Oxoacids of Chlorine: An oxoacid is a compound that contains at least one oxygen, hydrogen, and no less than one other element. These have at least one hydrogen molecule bound to oxygen. This hydrogen can separate into the \({{\rm{H}}^{\rm{ + }}}\) cation and the anion of the respective acid. Halogens form oxoacids of different types; the general formula of such oxoacids are Hypohalous acid \(\left( {{\rm{HOX}}} \right)\), halous acid \(\left( {{\rm{HOXO}}} \right)\), halic acid \(\left( {{\rm{HOX}}{{\rm{O}}_{\rm{2}}}} \right)\) and perhalic acid \(\left( {{\rm{HOX}}{{\rm{O}}_{\rm{3}}}} \right)\). These oxoacids are stable only in aqueous solutions or in the form of their salts. The acidic strength of oxoacids increases with an increase in the oxidation number of halogens. Fluorine is an exception to this property as it forms only one oxoacid, i.e., HOF, due to the absence of vacant d-orbitals. 

Chlorine forms four types of oxoacids such as hypochlorous acid \(\left( {{\rm{HOCl}}} \right)\)  Chlorous acid \({{\rm{HOClO}}}\) or \({{\rm{HCl}}{{\rm{O}}_{\rm{2}}}}\), Chloric acid \({{\rm{HOCl}}{{\rm{O}}_{\rm{2}}}}\) or \({\rm{HCl}}{{\rm{O}}_{\rm{3}}}\) and perchloric acid \({\rm{HOCl}}{{\rm{O}}_{\rm{3}}}\) or \({\rm{HCl}}{{\rm{O}}_{\rm{4}}}\). Oxoacids of chlorine are very useful in the manufacture of chlorine, ammonium chloride and glucose (from corn starch). It is also used in extracting glue from bones and purifying boneblack as well as in medicine and as a laboratory chemical. In this article, we will learn about all oxoacids of chlorine and their properties, uses, structures, and stability in detail.

Four Oxoacids of Chlorine

Apart from hydrochloric acid \(\left( {{\rm{HCl}}} \right)\), all are oxoacids of chlorine. The central atom in the oxoacids of chlorine is \({\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybridised and has essentially one \({\rm{Cl – OH}}\) bond. Whereas most oxoacids of chlorine have \({\rm{Cl =O}}\) bonds present in them. Different oxoacids vary in their acidic strength, structures, and other properties. The four oxoacids of chlorine are:

1. Hypochlorous acid \(\left( {{\rm{HClO}}} \right)\)
2. Chlorous acid \({\rm{(HOClO}}\) or \({\rm{HCI}}{{\rm{O}}_{\rm{2}}})\)
3. Chloric acid \({\rm{(HOCl}}{{\rm{O}}_{\rm{2}}}\) or \({\rm{HCl}}{{\rm{O}}_{\rm{3}}})\)
4. Perchloric acid \({\rm{(HOCl}}{{\rm{O}}_{\rm{3}}}\) or \({\rm{HCl}}{{\rm{O}}_{\rm{4}}})\)

Let’s have a detailed discussion about each of these oxoacids of chlorine one by one:

1. Hypochlorous acid (HClO)

Preparation

When chlorine gas is passed through water, a hydrolytic disproportionation reaction takes place that leads to the formation of hydrochloric acid and hypochlorous acid. It is a reversible reaction, so the hydrochloric acid is then precipitated using mercuric oxide, and hypochlorous acid is extracted. The chemical reaction involved is given below:

Chemical Formula and Structure 

Hypochlorous acid is a simple molecule that contains oxygen in the centre which connects with the chlorine and hydrogen atoms through a single bond.

