Hello, and welcome to today's chemistry lesson. Today, we are diving into Chapter Thirteen: Practical Work. This chapter is all about developing your skills in the laboratory — how to identify unknown substances, recognise different gases, distinguish between ions, and carry out systematic chemical analysis. By the end of this lesson, you will understand the key techniques for identifying gases by their preparation and properties, recognising cations and anions through specific tests, and distinguishing between common substances using simple chemical reactions.
Let us begin with the identification of gases. In practical chemistry, you will often need to recognise gases based on how they are prepared and what properties they display. We will examine ten important gases, looking at both their preparation methods and their characteristic tests.
First, consider hydrogen. This gas is prepared by adding dilute hydrochloric acid or dilute sulphuric acid to reactive metals such as magnesium, zinc, or iron. The reaction produces a salt and hydrogen gas. For example, zinc reacts with hydrochloric acid to form zinc chloride and hydrogen. Hydrogen is colourless, odourless, and neutral to litmus. The key test is the burning splint test: pure hydrogen burns with a pale blue flame, while a mixture of hydrogen and air produces a distinctive pop sound when ignited.
Next, oxygen. This gas can be prepared by heating higher metallic oxides such as red lead oxide, or by decomposing metal nitrates. When lead nitrate is heated, it breaks down into lead oxide, nitrogen dioxide, and oxygen. Oxygen is also colourless, odourless, and neutral to litmus. Its defining characteristic is that it rekindles a glowing splinter. Additionally, oxygen is absorbed by an alkaline solution of pyrogallol, turning it dark brown.
Carbon dioxide is an acidic gas. It is prepared by heating metallic carbonates — except sodium carbonate and potassium carbonate — or by adding dilute acid to any carbonate or hydrogen carbonate. When zinc carbonate is heated, it forms zinc oxide and carbon dioxide. Carbon dioxide turns moist blue litmus faint red. The most reliable test is passing the gas through lime water: the clear solution turns milky due to the formation of insoluble calcium carbonate. Interestingly, if you continue passing carbon dioxide, the milkiness disappears because the insoluble carbonate converts to soluble calcium hydrogen carbonate. Importantly, carbon dioxide has no effect on acidified potassium dichromate or potassium permanganate solutions.
Chlorine is a greenish-yellow gas with a sharp, pungent, choking odour. It is prepared by adding concentrated hydrochloric acid to strong oxidising agents such as manganese dioxide or red lead oxide. Chlorine turns moist blue litmus red and then bleaches it completely. This bleaching action occurs because chlorine reacts with water to form hypochlorous acid, which releases oxygen that destroys colour. Another definitive test is with starch-iodide paper: chlorine turns this paper blue-black by liberating iodine. Chlorine also forms a white precipitate when passed through silver nitrate solution.
Hydrogen chloride is a colourless gas with a pungent, choking odour. It is prepared by adding concentrated sulphuric acid to a metal chloride such as sodium chloride, and gently heating. The gas turns moist blue litmus red. A characteristic test involves bringing a glass rod dipped in ammonia solution near the gas: dense white fumes of ammonium chloride form instantly. Like chlorine, hydrogen chloride forms a white precipitate with silver nitrate solution, but this precipitate dissolves in excess ammonium hydroxide.
Sulphur dioxide has a suffocating odour of burning sulphur. It is prepared by adding dilute acid to metallic sulphites. Sulphur dioxide turns moist blue litmus red and bleaches it. Like carbon dioxide, it turns lime water milky. However, you can distinguish sulphur dioxide from carbon dioxide using two important tests: it decolourises pink potassium permanganate solution, and it turns orange acidified potassium dichromate solution green. Carbon dioxide does neither of these.
Hydrogen sulphide is unmistakable due to its foul smell of rotten eggs. It is prepared by adding dilute acid to metallic sulphides such as iron sulphide or zinc sulphide. The gas turns moist blue litmus red. The definitive test uses lead acetate paper: hydrogen sulphide turns this paper black due to the formation of lead sulphide. It also blackens lead nitrate solution.
Ammonia is a basic gas with a sharp, pungent characteristic smell. It is prepared by heating an ammonium salt with an alkali such as sodium hydroxide or calcium hydroxide. Ammonia turns moist red litmus blue — the opposite of acidic gases. When a glass rod dipped in concentrated hydrochloric acid is brought near ammonia gas, dense white fumes of ammonium chloride appear. Ammonia also turns colourless Nessler's reagent, which is potassium mercuric iodide, brown.
Water vapour is released when hydrated salts are heated. The vapour condenses as droplets on cooler parts of the test tube. This liquid is neutral to litmus. Two important tests confirm water: it turns white anhydrous copper sulphate blue, and it turns blue cobalt chloride paper pink.
Finally, nitrogen dioxide is a brown gas with an irritating, pungent odour. It is prepared by heating heavy metal nitrates such as copper nitrate or lead nitrate. The gas turns moist blue litmus red. It turns starch-iodide paper blue-black and changes green acidified ferrous sulphate solution to brown. Notably, sodium nitrate and potassium nitrate do not produce nitrogen dioxide when heated.
Now let us examine the action of heat on various substances, which provides valuable clues for identification.
