first brought up in 1923 and developed in 1938, compounds that can form a new bond by gaining a pair of electrons from a molecule or anion are called Lewis acids, and compounds that form a new bond by donating a pair of electrons to a molecule or cation are called Lewis bases. According to this definition, the proton is a Lewis acid. This definition covers many compounds that do not contain protons.

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Let us examine the reaction between the proton and the H₂O molecule within the framework of Lewis acid-base theory. Since the proton has an electron vacancy in its outer orbit, it can easily accept a pair of electrons and is a Lewis acid. The oxygen of water, on the other hand, is a Lewis base because it can form a new bond by giving its free electrons to another atom. The proton forms a new sigma bond binding to the oxygen with a pair of free electrons on the oxygen (5).

Lewis acid-base theory explains why BF₃ reacts with NH₃ to form a complex. Since the hybridization of the boron atom in the BF₃ molecule is sp², the structure of the molecule is plane and there is also an empty p orbital on the boron. Boron trifluoride is a Lewis acid because the number of electrons in the outer orbital of the boron atom is six. This orbital takes a pair of electrons and turns them into a structure complying with the octet rule. In the ammonia molecule, there is a free electron pair that does not bind to the nitrogen atom. Thus, ammonia is a Lewis base. In the complex formation between NH₃ and BF₃, the nitrogen atom crosschecks the octet rule transferring its free electrons to the vacant p orbital of the boron atom (6).

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Likewise, AlCl₃ is also a Lewis acid because it has only six electrons in its aluminum outer shell (7). Compounds such as ZnCl₂, FeCl₂, SnCl₄, and TiCl₄, which have an electron vacancy in their outer orbits, also fall into the Lewis acid group.

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In the Friedel-Crafts alkylation of aromatic compounds, compounds such as AlCl₃ are used as catalysts to increase the reactivity of alkyl halides. Aluminum trichloride forms a complex with a pair of nonbonding electrons of the alkyl halide. Aluminum trichloride is a Lewis acid because it has an electron vacancy in its outer shell. The halogen attached to the alkyl group is the Lewis base. The free electrons of the halogen attached to the alkyl group attack this vacancy and form a new covalent bond (8).

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Carbocations are also classified in Lewis acids. A carbocation has, as shown earlier, three sp² hybrid orbitals and one vacant p orbital. Carbocations are compounds that are extremely prone to bond with Lewis bases (9).

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Double bonds with low electron density, for example, enon-type compounds (alpha, beta-unsaturated systems) react with Lewis bases and form covalent bonds due to the electron vacancy formed in the beta-position (10).

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The carbon atoms of double bonds of tetracyanoethylene also have Lewis acid properties. Cyanide groups reduce the electron density of the double bond due to their strong electron-attracting properties. Therefore, nucleophiles (Lewis bases) attack carbon atoms of the double bond (11).

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The hydride ion is the typical Lewis base. It reacts instantly with Lewis acids and forms covalent bonds. Unpolarized double-bond electrons are also Lewis bases. Since Lewis bases have high electron density, they instantly react with Lewis acids.

Polarized bonds form a dipole. One end of these bonds is a Lewis acid and the other end is a Lewis base. For example; carbon-halogen bonds are polarized bonds. Since the halogen atom attracts the sigma electrons between carbon and halogen more strongly and has non-free bonding electrons, it is a Lewis base. The carbon atom with reduced electron density acts as a Lewis acid. There is also a polarization in the carbonyl group. Therefore, the carbonyl carbon is the Lewis acid and the carbonyl oxygen is the Lewis base (12).

