Core Concepts
Acid strength depends on a variety of chemical factors, including electronegativity, atomic radius, and resonance. Keep reading to learn all about the chemistry behind acid strength.
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Strong vs Weak Acids
When choosing between acids for an experiment or other purpose, chemists tend to look at acid strength as the most important factor. In terms of strength, acids tend to fall into two broad groups: strong acids and weak acids.
Strong acids are defined by their rapid dissociation, quickly releasing a proton and forming a highly stable conjugate base. Weak acids, by contrast, release their protons more slowly and form a relatively unstable conjugate base. While the strong and weak acid categorization is often useful to quickly understand the general strength of an acid, you may want to instead compare strengths between acids. For this, you’ll want to look at an acid’s pKa.
The technical definition of pKa is the negative-log of an acid’s proton dissociation constant (Ka), which represents the proportion of the acid that releases a proton at chemical equilibrium. This means that pKa decreases as acid strength increases. In simpler terms, stronger acids have lower pKa values, sometimes lower than 0, while weaker acids have higher pKa values.
However, while knowing which acids are stronger than others is valuable in the chemistry lab, understanding the factors behind acid strength allow for more advanced mastery of acid-base chemistry. So let’s delve into the underlying chemistry of acid strength!
Binary Acid Strength: Atomic Radius
First, let’s consider binary acids, which only involve one atom bound to proton, before we move on to more complicated polyatomic ion acids. The most common binary acids include the haloacids, which involve a halogen bound to a proton, such as HF, HCl, HBr, and HI. The most important factor affecting binary acid strength is anionic atomic radius. As we already know, atomic radius follows a periodic trend: atomic radius increases as you move down the periodic table.
To judge the strength of a binary acid, we specifically consider the atomic radius of the conjugate base. Generally, stronger binary acids involve larger atomic radii. This is because conjugate bases with higher atomic radii can better stabilize the negative charge resulting from proton dissociation than bases with smaller radii. With respect to the haloacids, this means that HI is the strongest, since I– has the largest atomic radius of the halides. The trend follows that HBr is the next strongest, followed by HCl, with HF being the weakest haloacid.
Polyatomic Ion Acid Strength
WIth acids involving polyatomic conjugate bases, the atomic radius trend still holds when considering the proton-bound atom. Specifically, thiols have lower pKas than alcohols since sulfur has a greater atomic radius than oxygen.
However, most acids with polyatomic anions involve elements like carbon, nitrogen, and oxygen bound to the acidic proton. Since these elements are similarly-sized, other factors outweigh atomic radius in determining acid strength, namely bond polarity and resonance.
Bond Polarity in Acid Strength
Generally, stronger acids involve more polar bonds. When comparing different atoms bound to an acidic proton, bond polarity corresponds to the periodic trend of electronegativty. In particular, electronegativity increases as you move right across the periodic table. The more electronegative the proton-bound atom, the more polar the bond, and therefore the stronger the acid. This is because more electronegative atoms, similar to larger atoms, can better stabilize negative charge than less electronegative elements. Thus, since nitrogen is more electronegative than carbon, ammonia is a stronger acid than methane. Because oxygen is yet more electronegative, water is an even stronger acid and hydrofluoric acid is even stronger than water.
Bond polarity also determines strength when comparing acids with the same atom bound to an acidic proton, but with different inductive effects. In particular, when we compare methane with fluoroform, we find that the highly electronegative fluorines in fluoroform pull electron density on the C-H bond. This results in a more highly polarized bond, and thus a stronger acid relative to methane.
Resonance in Acid Strength
Continuing with the theme of better stabilized conjugate bases corresponding to stronger acids, resonance stabilization can also increase acid strength. Specifically, stronger acids have resonance stabilized polyatomic bases relative to non-resonance stabilized bases. We can observe this when we compare the acidity of propane and propene. The sp3 C-H bond in propene is more acidic than the terminal C-H bonds of propane because the resulting allylic anion widely distributes its negative charge. The propanyllic anion, by contrast, has a more concentrated negative charge on one atom, resulting in less stabilization.
We can see yet stronger acids when the conjugate base is not only resonance stabilized, but aromatic. Cyclopentadiene is a famously strong organic acid because its result and cyclopentadienyl anion forms a super stable aromatic ring.
Oxyacids
One special type of polyatomic ion acid which follows different acid strength rules is the oxyacids. Oxyacids are defined as an acid with a conjugate base of a central atom bound to exclusively to one or more oxygens. Acidic protons can be bound to one or many of these oxygens. Common oxyacids include nitric acid, sulfuric acid, hypochlorous acid, and carbonic acid.
Oxyacid strength is determined by two fundamental factors: the electronegativity of the central atom and the number of oxygens.
In particular, stronger oxyacids have more electronegative central atoms. This is because electronegative central atoms have an inductive effect that makes the O-H bond more polar. This trend explains why nitric acid is stronger than phosphorous acid and why chloric acid is stronger than bromic acid.
Further, stronger oxyacids have more oxygens. This results from having more highly electronegative oxygens involved in the molecular structure also has an inductive effect that makes the O-H bond more polar. This trend explains why sulfuric and nitric acid are stronger than sulforous and nitrous acid, respectively.
