Chemistry

10 Types of acids

Acids are chemical compounds that have a pH of less than 7 on the pH scale. They are sour to taste and turn blue litmus paper red. They generally dissociate to release proton(H+) in aqueous solutions. They are highly corrosive in nature, hence should be handled with care. Acids can be classified based on their chemical reaction, the strength of dissociation, the number of protons dissociated, and the origins of the acids. Based on these classifications acids can be divided into 10 types. In this article, we discuss the 10 different types of acids with relevant examples.

Types of acids

Types of acid based on chemical reaction

Arrhenius acid
Brønsted-Lowry Acid
Lewis acid

Types of acid based on origin

Mineral acid
Organic acid

Types of acids based on acid strength

Strong acid
Weak Acid

Types of acid based on the number of dissociated H+ ions

Monoprotic acid
Diprotic acid
Triprotic acid

Classification based on the chemical reaction

Acids can be differentiated based on the definition of acids given by different scientists. The definition was based on the reaction of the chemical substances with other chemicals or in aqueous solutions. The basic characteristics like sour taste, pH below 7, etc will remain the same for all acids.

1) Arrhenius Acid

These acids were defined by a Swedish chemist named Svante Arrhenius after whom these acids are named. His suggestion was to classify chemical compounds based on different kinds of ions formed in aqueous solutions. Based on this definition, acids are chemical compounds that increase the concentration of hydronium(H₃O⁺) ions in aqueous solutions. And similarly, Arrhenius bases are chemical compounds that furnish hydroxyl (OH-) ions in aqueous solutions.

For example, Hydrochloric acid(HCl) in the aqueous solution can break down to give us H+ ions and Cl- ions. The H+ ions interact with water molecules to form hydronium ions. [H₂O + H⁺ → H₃O⁺]

HCl + H₂O → H₃O⁺+ Cl^– [This reaction increases the concentration of hydronium ions in the solution]

Examples of Arrhenius acids

Hydrochloric acid (HCl)
Sulfuric acid (H2So4)
Hydrobromic acid (HBr)
Nitric acid (HNO3)
Acetic acid (CH3COOH)

2)Brønsted-Lowry Acid

This type of acid was defined independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. According to their theory, an acid can be defined as any chemical compound that can donate a proton to another molecule during a reaction. The proton acceptor will be called a Brønsted-Lowry base. After the exchange, the acid will form a conjugate base and the molecule will form the conjugate acid.

So in a reaction between an acid and a base:

Acid + Base ⇌ Conjugate Base + Conjugate Acid

For example,

NH⁺⁴ + H₂O → NH₃ + H3O⁺

In the above reaction, ammonium ion (NH⁺⁴) is above to donate a proton to water(H₂O). So ammonium ion can be considered as a Bronsted-Lowry acid. This is similar to Arrhenius acid. The water molecule in this reaction accepts a proton for hydronium ions(H3O⁺). So, water will be called a Brønsted-Lowry base. The ammonia and hydronium ion will be called conjugate acid and conjugate base respectively.

Examples of Brønsted-Lowry acids

Hydrochloric acid (HCl)
Sulfuric acid (H2So4)
Hydrobromic acid (HBr)
Acetic acid (CH3COOH)
Carbonic acid (H2CO3)
Ammonium ion (NH₄⁺)

What is the difference between Arrhenius acid and Brønsted-Lowry acids?

Arrhenius acids and Bronsted-Lowry acids give a similar definition to acids as chemical compounds able to donate protons(H+). But Arrhenius defines it only for reactions in aqueous solutions. Whereas Brønsted-Lowry acids cover reaction with any molecule.

let us look at this example below to understand this further:
HCl +NH₃ →NH⁺⁴ + Cl⁻

In this example, hydrochloric acid(HCl) donates a proton to ammonia(NH3). So in this case HCl will not be called Arrhenius acid, because this is a reaction with water. Proton is donated to another molecule(NH3). So in this reaction, hydrochloric acid will be the Bronsted-Lowry acid and ammonium ion (NH⁺⁴) will be the conjugate acid.