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1.1.3 Lewis acid: Base definition

GN Lewis defined acids as substances that accept electron pairs and bases as substances that donate electron pairs. This definition is broader in scope than the Bronsted-Lowry definition. Reactions between compounds that do not contain protons are also acid-base reactions according to this definition. However, there is consistency between the two definitions. According to Bronsted-Lowry, acids are substances that donate positively charged particles (protons), whereas according to Lewis, acids are defined as substances that take negatively charged particles (electron pairs). According to the Lewis system, an example of a characteristic acid-base reaction is the reaction between trialkyl amine and boron trifluoride (13) [2].

$${\mathbf{R}}_{\mathbf{3}}\mathbf{N}\mathbf{:}\mathbf{+}{\mathbf{BF}}_{\mathbf{3}}\rightleftharpoons {\mathbf{R}}_{\mathbf{3}}\stackrel{\mathbf{+}}{\mathbf{N}}—\stackrel{-}{\mathbf{B}}{\mathbf{F}}_{\mathbf{3}}$$E13

Trialkyl amine has an unshared electron pair on the nitrogen atom. In boron trifluoride, the boron atom has not completed its octet and has an electron pair deficiency. As the molecules combine, the unshared electron pair on the nitrogen atom forms the NB covalent bond. Since the nitrogen atom donates electrons during the reaction, the R₃N molecule containing this atom is Lewis base, and the BF₃ molecule containing the B atom that gains electrons is Lewis acid. The nitrogen atom is called a donor atom, meaning that it donates electrons, and the boron atom is called an acceptor atom, which means that it takes electrons.

Since the central atom or ion receives electrons in the formation of coordination compounds, they are Lewis acids. Ligands attached to the central atom are Lewis bases because they donate electrons.

$$\mathbf{4}\mathbf{.}\mathbf{CO}\mathbf{+}\mathbf{Ni}\to \mathbf{Ni}{\left(\mathbf{CO}\right)}_{\mathbf{4}}$$E14

$$\mathbf{2}\mathbf{.}{\mathbf{NH}}_{\mathbf{3}}\mathbf{+}{\mathbf{Ag}}^{\mathbf{+}}\to {\left[\mathbf{Ag}{\left({\mathbf{NH}}_{\mathbf{3}}\right)}_{\mathbf{2}}\right]}^{\mathbf{+}}$$E15

In their reaction, CO and NH₃ are Lewis bases since they donate electron pairs, and Lewis acid, Ni, and Ag⁺ accept electron pairs. In the formation of coordination compounds, since the central atom or ion gains electrons, Lewis acid. Ligands attached to the central atom are Lewis bases because they donate electrons (14 and 15).

Lewis acids:

1. Since they can accept electron pairs, all cations are Lewis acid (16 and 17).

   $${\mathbf{Fe}}^{\mathbf{3}\mathbf{+}}\mathbf{+}\mathbf{6}{\mathbf{CN}}^{\mathbf{-}}\to {\left[\mathbf{Fe}{\left(\mathbf{CN}\right)}_{\mathbf{6}}\right]}^{\mathbf{3}\mathbf{-}}$$E16
   $${\mathbf{Cr}}^{\mathbf{3}\mathbf{+}}\mathbf{+}\mathbf{6}{\mathbf{NH}}_{\mathbf{3}}\to {\left[\mathbf{Cr}{\left({\mathbf{NH}}_{\mathbf{3}}\right)}_{\mathbf{6}}\right]}^{\mathbf{3}\mathbf{+}}$$E17

2. Compounds containing a central atom that gains electrons in the valence shell and can increase the coordination number act as a Lewis acid (18–20).

   $${\mathbf{AIF}}_{\mathbf{3}}\mathbf{+}\mathbf{3}{\mathbf{F}}_{\mathbf{6}}\to {\left[{\mathbf{AIF}}_{\mathbf{6}}\right]}^{\mathbf{3}\mathbf{-}}$$E18
   $${\mathbf{\text{SnCl}}}_{\mathbf{4}}\mathbf{+}\mathbf{2}{\mathbf{Cl}}^{\mathbf{-}}\to {\left[{\mathbf{\text{SnCl}}}_{\mathbf{6}}\right]}^{\mathbf{6}\mathbf{-}}$$E19
   $${\mathbf{SbF}}_{\mathbf{5}}\mathbf{+}{\mathbf{F}}^{\mathbf{-}}\to {\left[{\mathbf{SbF}}_{\mathbf{6}}\right]}^{\mathbf{-}}$$E20

3. Molecules such as CO₂ and SO₃ that have one or more multiple bonds in their central atoms act as Lewis acids (21).

$${\mathbf{CO}}_{\mathbf{2}}\mathbf{+}{\mathbf{OH}}^{\mathbf{-}}\to {\mathbf{HCO}}_{\mathbf{3}}^{\mathbf{-}}$$E21

General grouping can be made for Lewis bases.