IUPAC name	Hypochlorous acid or chloric (I) acid or chloranol.
Chemical Formula	\({{\rm{HClO}}}\)
Molar mass	\({\rm{52}}{\rm{.46\, g\, mo}}{{\rm{l}}^{{\rm{ – 1}}}}\)
\({\rm{p}}{{\rm{K}}_{\rm{a}}}\) value	\(7.53\)
Solubility	Water-soluble
Salts of Hypochlorous acid	Hypochlorites \(\left( {{\rm{CI}}{{\rm{O}}^ – }} \right)\)

Uses

1. Hypochlorous acid is used as a disinfectant and sanitiser.
2. It heals the wound.
3. It is used for water treatment in swimming pools.
4. Hypochlorous acid is tested to remove up to \(99\% \) of foul odours, including garbage, rotten meat, toilet, stool, and urine.

2. Chlorous acid (\({\rm{HOClO}}\) or \({\rm{HCl}}{{\rm{O}}_{\rm{2}}}\))

Preparation

Chlorous acid is prepared through the reaction of barium or lead chlorite and dilute sulphuric acid. It is a powerful oxidising agent.

\({\rm{Ba}}{\left( {{\rm{Cl}}{{\rm{O}}_{\rm{2}}}} \right)_{\rm{2}}}{\rm{ + }}{{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{4}}} \to {\rm{BaS}}{{\rm{O}}_{\rm{4}}}{\rm{ + 2HCl}}{{\rm{O}}_{\rm{2}}}\)

\({\rm{Pb}}{\left( {{\rm{Cl}}{{\rm{O}}_{\rm{2}}}} \right)_{\rm{2}}} + {{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{4}}} \to {\rm{PbS}}{{\rm{O}}_{\rm{4}}}{\rm{ + 2HCl}}{{\rm{O}}_{\rm{2}}}\)

Chemical Formula and Structure 

Chlorous acid is an inorganic compound having a weak acidic character, it contains one \({{\rm{CI = O}}}\) bond and another oxygen connect with the chlorine and hydrogen atoms through a single bond.

Properties

IUPAC name	Chlorous acid
Chemical Formula	\({\rm{HCI}}{{\rm{O}}_{\rm{2}}}\)
Molar mass	\({\rm{68}}{\rm{.46\, g\, mo}}{{\rm{l}}^{{\rm{ – 1}}}}\)
\({\rm{p}}{{\rm{K}}_{\rm{a}}}\) value	\(1.96\)
Solubility	Water-soluble
Salts of Chlorous acid	Chlorite

Uses

1. Chlorous acid acts as an effective bleaching agent. Because of this, it is used for the removal of lignin and non-carbohydrate components from wood pulp without any action on carbohydrates. 
2. The sodium salt of chlorite ion is used in the industrial production of chlorine dioxide.

3. Chloric acid (\({\rm{HOCl}}{{\rm{O}}_2}\) or \({\rm{HCl}}{{\rm{O}}_3}\))

Preparation

Chloric acid is prepared by the reaction of sulphuric acid with barium chlorate; later on, the insoluble barium sulphate is removed by precipitation:

\({\rm{Ba}}{\left( {{\rm{Cl}}{{\rm{O}}_3}} \right)_{\rm{2}}}{\rm{ + }}{{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{4}}} \to {\rm{BaS}}{{\rm{O}}_{\rm{4}}}{\rm{ + 2HCl}}{{\rm{O}}_3}\)

Another method for producing chloric acid and hydrogen chloride is by heating hypochlorous acid.

\({\rm{3HClO}} \to {\rm{HCl}}{{\rm{O}}_3} + {\rm{2HCl}}\)

Chemical Formula and Structure

Chloric acid is an oxoacid of chlorine that contains two \({{\rm{Cl= O}}}\) bonds and one \({{\rm{Cl -OH}}}\) bond along with a lone pair of electrons.

Properties

IUPAC name	Chloric (V) acid or Chloric acid
Chemical Formula	\({\rm{HCI}}{{\rm{O}}_{\rm{3}}}\)
Molar mass	\({\rm{84}}{\rm{.459\,g\, mo}}{{\rm{l}}^{{\rm{ – 1}}}}\)
\({\rm{p}}{{\rm{K}}_{\rm{a}}}\) value	Approx \(( -1)\)
Solubility	soluble in water
Salt of chloric acid	Chlorate

Uses

1. Chloric acid is used as a reagent in chemical analysis and to prepare other chemicals. 
2. It will accelerate the combustion of combustible materials, and most of them can be ignited on contact only.