3)Lewis Acid

Any chemical compound having an empty orbital, which accepts a pair of a non-bonded electrons during a chemical reaction can be defined as a Lewis acid. This definition was proposed by Gilbert N. Lewis in 1923. In the same manner, we can define a lewis base as a chemical that donates electron pair to another molecule during a chemical reaction. After accepting the electron pair the Lewis acid forms a Lewis adduct. Lewis acids are electrophilic and have low energy orbitals.

Ions like H3O+ and H+ can be considered as lewis acids as they can accept a pair of electrons during neutralization. Metal cations coordinating with ligands can be called Lewis acids, as they have empty orbitals to accept electrons(eg: Mg2+, Ca2+, etc). Many metal-organic complex frameworks and zeolites have many lewis acid sites in their framework.

Examples of Lewis acids

Borane (BH3)
Ethylaluminium sesquichloride [(C2H5)2AlCl*Cl2AlC2H5]
Hydrogen ion (H+)
Hydronium ions (H3O+)
Magnesium(II) ion (Mg2+)

Application of Lewis acids

Metal-organic frameworks like MIL-101(Cr) are used in catalyst reactions. Such metal-organic frameworks are made up of metal cations coordinating with ligands, and have many Lewis acid sites, making them a favorite for catalyst reactions. You can read this article for more details.

Lanthanide (Ln)-containing polyoxometalates have been studied as Lewis acid catalysts for cyanosilylation of benzaldehydes. You can read this article for more details.

You can read this article titled “A DFT study for catalytic deoxygenation of methyl butyrate on a Lewis Acid site of ZSM-5 zeolite” to understand the catalyst mechanism by Lewis acid zeolitic catalysts.

What is the difference between Brønsted-Lowry acids and Lewis acids?

Brønsted acid-base reactions are based on proton transfer from acid to base. But, on the other hand, Lewis acids are based on electron transfer from Base to acid. Most of the Lewis acids don’t qualify as Brønsted acid, as such compounds can not furnish H+ ions like for example Al(3+).

Classification based on the origin of acids

1) Mineral acids

Mineral acids are acids derived from inorganic compounds. Common acids like HCl, HNO3, and H2SO4 are all mineral acids. These include mono-di or tri protic acids. These acids are highly corrosive and should be handled with care. The application of mineral acids is huge across various industries. Since these acids are inorganic in nature, they will generally lack carbon. They are generally soluble in water and insoluble in most organic solvents. They include both strong acids (like Fluoroantimonic acid) and weak acids (like fluoric acid).

Examples of mineral acids

Chlorosulphonic acid (HClSO₃)
Hydroiodic acid (HI)
Perchloric acid (HClO₄)
Hydrochloric acid (HCl)
Selenic acid (H₂SeO₄)
Hydrofluoric acid (HF)
Hyposelenic acid (H₂SeO₃)
Nitric acid (HNO₃)
Fluorophosphoric acids (HF₆PF)
Boric acid (H₃BO₃)

2) organic acids

Organic acids are acids that are organic in nature(contain hydrocarbons). A very famous example of organic acids is acetic acid with the formula CH3COOH. The groups attached to the hydrocarbon attribute the acidic character to these acids. For example, in acidic acid, the Cooh group gives it an acidic nature. Similarly, groups like, -OH, -SO2OH, -SH, etc can provide acidic character to organic molecules.

Organic acids are found in many fruits, vegetables, and also in our bodies. These acidic give the characteristic taste and smell to many food items. For example, citric acid is found in citrus fruits, lactic acid is found in milk, malic acid is found in tomato, tartaric acid is found in turmeric, etc.

You can read this article titled “Taste and palatability of strawberry jam as affected by organic acid content” to understand how the taste of items is affected by the organic acid content in them.

