Introduction
Although it may seem obvious to many what an acid or a base is, it is more difficult to describe it in scientific terms; it is difficult to set a benchmark for what is considered an acid-base reaction and what is not, and this has given rise to many theories about what an acid or a base is, each with varying scopes. As of today, there are three widely accepted theories to explain acids and bases: the Bronsted-Lowry theory, the Lewis theory and the Arrhenius theory. Although they vary in scope, they do not disagree with each other; there is always a certain class of compounds which can fall into all three theories. In this article, we will be discussing in depth only two of the three theories, the Bronsted-Lowry and Lewis theories, as they are the more relevant theories in today’s chemistry. The reason why we will not explain in depth the Arrhenius theory is because it is very limited in scope, as we will explain later.
Arrhenius Theory
As we have mentioned earlier, the Arrhenius theory is one of the acid-base theories that are most limited in scope, because it defines acids as chemical species which dissociate to form H+ ions in solution, and bases as chemical species which dissociate to form OH- ions in solution. We will give some examples of Arrhenius bases and acids: a good example for an Arrhenius acid is the hydrogen halides, HX, which dissociate in solution to form H+ and X- ions. A good example for an Arrhenius base would then be any OH- salt, such as NaOH or Mg(OH)2, which dissociate to form the metal cation and the OH- anion. This is not to say, of course, that molecules without H+ or OH- substituents cannot be Arrhenius acids and bases, respectively. For example, NH3 is still considered an Arrhenius base, because in solution, it can react with water by taking away a H+ (proton) from H2O, resulting in NH4+ and OH-. As such, the OH- ion is still generated and NH3 can be considered an Arrhenius base. For acids, the protonated water molecule, H3O+, is considered as H+ because H+ cannot exist in solution. Below, we show the example of an Arrhenius acid-base reaction (Fig. 1).
Fig. 1: Arrhenius acid-base reaction.
Water is a suitable solvent to base the Arrhenius theory in, because it is amphoteric, meaning that it is (loosely) both acidic and basic, i.e. it can get protonated (from H2O to H3O+) or get deprotonated (from H2O to OH-). Of course, since the Arrhenius theory is defined as ‘in solution’, any solvent can be used. The Arrhenius theory has a terribly limited scope because of this same term. A requirement for all Arrhenius acids and bases is that they must be able to dissolve in solution, and of course there are some compounds which may be acidic or basic but nevertheless are unable to dissolve in water. This is also important where we consider an Arrhenius acid-base reaction, because in cases where only one of the reactants are able to dissolve, or that they are only able to dissolve in different solvents, reaction cannot occur and thus it can no longer be considered an acid-base reaction.
Bronsted-Lowry Theory
The next theory that we will discuss, and the first in detail, is the Bronsted-Lowry theory. This theory is relatively simple; it suggests that an acid is a proton donor and a base is a proton acceptor. Let us take a look at the definition of a proton acceptor first. The proton (H+) possesses no electrons, but a bond to be formed between the proton and the base must contain two electrons. This means that both electrons must come from the base. As such, it is typical for the base to contain an unshared pair of electrons (lone pair), which is sometimes present in a pi orbital. By this theory, we may consider the acid-base reaction as a proton transfer, with the acid transferring the proton to the base. We can also take note of one key thing: when the acid donates its proton, it will form a basic species, because after donating the proton, it will always be able to take its proton back to reform the original acidic species. The more acidic the Bronsted-Lowry acid, the more stable the resulting basic species. The basic species is known as a conjugate base. This also applies to the original basic species, which accepts the proton to form an acidic species, known as a conjugate acid (Fig. 2).
Fig. 2: Conjugate acids and bases.
The fine line between acids and bases can be very confusing in the Bronsted-Lowry acid-base theory. Typically, acid-base reactions would be considered reactions between an acidic compound and a basic compound. But what is the ‘line’ between acids and bases? How does one determine whether a substance is basic or not? Where a stronger acid interacts with a weaker acid, the stronger acid will end up protonating the weaker acid, since it has the stronger tendency to lose the proton, thus the weaker acid has acted as a base in the reaction. But is it really a ‘base’? To clear up this confusion, a quantitative treatment of the strengths of acids and bases is needed. This comes in the form of a pKa table, where each acid or base is assigned a value based on how strongly acidic or basic they are. The pKa table is not absolute because the quantities vary based on some variables, such as temperature or between different states (i.e. solid state vs gaseous state).
By the pKa table, we note that water, which is amphoteric, has a pKa of 14. Below that, we have alcohols, with a pKa of 17; we note that it is quite common to see the alcohol being protonated to form RH2O+, a better leaving group (H2O). Below the alcohol, we have the ketones and aldehydes, which both have the acidic alpha-proton. The acidity comes from keto-enol tautomerism (Fig. 3); when the alpha-proton is lost, the negative charge is still relatively stable because of the existence of resonance forms involving enols. The resonance form delocalizes the negative charge on the carbanion over to the more electronegative oxygen, which is able to bear the negative charge well. This makes the alpha-carbon relatively acidic, more than to be expected from other carbons with carbon-hydrogen bonds. For example, alkanes have a pKa of 50.
