Chlorine and its Compounds: Properties, Preparation, and Industrial Applications

Reaction with Sodium Hydroxide

The reaction between chlorine gas and sodium hydroxide solution is significant, particularly in the formation of specific salts. When chlorine gas is reacted with sodium hydroxide, the equation is represented as $3Cl_2(g) + 6NaOH(aq) \rightarrow NaClO_3(aq) + 5NaCl(aq) + 3H_2O(l)$. In this reaction, one of the primary products is sodium trioxochlorate(V). From an ionic perspective, the reaction is written as $3Cl_2(g) + 6OH^-(aq) \rightarrow ClO_3^-(aq) + 5Cl^-(aq) + 3H_2O(l)$. This demonstrates the disproportionation of chlorine in a basic medium, leading to both chloride and trioxochlorate(V) ions.

Tests for Chlorine Gas

Chlorine is identified physically as a greenish-yellow gas that possesses a sharp and irritating smell. Its presence in a laboratory setting can be confirmed through specific chemical actions on test papers, notably damp litmus paper and starch-iodide paper. In the damp litmus paper test, a piece of damp blue litmus paper is dropped into a gas jar containing the unknown gas. If the gas is indeed chlorine, the litmus paper will first turn pink due to its acidic nature in moisture and then become rapidly bleached to white because of the strong oxidizing power of chlorine.

In the starch-iodide paper test, a piece of damp starch-iodide paper is introduced to the unknown gas. If the paper turns dark blue, the presence of chlorine is confirmed. This color change occurs because chlorine is more reactive and displaces iodine from the iodide ions present on the paper. The displacement reaction is represented as $2KI(aq) + Cl_2(g) \rightarrow 2KCl(aq) + I_2(s)$. Once the iodine is liberated, it reacts with the starch in the paper to produce the characteristic dark blue color.

Industrial and Domestic Uses of Chlorine

Chlorine is characterized as a powerful germicide due to its strong oxidizing nature. This property makes it essential in the sterilization of water for domestic use and in the treatment of water in swimming pools. Industrially, it serves as a critical bleaching agent for various materials such as cotton, linen, and wood pulp used in paper manufacturing. It is also an indispensable reagent in the manufacture of several chemical products. This includes organic solvents like trichloromethane ($CHCl_3$), tetrachloromethane ($CCl_4$), trichloroethene ($C_2HCl_3$), and trichloroethanal ($CCl_3CHO$), which is also known as chloral. Trichloroethanal is particularly important for the manufacture of the insecticide D.D.T.

Further applications of chlorine involve the production of widely used plastics such as poly(chloroethene), commonly known as polyvinyl chloride or PVC, as well as synthetic rubber and hydrochloric acid. Chlorine is used in the preparation of potassium trioxochlorate(V), which is utilized in making matches and fireworks, and sodium trioxochlorate(V), which serves as a potent weed-killer. In the textile and laundry sectors, bleaching agents like bleaching powder and sodium oxochlorate(I) are widely utilized. Additionally, domesticated antiseptics often contain acidified sodium oxochlorate(I) solution.

Hydrogen Chloride: Origins and Preparation

Hydrogen chloride exists as a gas at standard temperature and pressure ($s.t.p.$). When this gas dissolves in water, it forms hydrochloric acid. Historically, pure hydrogen chloride was first prepared in 1772 by the scientist Priestley. He referred to it as marine-acid gas because it was originally synthesized from sea salt. Naturally, hydrogen chloride occurs in only trace amounts in the atmosphere as an industrial by-product. While it is classified as an air pollutant, it is readily washed out of the atmosphere by rain because of its extreme solubility in water.

In the laboratory, hydrogen chloride is prepared by the action of hot concentrated tetraoxosulphate(VI) acid on any soluble chloride, such as sodium chloride ($NaCl$). The reaction occurs in stages depending on the temperature. Initially, at lower temperatures, sodium hydrogentetraoxosulphate(VI) ($NaHSO_4$) is formed along with hydrogen chloride gas, represented by the equation $NaCl(s) + H_2SO_4(aq) \rightarrow NaHSO_4(aq) + HCl(g)$. As the temperature increases, the reaction proceeds to completion, where sodium hydrogentetraoxosulphate(VI) reacts with more sodium chloride to form sodium tetraoxosulphate(VI) ($Na_2SO_4$) and additional hydrogen chloride: $NaCl(s) + NaHSO_4(aq) \rightarrow Na_2SO_4(aq) + HCl(g)$. The overall preparation equation is $2NaCl(s) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + 2HCl(g)$. Ionically, this is expressed as $Cl^-(s) + H^+(aq) \rightarrow HCl(g)$.

Industrially, hydrogen chloride is manufactured in large quantities by the direct combination of hydrogen and chlorine gases. These gases are typically obtained from the electrolysis of brine. The process involves burning a jet of hydrogen in an atmosphere of chlorine, producing very pure hydrogen chloride gas according to the equation $H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$. This pure gas is then dissolved in water to produce hydrochloric acid.

Physical Properties of Hydrogen Chloride

Pure hydrogen chloride gas is colourless and possesses a sharp, irritating smell. It exhibits acidic properties in the presence of water, as evidenced by its ability to turn damp blue litmus paper red. The gas is approximately 1.25 times denser than air. One of its most notable characteristics is its high solubility in water; this is famously demonstrated in the fountain experiment (Fig. 21.5), where the high solubility of the gas in a flask causes an intake of water that turns red upon contact with the gas, indicating the formation of an acid. In moist air, hydrogen chloride forms misty fumes because it dissolves in atmospheric moisture to form tiny droplets of hydrochloric acid.

