Hello, and welcome to your chemistry lesson for Class 10. Today, we are going to explore an important compound: hydrogen chloride. This is a fascinating substance that bridges the world of gases and acids, and by the end of this lesson, you will understand how we prepare it, why it behaves the way it does, and how it transforms into hydrochloric acid.
Let us begin with the basics. Hydrogen chloride has the molecular formula HCl, with a molecular mass of 36.5 atomic mass units. It is a covalent compound, formed when hydrogen and chlorine share a pair of electrons. In its pure state, it exists as a colourless gas with a sharp, pungent odour that can irritate your nose and throat.
Now, how do we actually make hydrogen chloride in the laboratory? The standard method uses sodium chloride and concentrated sulphuric acid.
Here is what happens. When you heat common salt, which is sodium chloride, with concentrated sulphuric acid below 200 degrees Celsius, you get sodium hydrogen sulphate and hydrogen chloride gas. The reaction is: NaCl plus H₂SO₄ gives NaHSO₄ plus HCl gas, at temperatures below 200 degrees Celsius.
If you continue heating above 200 degrees Celsius, the sodium hydrogen sulphate reacts with more sodium chloride to form sodium sulphate and additional hydrogen chloride. When heated above 200 degrees Celsius, 2 NaCl reacts with H₂SO₄ to form Na₂SO₄ and 2 molecules of HCl gas.
Notice why we use sulphuric acid and not nitric acid. Nitric acid is volatile, so it would escape along with the hydrogen chloride gas, contaminating our product. Sulphuric acid stays put because it has a high boiling point.
Once we have generated the gas, we need to collect it properly. Hydrogen chloride is about 1.28 times heavier than air, with a vapour density of 18.25 compared to air's 14.4. Therefore, we use downward delivery, also called upward displacement of air. We cannot collect it over water because it dissolves far too readily. One volume of water dissolves 452 volumes of the gas at room temperature.
Before collection, the gas must be dried. We pass it through concentrated sulphuric acid. We cannot use phosphorus pentoxide or quick lime as drying agents because they actually react with hydrogen chloride. Phosphorus pentoxide forms phosphorus oxychloride and metaphosphoric acid. Calcium oxide, being basic, forms calcium chloride and water.
Here is a beautiful demonstration of hydrogen chloride's extreme solubility: the fountain experiment.
Imagine a dry round-bottomed flask completely filled with dry hydrogen chloride gas. We fit it with a rubber stopper containing two holes. Through one hole, we insert a long jet tube that dips into a beaker of blue litmus solution. Through the other hole, we place a dropper containing a few drops of water.
When we press the dropper, water enters the flask. Instantly, the hydrogen chloride gas dissolves, creating a partial vacuum inside. The higher atmospheric pressure outside forces the blue litmus solution up through the jet tube, and it sprays into the flask as a spectacular red fountain. The colour change to red confirms the acidic nature of the dissolved gas.
Now let us turn to the chemical properties of hydrogen chloride gas. First, it is neither combustible nor a supporter of combustion. A burning splint placed in it will be extinguished.
When heated above 500 degrees Celsius, it undergoes thermal dissociation into hydrogen and chlorine.
Above 500 degrees Celsius, 2 molecules of HCl dissociate into H₂ and Cl₂.
Active metals that lie above hydrogen in the electrochemical series react with hydrogen chloride when heated. Sodium, calcium, magnesium, zinc, and iron all form their chlorides and liberate hydrogen gas. For example, when heated, zinc reacts with hydrogen chloride to form zinc chloride and hydrogen gas.
One of the most striking reactions occurs with ammonia. When hydrogen chloride gas meets ammonia gas, they combine instantly to form dense white fumes of ammonium chloride. This is a classic test: NH₃ gas combines with HCl gas to form solid NH₄Cl. Two gases reacting to form a solid, a phenomenon you can actually see.
When hydrogen chloride dissolves in water, we get hydrochloric acid. This is where things get interesting from an electronic perspective.
In water, the polar covalent HCl molecule ionises to form hydronium ions and chloride ions. The hydronium ion, H₃O⁺, is what makes the solution acidic.
Here is a crucial point: dry hydrogen chloride gas, or even liquefied hydrogen chloride, does not contain free ions. Therefore, dry gas does not turn blue litmus red, and it does not conduct electricity. Only in aqueous solution does it show acidic properties.
Similarly, when dissolved in organic non-polar solvents like toluene, hydrogen chloride remains covalent and molecular. It does not ionise, so it neither affects litmus nor conducts electricity. This behaviour confirms its covalent nature.
Let us look at how we actually prepare hydrochloric acid in the laboratory. We cannot simply bubble the gas directly into water because of a serious problem called back suction.