1. All anions are Lewis bases. Increasing the charge density increases the base force (22).

   $${\mathbf{H}}^{+}\mathbf{+}{\mathbf{OH}}^{\mathbf{-}}\to {\mathbf{H}}_{\mathbf{2}}\mathbf{O}{\mathrm{OH}}^{-}\text{is\hspace{0.17em}\hspace{0.17em}lewis\hspace{0.17em}\hspace{0.17em}base}.$$E22

2. Molecules with unshared electron pairs, such as water, alcohol, and ether, act as Lewis bases (23).

   $${\mathbf{Cu}}^{\mathbf{2}\mathbf{+}}\mathbf{+}\mathbf{4}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\to {\left[\mathbf{Cu}{\left({\mathbf{H}}_{\mathbf{2}}\mathbf{O}\right)}_{\mathbf{4}}\right]}^{\mathbf{2}\mathbf{+}}$$E23

3. Alkynes that can form coordinate covalent bonds with metal ions acts as a Lewis base (24 and 25).

$${\mathbf{Ag}}^{\mathbf{+}}\mathbf{+}{\mathbf{C}}_{\mathbf{2}}{\mathbf{H}}_{\mathbf{4}}\to {\left[\mathbf{Ag}\left({\mathbf{C}}_{\mathbf{2}}{\mathbf{H}}_{\mathbf{4}}\right)\right]}^{\mathbf{+}}$$E24

$$\mathbf{Cr}\mathbf{+}\mathbf{2}{\mathbf{C}}_{\mathbf{6}}{\mathbf{H}}_{\mathbf{6}}\to \mathbf{Cr}{\left({\mathbf{C}}_{\mathbf{6}}{\mathbf{H}}_{\mathbf{6}}\right)}_{\mathbf{2}}$$E25

1.1.4 Lewis definition

Lewis acid; electron pair acceptor; some defined it as a chemical species with an electron pair acceptor. A coordinate-covalent bond is formed as a result of electron pair transfer from B to A, with A Lewis acid and B Lewis base (26):

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Lewis acid-base reaction from a two-electron reduction-oxidation (redox) reaction is the formation of a coordinate-covalent bond in the first. The oxidation numbers of atoms do not change in the formation of coordinate-covalent bonds. However, the characteristic feature of redox reactions is that the oxidation numbers change. Lewis acid-base reactions can be divided into three main groups:

1. Complex formation reactions: These are reactions in which Lewis acids and bases interact to form an addition product, as in the reaction. In the following reactions, the terms complex compound and complex ion are used for some of the products.

2. Pay attention (27–29):

   

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3. Substitution reactions: In these reactions, a Lewis base is replaced by another Lewis base, or a Lewis acid is replaced by another Lewis acid. The following reactions can be given as examples for the first group (30 and 31);

   

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   An acid can be replaced by another acid, as seen in the following reaction (32).

   

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4. Metathesis (double displacement) reactions: Acids and bases are replaced in these reactions (33).

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The Lewis acid-base concept is a very comprehensive one. It includes many chemical species defined as acids or bases according to Arrhenius, Brönsted-Lowry, and solvent system concepts. This definition is quite helpful in predicting whether two chemical species interact and, if so, the products formed. Therefore, it is widely used in inorganic chemistry and organic chemistry. Another benefit is that it allows us to explain acidity and basicity more easily with the molecular orbital approach. Some Lewis acids and bases listed below will provide more detailed information about Lewis acid-base reactions.