4. Perchloric acid (\({\rm{HOCl}}{{\rm{O}}_3}\) or \({\rm{HCl}}{{\rm{O}}_4}\))

Preparation

Treatment of barium perchlorate with sulphuric acid gives the precipitate of barium sulphate, leaving behind perchloric acid. The chemical reaction is given below:

\({\rm{Ba}}{\left( {{\rm{Cl}}{{\rm{O}}_4}} \right)_{\rm{2}}}{\rm{ + }}{{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{4}}} \to {\rm{BaS}}{{\rm{O}}_{\rm{4}}}{\rm{ + 2HCl}}{{\rm{O}}_4}\)

Chemical Formula and Structure 

Perchloric acid is an oxoacid of chlorine that contains three \({{\rm{Cl= O}}}\) bonds and one \({{\rm{Cl – OH}}}\) bond.

Properties

IUPAC name	Perchloric acid or Chloric (VII) acid or Hyperchloric acid
Chemical Formula	\({\rm{HCI}}{{\rm{O}}_{\rm{4}}}\)
Molar mass	\({\rm{100}}{\rm{.46\,g\,mo}}{{\rm{l}}^{{\rm{ – 1}}}}\)
\({\rm{p}}{{\rm{K}}_{\rm{a}}}\) value	\(-10\)
Solubility	Water-soluble
Salt of Perchloric acid	Perchlorate

Uses

1. Perchloric acid is used in rocket fuel as it is mainly produced as a precursor to ammonium perchlorate.
2. Perchloric acid is one of the most proven materials for cutting metals to create designs or a pattern in them, which is used in critical electronics applications as well as for ore extraction.
3. Perchloric acid may also be called a superacid; it is one of the strongest Bronsted–Lowry acids. 
4. It is used as an eluent in ion-exchange chromatography.
5. It is used for electropolishing or etching metals like aluminium, molybdenum, etc.

Stability of Oxoacids of Chlorine

The stability order of the oxoacids of chlorine are:

\({\rm{HOCI < HCl}}{{\rm{O}}_2} < {\rm{HCl}}{{\rm{O}}_{\rm{3}}}{\rm{ < HCl}}{{\rm{O}}_{\rm{4}}}\)

Perchloric acid is the most stable oxoacid of chlorine because it has a maximum number of resonating structures of its conjugate base. It has the lowest oxidising power too. 

Acidic Strength of Oxoacids of Chlorine

The acidic strength of oxoacids of chlorine depends upon the oxidation number. As the oxidation number increases, the acidic character also increases. The increasing order of their acidic strength is:

\({\rm{HOCI < HCl}}{{\rm{O}}_2} < {\rm{HCl}}{{\rm{O}}_{\rm{3}}}{\rm{ < HCl}}{{\rm{O}}_{\rm{4}}}\)

Summary

In short, we can say that an oxoacid is a compound that contains at least one oxygen, hydrogen, and one another element. One hydrogen atom is attached to the oxygen, and this hydrogen can dissociate into the hydrogen cation and anion of the acid in an aqueous solution. Most of the halogens form oxoacids, such as Hypohalous acid, halous acid, halic acid, and perhalic acid. As an exception to other halogens, Fluorine forms only one oxoacid, i.e., \({\rm{HOF}}\), due to the absence of vacant d-orbitals. 

