Recap
In the previous part of the article, we discussed three different prevalent theories to explain acids and bases: the Arrhenius theory, the Bronsted-Lowry theory, as well as the Lewis theory. The Arrhenius theory suggests that chemical species which dissociate to form H+ ions in solution are called acids while chemical species which dissociate to form OH- ions in solution are called bases. The scope of this theory is limited, as we noted, because acids and bases can only be defined by the Arrhenius theory if they can dissolve. Many acids and bases which cannot dissolve in the same solvent are not covered by the Arrhenius theory. The second theory, and arguably the most widely used theory in organic chemistry, is the Bronsted-Lowry theory. It defines acids as proton donors and bases as proton acceptors. In a Bronsted-Lowry acid-base reaction, the acid donates a proton, forming a basic species known as a conjugate base. The base accepts the proton, forming an acidic species known as a conjugate acid.
We also mentioned the Lewis theory and its related application, the hard-soft acid-base (HSAB) theory. The Lewis theory tells us that acids are any chemical species with a vacant orbital, while a base is any chemical species with an available pair of electrons for donation. By the Lewis theory, a Lewis acid-base reaction is regarded as an exchange of electrons. Importantly, for Lewis theory, we only regard the proton as the acid and not the entire molecule containing the proton. Because the definition of an acid-base reaction is now about electron exchanges, a nucleophilic substitution reaction (Fig. 1) can also be considered a Lewis acid-base reaction, shown in the SN2 mechanism especially (discussed here). In the first step of the SN2 mechanism, the nucleophile uses its pair of electrons to attack the vacant orbitals of the target atom on the substrate, forming a bond with it. This is considered a Lewis acid-base reaction, because there exists an exchange of electrons.
Fig. 1: Example of the SN2 mechanism.
The Lewis theory has the notable application of hard-soft acid-base theory. This refers to the idea that acids and bases can be classified by the unique quality of hardness or softness, and the rate of a Lewis acid-base reaction will depend on this unique quality. Soft bases refer to basic atoms which have small electronegativities and high polarizabilities, while hard bases refer to basic atoms with the exact opposite, high electronegativities and small polarizabilities. In contrast, soft acids are larger and have higher polarizabilities, while hard acids are smaller and have lower polarizabilities. This classification is useful because hard acids prefer hard bases and soft acids prefer soft bases, forming ionic and covalent bonds respectively. With preferences comes faster reaction rates.
Applications of HSAB Theory
As mentioned in the previous article, the HSAB theory has some utility in the field of coordination chemistry. In a coordination compound or complex (Fig. 2), there is a central metal atom; the substituents of this atom, known as ligands, form dative covalent bonds to the metal atom. Dative covalent bonds refer to covalent bonds where both of the electrons are contributed by one atom; in this case, it is the ligands. What is so significant about this? Remember that the definition of a Lewis acid is an atom with vacant orbitals for bonding. The central metal atom does indeed have a large number of orbitals since it is usually further down the periodic table. More importantly, the ligands are considered Lewis bases because they are ‘electron pair donors’, they donate electron pairs to the central metal atom by forming a bond with it. Because of this, the HSAB theory can be applied to coordination chemistry and have its effects on the rate predicted or analyzed accordingly.
Fig. 2: Structure of a coordination complex.
While we will not delve into the nomenclature of the coordination complexes, we can observe something important about the complex shown in Fig. 2, which is Zeise’s salt. What is interesting is that we can see the alkene acting as a ligand and donating electrons to the central platinum atom. This can work because an alkene contains pi electrons that can be donated. The same applies for aryls, which have a pi system and can coordinate as well. Because of this, alkenes and aryls may be considered under HSAB theory. Both alkenes and aryls display little electronegativity and generally have relatively large polarities, making them soft bases. This is very important as it means that if alkenes and aryls act as ligands, they will always prefer, and have higher reaction rates, when reacting with soft acids (metal atoms). As we have noted before, hard acids are generally smaller (higher polarizabilities); as such, metal atoms with smaller radii such as Na+, K+ and Fe3+ are considered hard acids. In contrast, soft acids are generally larger (lower polarizabilities), thus metal atoms with large radii such as Cu+, Ag+ and Pt2+ are considered soft acids.