While hydrogen chloride is covalent and does not ionize in non-polar solvents like chloroform and toluene, it readily dissolves in them. Such non-polar solutions do not conduct electricity and show no acidic properties. However, in aqueous solution, the covalent $HCl$ molecules react with polar water molecules, ionizing to form chloride ions and oxonium ions. The equation for this ionization is $HCl(g) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$. The oxonium ions ($H_3O^+$) are responsible for the acidic behavior of the aqueous solution. Notably, hydrogen chloride gas neither burns nor supports combustion, and it will extinguish a burning splinter.

Chemical Properties and Reactions of Hydrogen Chloride Gas

Hydrogen chloride gas undergoes a direct combination reaction with ammonia. When a gas jar of hydrogen chloride is inverted over a jar of ammonia gas ($NH_3$), dense white fumes are immediately formed. These fumes consist of fine particles of solid ammonium chloride suspended in the air. The reaction is written as $HCl(g) + NH_3(g) \rightarrow NH_4Cl(s)$. This reaction is a common diagnostic test for identifying the gas.

Additionally, hydrogen chloride reacts with many metals, particularly when they are heated. This reaction results in the formation of the respective metal chloride and the liberation of hydrogen gas. For example, with zinc, the reaction is $Zn(s) + 2HCl(g) \rightarrow ZnCl_2(s) + H_2(g)$. In cases where a metal can form more than one type of chloride, the reaction with hydrogen chloride gas typically produces the lower chloride. A specific example is the reaction with iron, which yields iron(II) chloride: $Fe(s) + 2HCl(g) \rightarrow FeCl_2(s) + H_2(g)$.

Properties and Solutions of Hydrochloric Acid

Hydrochloric acid is formed by dissolving hydrogen chloride gas in water. When concentrated hydrochloric acid is heated, it boils and liberates fumes of hydrogen chloride gas, which causes the remaining solution to become more dilute. Conversely, when dilute hydrochloric acid is heated, water is lost through evaporation, causing the residual solution to become more concentrated. Eventually, both types of solutions reach a specific point where they contain approximately 20.24% hydrochloric acid. This specific composition is known as an azeotropic mixture or a constant boiling mixture, as further heating will not change the concentration of the solution.

In its aqueous form, hydrochloric acid exhibits the standard properties of acids. It reacts with metals that are more electropositive than hydrogen to liberate hydrogen gas. For instance, $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$, or ionically $Zn(s) + 2H^+(aq) \rightarrow Zn^{2+}(aq) + H_2(g)$. This property is utilized to separate mixtures, such as zinc and copper dust; since zinc is more electropositive, it reacts with the acid and dissolves, while the copper remains unchanged. It also reacts with bases and alkalis to produce a salt and water, such as $ZnO(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2O(l)$ and $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$. The ionic representation of the neutralization of an alkali is $OH^-(aq) + H^+(aq) \rightarrow H_2O(l)$.

Hydrochloric acid also reacts with trioxocarbonates(IV) and hydrogentrioxocarbonates(IV) to liberate carbon(IV) oxide gas. Examples include $2HCl(aq) + Na_2CO_3(s) \rightarrow 2NaCl(aq) + H_2O(l) + CO_2(g)$ and $HCl(aq) + NaHCO_3(s) \rightarrow NaCl(aq) + H_2O(l) + CO_2(g)$. The respective ionic equations are $2H^+(aq) + CO_3^{2-}(s) \rightarrow H_2O(l) + CO_2(g)$ and $H^+(aq) + HCO_3^-(s) \rightarrow H_2O(l) + CO_2(g)$. Furthermore, hydrochloric acid serves as a precipitating agent; adding it to silver trioxonitrate(V) or lead(II) trioxonitrate(V) solutions results in white precipitates of silver chloride ($AgCl$) or lead(II) chloride ($PbCl_2$). The ionic reactions are $Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)$ and $Pb^{2+}(aq) + 2Cl^-(aq) \rightarrow PbCl_2(s)$.

Analytical Tests for Hydrogen Chloride

If a gas is suspected to be hydrogen chloride, it is usually colourless, has an irritating smell, fumes in moist air, and turns blue litmus paper red. To confirm its presence, two specific tests are conducted. First, the ammonia test involves placing a glass rod dipped in ammonia solution at the mouth of the gas jar. If the gas is hydrogen chloride, dense white fumes of ammonium chloride ($NH_4Cl$) will form. Second, adding a few drops of silver trioxonitrate(V) solution to the gas jar and shaking it will produce a white precipitate of silver chloride if the gas is hydrogen chloride. Lead(II) trioxonitrate(V) can be utilized as an alternative to silver trioxonitrate(V) for this precipitation test.

Concentrated hydrochloric acid should not be used to acidify solutions of strong oxidizing agents like potassium tetraoxomanganate(VII). This is because the acid is easily oxidized by such agents, resulting in the liberation of chlorine gas. Instead, tetraoxosulphate(VI) acid is the preferred acidifying agent in these scenarios.

Industrial Applications of Hydrochloric Acid

Hydrogen chloride gas and its aqueous acid have numerous industrial and laboratory uses. The gas is used in the synthesis of chloroethene, which is a vital monomer in the production of various plastics. Hydrochloric acid is extensively used in the food and materials industry, including the manufacture of gelatin and glue from animal tendons and the conversion of starch into glucose. It also plays a role in the production of textiles and dyes.

In the metallurgical industry, hydrochloric acid is used for a process called pickling, which involves removing oxide layers from the surfaces of metals before they undergo electroplating. This technique is particularly common in the steel-making industry. Finally, within a laboratory setting, hydrochloric acid is a staple reagent used for chemical testing, quantitative and qualitative analysis, and the preparation of various other chemical compounds.