Examples of organic acids

Citric acid (C6H8O7)
Lactic acid (C3H6O6)
Acetic acid (CH3COOH)
Malic acid (C4H6O5)
Tartaric acid (C4H6O6)
Formic acid (HCOOH)
Uric acid (C5H4N4O3)
Oxalic acid (C2H2O4)
Phthalic acid (C6H4(CO2H)2)
Oxaloacetic acid (C4H4O5)
Succinic acid (C4H6O4)
Gluconic acid (C6H12O7)

Classification based on acid strength

1) Strong acid

A strong acid undergoes complete dissociation in aqueous solutions. It means that it can donate a proton easily in aqueous solutions. The strength is determined based on the ability of the acid to donate H+ ions freely. The H+ ions then react with water to form hydronium ions. A common example of strong acid would be hydrochloric acid.

HCl + H₂O → H₃O⁺+ Cl^–

Strong acids have a large dissociation constant (Ka). They have a small logarithmic constant (pKa) value, generally less than -2.

Examples of strong acids

Fluoroantimonic acid (SbHF6)
Perchloric acid (HClO4)
Triflic acid [H(CF₃SO₃)]
Fluorosulfuric acid (H[FSO₃])
Carborane superacid [H(CHB₁₁Cl₁₁)]
Hydrochloric acid (HCl)

2) Weak acid

A weak acid undergoes incomplete dissociation in aqueous solutions. It means that it cannot donate a proton easily as compared to the strong acids. Weak acids are common than strong acids because it is very difficult for complete dissociation. They usually show a pH value between 3-6. Acetic acid (CH3COOH) is one of the most commonly used weak acids. Most of the organic acids can be included in the list of weak acids.

For example, let us consider the dissociation of formic acid:

HCOOH + H₂O ⇆ HCOO^– + H₃O⁺

The reaction can proceed both ways. The backward reaction will be more favored in the case of a weak acid. So, they have a lower value of dissociation constant (Ka). But, there will be some forward reaction also but less favorable compared to the backward reaction.

Are weak acids actually weak in nature?

The strength of an acid is an indication of the ability to dissociate H+ ions. Although the name suggests that these acids are weak, but they can be quite acidic in their concentrated form and should be handled with care. For example, acetic acid in its concentrated form can attain a pH of around 2.4.

Examples of Weak acids

Trichloracetic acid (CCl3COOH)
Hydrogen sulfide (H2S)
Formic acid (HCOOH)
Nitrous acid (HNO₂)
Methanoic acid (HCO2H)

Classification based on the number of dissociated H+ ions

Monoprotic acids

Monoprotic acids can donate only one proton(H+) ions.

HBr + H₂O → H₃O⁺+ Br^– [Hydrobromic acid is able to donate a single H+ ion]

Examples of Monoprotic acids

Hydrochloric acid (HCl)
Nitric acid (HNO₃)
Acetic acid (CH3COOH)
Hydroiodic acid (HI)
Hydrobromic acid (HBr)

Diprotic Acids

Diprotic acids can donate two protons(H+) ions.

H2SO4 + 2H2O → 2H₃O⁺ + SO4 – [Sulfuric acid is able to donate two H+ ions]

Examples of Diprotic acid

Sulfuric acid (H₂SO₄)
Carbonic acid (H₂CO₃)
Chromic acid (H2CrO4)
Glycine (C₂H₆NO₂)
Oxalic acid (C₂H₂O₄)

Triprotic acids

Triprotic acids (also called polyprotic along with diprotic acids) can donate three protons(H+) ions.

H₃PO₄ + 3H₂O ⇌ 3H₃O⁺ + H₂PO₄⁻ [phosphoric acid is able to donate three H+ ions]

Examples of Triprotic acids

Arsenic acid (H₃AsO₄)
Citric acid (C₆H₈O₇)
Phosphoric acid (H₃PO₄)