1.1.5 Lewis acids

For a chemical species to act as a Lewis acid, it must have either a low-energy vacant orbital to accept the electron pair or a positive center to attract the electron pair. According to the valence bond theory, the low energy orbital is the valence orbital. In the molecular orbital approach, it is the relatively low energy LUMO (lowest unoccupied molecular orbital). Some Lewis acids with these properties are given below [3].

1. H⁺ Ion: It is the simplest Lewis acid, and its 1 s orbital is empty. It receives an electron pair from a Lewis base into its orbital. For example, its reaction with a Lewis base, HS⁻, is (34):

   

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2. Major cations with high charge/radius ratio: Examples of these are Be²⁺, Mg²⁺, and A1³⁺. The valence orbitals of these ions are empty. Be²⁺ has four vacant orbitals, and the other two have six (two of the 3d orbitals are considered low-energy. Look Coordination Chemistry section). Be²⁺ binds with four molecules of water, a Lewis base, and A1³⁺ with six molecules (35 and 36).

   $${\mathbf{Be}}^{\mathbf{2}\mathbf{+}}\mathbf{+}\mathbf{4}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\to {\left[\mathbf{Be}{\left({\mathbf{H}}_{\mathbf{2}}\mathbf{O}\right)}_{\mathbf{4}}\right]}^{\mathbf{2}\mathbf{+}}$$E35
   $${\mathbf{Al}}^{\mathbf{3}\mathbf{+}}\mathbf{+}\mathbf{6}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\to {\left[\mathbf{Al}{\left({\mathbf{H}}_{\mathbf{2}}\mathbf{O}\right)}_{\mathbf{6}}\right]}^{\mathbf{3}\mathbf{+}}$$E36

   Cations take an electron pair from each water molecule. As shown below, the formal charge of oxygen in the coordinating water molecule is (+1). If it donates the second pair of electrons, the formal charge becomes (+2). Ice is energetically unsuitable for oxygen, which is highly electronegative (37).

   

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   Instead, the following decomposition takes place (38):

   

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   As can be seen, the acidity of aqueous solutions of such cations is a consequence of this dissociation. It is assumed that cations with low charge/radius ratios, such as Na⁺, Rb^+, and Ba²⁺, engage in ion-dipole interactions with Lewis bases, such as water, rather than coordinate-covalent bonding.

   Transition metal cations: These cations contain empty valence orbitals. They receive electron pairs from Lewis bases such as NH₃, H₂O, F⁻, and Cl⁻ (39 and 40).

   $${\mathbf{Co}}^{\mathbf{2}\mathbf{+}}\mathbf{+}\mathbf{6}{\mathbf{NH}}_{\mathbf{3}}\to {\left[\mathbf{Co}{\left({\mathbf{NH}}_{\mathbf{3}}\right)}_{\mathbf{6}}\right]}^{\mathbf{2}\mathbf{+}}$$E39
   $${\mathbf{Fe}}^{\mathbf{3}\mathbf{+}}\mathbf{+}\mathbf{4}{\mathbf{Cl}}^{-}\to {\mathbf{\text{FeCl}}}_{\mathbf{4}}$$E40

3. Molecules with incomplete octets: These are molecules such as BeCI₂, BF₃, and AlCl₃. In the primary resonance structures of these molecules, the central atoms contain vacant valence orbitals. For example, BeCl₂ has two and one empty orbitals. For this reason, BeCl₂ reacts with two ether molecules and B(CH₃)₃ with one pyridine (C₅H₅N)) molecule (41 and 42).

   

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4. Molecules with empty d orbitals: As mentioned earlier, two of the valence d orbitals of the main elements can enter into a bonding. Molecules with at least one empty orbital behave as Lewis acids. Examples of these are SnCl₂, PCl₃, SiCl_4, and SbF₅. Since SbF₅ has a low energy d orbital, it receives a lone pair from the Lewis base (43).

   

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   SiCl₄ has two low-energy d orbitals. To highlight the importance of the Lewis definition, let us consider the mechanism of the hydrolysis reaction of SiCl₄. Since SiCl₄ has two vacant d orbitals, it first binds to two H₂O molecules (44).