Because alkenes and aryls are both soft bases, they prefer to react with soft acids, and this means that they prefer to react with larger metals. As demonstrated by Fig. 2, the alkene forms the dative covalent bond with platinum, which is large and has been listed as a soft acid above. Unusually, both chromium and iron, which are relatively much smaller than platinum, are nevertheless able to form complexes with alkenes and aryls present in them as ligands. Perhaps the most well-known such complex is that of ferrocene, which involves the central iron atom bonding to two aryls, which are soft. The reasoning behind this is that there is more than one factor affecting the hardness and softness of acids and bases; as with acidity and basicity, there is no benchmark for whether an acid or base is soft or hard. Chromium ions with lower oxidation states may form bonds with alkenes or aryls, as can some iron ions.
Fig. 2: Structure of ferrocene.
Factors for Acidity or Basicity
Acidity and basicity are closely interlinked. We note that no proper benchmark exists for what is an acid and for what is a base. The more acidic a compound, the less basic, and the more basic a compound, the less acidic. While it is difficult to quantitatively rank acids or base strengths, we can note some structural differences and effects that can drastically affect these strengths. This can allow us to predict the relative acidities or basicities of two different compounds. Commonly, the effects are related or directly linked to stabilization of a conjugate, thus the Bronsted-Lowry theory is mainly used in this part of the article, because it is more important here. The more stable the conjugate base or acid, the more energetically favorable the donation or acceptance of a proton by the acid or base, respectively, and this will be able to influence the rate of a reaction.
The first of these effects are field effects (Fig. 3), which we have already (discussed here). When electron-withdrawing groups are present in a molecule, they are able to increase its acidity (and decrease its basicity), while for electron-donating groups, the opposite is true. The reasoning behind this is that the electron-withdrawing group stabilizes the conjugate base while the electron-donating group destabilizes the conjugate base. Where a proton leaves, it is typical for the conjugate base to have a negative charge. The existence of this change, especially on an atom like carbon, is destabilizing. The electron-withdrawing groups present on the molecule are able to withdraw some electron density from the negative charge, stabilizing it. The field effect works through space, so it is affected by the proximity of the electron-withdrawing group from the negatively-charged portion of the molecule. This contrasts it with a very similar effect, the inductive effect, which works through the carbon backbone instead.
Fig. 3: Example of the field effect.
A cousin of the field effect is the inductive effect. For the inductive effect, it operates through bonds but not space, so for the inductive effect to work, the electron-withdrawing or electron-donating group has to be directly bonded to the atom which is giving away or accepting the proton. The effects are generally the same as that of the field effects, however they are obviously much stronger as there is both a closer proximity and a more direct connection with the negatively-charged or positively-charged atom. Where the group works by the inductive effect, it will also simultaneously work by the field effect, although the stronger effect of inductivity usually covers up the weaker field effect. Besides this, the effects for the inductive and field effects are the same: the presence of an electron-withdrawing substituent on the negatively-charged atom in the conjugate base A will increase the acidity of H-A, while the presence of an electron-donating substituent on the negatively-charged atom would instead decrease the acidity, and increase the basicity, of H-A.
There is also the resonance effect. The resonance effect refers to cases where a molecule’s structure is weighted over several resonance forms (Fig. 4). The actual structure of the molecule would be the resonance hybrid, a combination of all the resonance forms in the molecule. Because of this, where resonance exists in a molecule with a negative charge, and different resonance forms happen to move this negative charge to other atoms in the molecule, the negative charge will be stabilized. The reason behind this is that the negative charge is now delocalized over several atoms instead of concentrated at one atom, making it more stable. Another possible reason could be that the delocalization occurs to move some electron density over to a more electronegative atom, which can bear the negative charge better. We note that in the conjugate base of an acid, some electron density will be born by the donor atom. Where resonance exists, this negative charge can be delocalized, thus making the negatively-charged molecule more stable. Importantly, the resonance must exist only in the conjugate base, and not the acid, otherwise the reaction will not be energetically favorable as both the reactant and the product are stabilized. The resonance effect is the reason why compounds like carboxylic acid or phenols possess unexpectedly high acidities.
Fig. 4: Structure of the resonance forms of carbonate.
Part 1 of this article is here.