Chlorine forms four types of oxoacids such as hypochlorous acid \({\rm{HOCl}}\), chlorous acid \({\rm{HOClO}}\) or \({\rm{HCl}}{{\rm{O}}_2}\), chloric acid \({\rm{HOCl}}{{\rm{O}}_2}\) or \({\rm{HOCl}}{{\rm{O}}_3}\) and perchloric acid \({\rm{HOCl}}{{\rm{O}}_3}\) or \({\rm{HOCl}}{{\rm{O}}_4}\). Oxoacids of chlorine are prepared both in the laboratory and on an industrial scale. Some important uses of these oxoacids are: Hypochlorous acid is used as disinfectant and sanitiser, chlorous acid is an effective bleaching agent that is used in the paper industry, chloric acid and perchloric acid is highly flammable and used in rocket fuels, in etching of metals, etc. Perchloric acid is a very dangerous explosive and needs to be handled with care.

Frequently Asked Questions (FAQs) 

Q.1. How many oxoacids are formed by chlorine?
Ans. There are four oxoacids of chlorine; these are hypochlorous acid \({\rm{(HOCl)}}\), chlorous acid \({\rm{HOClO}}\) or \({\rm{HCl}}{{\rm{O}}_2}\), chloric acid \({\rm{HOCl}}{{\rm{O}}_2}\) or \({\rm{HCI}}{{\rm{O}}_3}\) and perchloric acid \({\rm{HOCl}}{{\rm{O}}_3}\) or \({\rm{HCl}}{{\rm{O}}_4}\).

Q.2. What are the formulas of oxoacids of chlorine?
Ans: The chemical formulas of four oxoacids of chlorine are: Hypochlorous acid \(\left( {{\rm{HOCl}}} \right)\), Chlorous acid \({\rm{HOClO}}\) or \({\rm{HCl}}{{\rm{O}}_2}\), Chloric acid \({\rm{HOCl}}{{\rm{O}}_2}\) or \({\rm{HCl}}{{\rm{O}}_3}\) and Perchloric acid \({\rm{HOCl}}{{\rm{O}}_3}\) or \({\rm{HCl}}{{\rm{O}}_4}\).

Q.3. What are the different oxoacids formed by chlorine, and what is the oxidation state of chlorine in each of them?
Ans: There are four different types of oxoacids of chlorine. The oxidation state of chlorine in hypochlorous acid \(\left( {{\rm{HOCl}}} \right)\) is \(+1\), the oxidation state of chlorine in chlorous acid \({\rm{HOClO}}\) or \({\rm{HCI}}{{\rm{O}}_2}\) is \(+3\), the oxidation state of chlorine in chloric acid \({\rm{HOCl}}{{\rm{O}}_2}\) or \({\rm{HCI}}{{\rm{O}}_3}\) is \(+5\), and the oxidation state of chlorine in perchloric acid \({\rm{HOCl}}{{\rm{O}}_3}\) or \({\rm{HCl}}{{\rm{O}}_4}\) is \(+7\).

Q.4. What are the oxoacids of halogen?
Ans: Halogens form oxoacids that are of different types; the general formula of such oxoacids are Hypohalous acid \(\left( {{\rm{HOX}}} \right)\), halous acid \(\left( {{\rm{HOXO}}} \right)\), halic acid \(\left( {{\rm{HOX}}{{\rm{O}}_2}} \right)\) and perhalic acid \({{\rm{HOX}}{{\rm{O}}_3}}\). These oxoacids are stable only in aqueous solutions or in the form of their salts. The acidic strength of oxoacids increases with an increase in the oxidation number of halogens. Fluorine is an exception to this property as it forms only one oxoacid, i.e., HOF, due to the absence of vacant d-orbitals. 

Q.5. Which oxoacid of chlorine is the strongest acid?
Ans: Perchloric acid is the strongest acid among all the oxoacids of chlorine. It is highly reactive and explosive in nature. As the thermal stability of perchloric acid \(\left( {{\rm{HCl}}{{\rm{O}}_4}} \right)\) is very high, and its oxidising power will be the least.

Q.6. Which oxoacid of chlorine is most stable?
Ans: Perchloric acid is the most stable oxoacid of chlorine because it has a maximum number of resonating structures of its conjugate base. It has the lowest oxidising power too. 

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