   

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   In the resulting structure, oxygens (+1) and silicon (−2) have formal charges. Therefore, the structure is unstable, and 2H⁺ and 2Cl⁻ leave the structure, as shown below (45).

   

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   Since it has two vacant d orbitals in SiCl₂(OH)₂, it binds two H₂O molecules again. 2H⁺ and 2Cl⁻ are separated from this structure and Si(OH)₄ is obtained (46).

   

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   The hydrolysis reaction of SiC1₄ is (47),

   $${\text{SiCl}}_4+{\mathrm{H}}_2\mathrm{O}\to \mathrm{Si}{\left(\mathrm{OH}\right)}_4+4\mathrm{HCI}$$E47

   when reactions are summed side by side. This mechanism also explains why CCl₄ does not hydrolyze. It is because the valence shell of carbon does not contain the d orbital.

5. Molecules with multiple bonds between atoms with different electronegativity: In these molecules, the electropositive atom is partially positively charged. As a result of the attack of the electron pair of the Lewis base to this positive center, the electronegative atom receives the bond electrons as a lone pair. Thus, the Lewis base binds to the molecule. Examples of molecules in this group are CO₂, SO₂, SO₃, and R₂C=O let us consider the reaction between CO₂ and H₂O as an example. Since carbon is the positive center, H₂O attacks this center, and an oxygen atom in CO₂ takes up the electrons of a π bond as a lone pair (48) [4].

   

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   Since oxygen has a positive formal charge in the resulting structure, it takes the electrons of an O—H bond and H⁺ leaves. This ion binds to (−1) charged oxygen (49):

   

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   The sum of the two reactions gives the (50)

   $${\mathbf{CO}}_{\mathbf{2}}\mathbf{+}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\to {\mathbf{H}}_{\mathbf{2}}{\mathbf{CO}}_{\mathbf{3}}$$E50

   reaction. As seen in the first step, CO₂ is a Lewis acid because it takes a lone pair from H₂O and forms a coordinate covalent bond.

6. Some transition elements: Atoms such as Cr, Fe, and Ni receive lone pairs from Lewis bases such as CO and NO (51) [5]:

   $$\mathbf{Cr}\mathbf{+}\mathbf{4}\mathbf{NO}\to \mathbf{Cr}{\left(\mathbf{NO}\right)}_{\mathbf{4}}$$E51

1.1.6 Acidity and basicity according to the molecular orbital approach

A significant role of the Lewis definition is that it allows us to explain acidity and basicity with molecular orbital theory. The highest occupied molecular orbital of a molecule is denoted by HOMO, and the lowest unoccupied molecular orbital is denoted by LUMO. According to the molecular orbital theory, the acidity of a molecule depends on LUMO, and its basicity depends on HOMO. Those with relatively low energy LUMOs act as acidic, and those with relatively high energy HOMOs and especially non-bonding molecular orbitals behave as basic. Because of non-bonding molecules, the electron pair in the orbital is under the influence of a single nucleus. The molecule has a higher tendency to donate these electrons. Molecules such as BF₃ and SiCl₄ behave as acids because they contain relatively low-energy LUMO. Likewise, the behavior of molecules such as CO and H₂C=CH₂ as π-acceptors is due to the relatively low energy π* molecules orbital (LUMO).

The molecular orbital theory explains the formation of coordinate-covalent bonds by the interaction of the LUMO of the acid with the HOMO of the base. A bonding and an antibonding molecular orbital are formed from the interaction of these two orbitals. The lone pair lower energy bond in the HOMO enters the molecular orbital. Since the counter-bonding molecular orbital is empty, a covalent bond is formed in the interaction of HOMO and LUMO. One point to consider here is the difference between the energy levels of HOMO and LUMO. If this difference is too high complete electron transfer from HOMO to LUMO will occur, and instead of covalent interaction, there will be electrostatic interaction. Since the Lewis definition requires covalent interaction, the chemical species would not be defined as acid and base in such a case [6].