Because hydrogen chloride is so soluble, water would rush back up the delivery tube into the generating flask.
This is dangerous because the water would react violently with the hot concentrated sulphuric acid present there.
The solution is elegant: we use an inverted funnel arrangement. The funnel is placed just touching the water surface in a trough. As the gas dissolves, water rises slightly into the funnel, but then an air gap forms between the funnel rim and the water surface. This equalises pressure, and the water falls back. The cycle repeats until the water becomes saturated.
The funnel serves two purposes: it prevents back suction, and it provides a large surface area for efficient absorption. Sometimes, an empty flask is placed between the generator and the water trough as an additional safety measure, an anti-suction device.
When we try to concentrate dilute hydrochloric acid by boiling, we encounter something peculiar. The concentration rises only until it reaches about 22.2 percent hydrogen chloride by weight. At this point, the solution contains 77.8 percent water. The boiling point becomes 110 degrees Celsius. The mixture becomes a constant boiling mixture or azeotrope, which boils without change in composition. No matter how long you boil it, the composition stays the same because hydrogen chloride gas escapes along with water vapour.
Hydrochloric acid shows all the typical reactions of a strong acid.
With metals above hydrogen in the activity series, it produces metal chlorides and hydrogen gas. With metal oxides and hydroxides, it undergoes neutralisation to form salts and water. For example, black copper oxide reacts with hydrochloric acid to form a blue solution of copper chloride and water.
With carbonates and hydrogen carbonates, it liberates carbon dioxide with brisk effervescence. With sulphites, it gives sulphur dioxide. With sulphides, it releases hydrogen sulphide gas with its characteristic rotten egg smell.
Thiosulphates react differently. With dilute hydrochloric acid, they produce sulphur dioxide and precipitate yellow sulphur. This distinguishes them from sulphites, which do not give a sulphur precipitate.
Now for some important precipitation reactions. Silver nitrate solution reacts with hydrochloric acid to give a thick, curdy white precipitate of silver chloride. The reaction is: AgNO₃ plus HCl gives AgCl precipitate plus HNO₃.
This silver chloride precipitate is insoluble in nitric acid but dissolves in ammonium hydroxide to form a complex salt. The formula is [Ag(NH₃)₂]⁺Cl⁻, called diammine silver one chloride. When you add dilute nitric acid to this solution, the white precipitate of silver chloride reappears.
Silver chloride has a fascinating property: it is photosensitive. When exposed to light, it decomposes into silver metal and chlorine, turning from white to black. The fine black powder you see is metallic silver.
Lead nitrate and mercury one nitrate also give white precipitates with hydrochloric acid. Lead chloride and mercury one chloride form respectively. The precipitate of lead chloride is distinctive because it dissolves on heating.
Concentrated hydrochloric acid acts as a reducing agent with strong oxidising agents. When heated with manganese dioxide, lead dioxide, or red lead, it liberates chlorine gas. For example, when heated with concentrated hydrochloric acid, MnO₂ forms MnCl₂, water, and Cl₂ gas. The greenish-yellow chlorine gas evolved turns moist starch iodide paper blue-black.
Before we conclude, let us recap the essential points you need to remember.
First, hydrogen chloride is prepared in the laboratory by heating sodium chloride with concentrated sulphuric acid. Below 200 degrees Celsius, sodium hydrogen sulphate forms. Above 200 degrees Celsius, sodium sulphate forms. It is collected by downward delivery because it is 1.28 times heavier than air. It is dried using concentrated sulphuric acid.
Second, the fountain experiment dramatically demonstrates its extreme solubility in water, and the resulting solution is acidic.
Third, dry hydrogen chloride is covalent and non-acidic, but in water it ionises to form hydronium ions, showing acidic properties.
Fourth, hydrochloric acid is prepared by dissolving the gas in water using an inverted funnel to prevent dangerous back suction.
Fifth, hydrochloric acid reacts with metals above hydrogen in the activity series, their oxides, hydroxides, carbonates, hydrogen carbonates, sulphites, and sulphides. It also gives characteristic precipitates with silver nitrate and lead nitrate.
Sixth, concentrated hydrochloric acid reduces strong oxidising agents like manganese dioxide, lead dioxide, and red lead to liberate chlorine gas.
That brings us to the end of our lesson on hydrogen chloride and hydrochloric acid. You have learned about preparation, physical and chemical properties, and the subtle but crucial difference between the gaseous compound and its aqueous solution. Keep practising the equations, and pay special attention to the conditions and observations that distinguish one reaction from another. Until next time, stay curious and keep exploring the fascinating world of chemistry.