Fig. 3: Keto-enol tautomerism of acetone.
Lewis Theory
Next, we will be discussing the Lewis theory of acids and bases. This theory is becoming increasingly more relevant over the Bronsted-Lowry acid-base theory, because one of its key applications, the Hard-Soft Acid-Base (HSAB) theory, is very useful in chemical calculations. Let us first proceed with the definitions of Lewis acids and bases. A Lewis acid is defined as any species with a vacant orbital, while a Lewis base is defined as any species with an available pair of electrons for electron donation. The Lewis acid-base reaction is thus different from the Bronsted-Lowry acid-base reaction in that the Lewis acid-base involves a transfer of electrons, while the Bronsted-Lowry involves a transfer of a proton. In the Lewis case, there must be some degree of interaction between the vacant orbitals of the Lewis acid and the lone pairs of the Lewis base. A common example of a Lewis acid are the boron trihalides (Fig. 4), while an example of a Lewis base are simple anions or any other molecule containing an atom with a lone pair.
Fig. 4: Structure of a boron trihalide.
We should note that the Lewis theory is not completely different from the Bronsted-Lowry theory. By the Lewis theory, a Bronsted-Lowry base will always be a Lewis base (both have a pair of electrons that can be formed to the proton, as we have mentioned before), while all Bronsted-Lowry acids are Lewis acids. The similarity between the Lewis base and the Bronsted-Lowry acid is less clear here, because the Lewis acid has a vacant orbital while it may seem that the Bronsted-Lowry acid does not always have a vacant orbital. The answer to this lies in the fact that the two theories treat the definition of the acid differently. In the Bronsted-Lowry theory, the whole molecule containing the acidic hydrogen is treated as the acid. However, in the Lewis theory, only the proton is treated as the acid. The proton has the vacant orbital, in which it can accept electrons to form a bond with the Lewis base. And as such, this allows us to demonstrate the similarities between the Lewis theory and the Bronsted-Lowry theory.
The Lewis theory has a noticeably large scope because many reactions can be classed as Lewis acid-base reactions. For example, nucleophilic substitution involves a nucleophile and a substrate (or an electrophile, but this can be relative just like with acids and bases). In the reaction, the nucleophile attacks the atom of the substrate with a vacant orbital, and uses its unshared pair of electrons to form a bond with the substrate. By the Lewis theory, the nucleophile is the Lewis acid while the substrate (or electrophile) is the Lewis base. As such, a nucleophilic substitution reaction can be considered a Lewis acid-base reaction. Because of the large scope of the Lewis theory, there is a lack of measurement of Lewis acid strength, although approximations have been made: AlX3 > FeX3 > SnX4 > AsX5 > ZnX2 > HgX2. Since the approximation of Lewis acid strength is lacking, we may expect that it would be difficult to calculate the feasibility or rate of a Lewis acid-base reaction. However, since Lewis acid-base reactions can include nucleophilic substitutions, the feasibility or rate depends on a separate, unique quality, the hardness or softness of an acid or base. This forms the basis of the theory, hard-soft acid-base (HSAB) theory.
Hard-Soft Acid-Base (HSAB) Theory
Under HSAB theory, acids and bases fall into four types: hard acids, soft acids, or hard bases, soft bases. Soft bases refer to cases where the donor (basic) atom has a small electronegativity and high polarizability. Hard bases refer to cases where the donor atom has a high electronegativity and a low polarizability. Polarizability is the exact opposite of electronegativity; as such, the elements on the top right of the periodic table will be the least polarizable. As a result, soft bases are easier to oxidize than hard bases. Soft acids are large, have low positive charges and have high polarizability. Hard acids are small, have higher positive charges and have low polarizability. Let us attempt to classify H+, the proton. It is obviously a Lewis acid, since it has vacant orbitals. At the same time, it is very small, and thus not polarizable at all. Thus, we would classify H+ as a hard acid.
Why is this theory highly useful? The reason for this is because of the rule that hard acids prefer hard bases, forming ionic bonds, while soft acids prefer soft bases, forming covalent bonds. This is because hard bases have high electronegativities, making them withdraw more electron density from the hard acid it forms a bond with, causing the bond to have an increased ionic character and to be classified as ionic. The opposite is true for soft bases thus there would be less ionic character. The application of the theory is shown in coordination chemistry, but this will be discussed more in the next part of this article.
Oh by the way, articles will be published every three days instead of every two days from now on. Part 2 of this article will be released on 8 Nov here.