When copper carbonate — a light green amorphous powder — is heated strongly, it turns black and releases a colourless, odourless gas. This gas extinguishes a burning splinter, turns lime water milky, and has no effect on acidified potassium dichromate or potassium permanganate. The residue is black copper oxide, and the gas is carbon dioxide. Therefore, the original substance was copper carbonate.
Similarly, zinc carbonate — a white amorphous solid — turns pale yellow when heated, then white on cooling. It too releases carbon dioxide. The residue is zinc oxide, which is yellow when hot and white when cold.
Metal nitrates behave differently. Zinc nitrate hexahydrate is a white crystalline deliquescent solid. On heating, it melts to a white sticky mass and gives off steamy vapours. Stronger heating produces reddish-brown nitrogen dioxide fumes and a gas that relights a glowing splinter — oxygen. The residue is pale yellow zinc oxide when hot, turning white when cold.
Copper nitrate hexahydrate is a bluish-green crystalline solid. It melts to a bluish-green mass with steamy vapours, then decomposes to black copper oxide with reddish-brown nitrogen dioxide and oxygen.
Lead nitrate is a heavy white crystalline solid that crumbles with a crackling noise when heated. It produces reddish-brown nitrogen dioxide and oxygen, leaving a residue that is reddish-brown when hot, yellow when cold, and stains the glass yellow. This residue is lead oxide.
Several general principles emerge from heating tests. Ammonium salts with alkali produce ammonia. Ammonium nitrate and ammonium chloride leave no residue. Ammonium dichromate decomposes to a greenish-grey residue of chromium oxide. Strong oxidising agents like lead dioxide, red lead oxide, and mercury oxide produce oxygen when heated. Carbonates and bicarbonates evolve carbon dioxide, except sodium and potassium carbonates. Hydrated salts release water vapour. Lead compounds typically yield lead monoxide, which is brown when hot and yellow when cold. Zinc compounds give zinc oxide, yellow when hot and white when cold. Copper compounds form black copper oxide.
Moving on to the recognition of substances by their physical characteristics.
Colour provides immediate clues. Blue or bluish-green substances suggest copper two plus ions. Light green indicates iron two plus. Yellow or yellowish-brown points to iron three plus. White or colourless substances may contain lead two plus, zinc two plus, calcium two plus, sodium plus, potassium plus, or ammonium ions.
Odour is equally telling. The smell of ammonia indicates ammonium ions. The smell of rotten eggs suggests sulphide ions. The smell of burning sulphur indicates sulphite ions.
Physical state also helps: amorphous salts often contain carbonate ions, while hygroscopic or deliquescent substances typically contain chloride or nitrate ions.
Now we turn to systematic identification of ions in solution.
Cations are identified using alkalies — sodium hydroxide or ammonium hydroxide. These produce characteristic coloured precipitates of metallic hydroxides. Always add the alkali slowly at first, drop by drop, because some precipitates dissolve in excess and you might miss them if you add too quickly.
With sodium hydroxide solution: calcium forms a white curdy precipitate that remains insoluble in excess. Lead forms a white chalky precipitate that dissolves in excess. Zinc forms a white gelatinous precipitate that dissolves in excess. Copper forms a pale blue precipitate that stays insoluble. Iron two plus forms a pale green precipitate that turns brown, remaining insoluble. Iron three plus forms a reddish-brown precipitate that is insoluble.
With ammonium hydroxide: calcium shows no precipitate. Lead forms a white chalky precipitate that stays insoluble. Zinc forms a white gelatinous precipitate that dissolves in excess. Copper forms a pale blue precipitate that dissolves in excess to give a deep blue solution. Both iron two plus and iron three plus give insoluble precipitates — pale green turning brown for iron two plus, and rust brown for iron three plus.
The ammonium ion has special tests. When caustic alkali is added to any ammonium salt, ammonia gas is evolved. Additionally, Nessler's reagent turns brown in the presence of ammonium ions.
Anions are tested by reaction with acids or specific reagents. Adding dilute sulphuric acid: carbonates produce brisk effervescence of carbon dioxide, which extinguishes a burning splinter and turns lime water milky. Sulphides produce hydrogen sulphide with its characteristic rotten egg smell, turning lead acetate paper black. Sulphites produce sulphur dioxide with its suffocating odour, turning acidified potassium dichromate paper from orange to green.
Adding concentrated sulphuric acid: chlorides evolve hydrogen chloride gas, which forms dense white fumes with ammonia and can be confirmed with manganese dioxide to produce chlorine, or with silver nitrate to give a white precipitate soluble in ammonium hydroxide. Nitrates produce reddish-brown nitrogen dioxide fumes, confirmed by the brown ring test with ferrous sulphate and concentrated sulphuric acid.
For sulphate ions, add barium chloride solution after acidifying with nitric acid: a white precipitate insoluble in mineral acids confirms sulphate. Alternatively, use lead acetate with acetic acid: a white precipitate soluble in excess ammonium acetate also confirms sulphate.
Let us now consider how to distinguish between dilute acids and alkalis when both are colourless solutions.
Using indicators: acids turn blue litmus red, while alkalis turn red litmus blue. Methyl orange turns pink in acid and yellow in alkali. Phenolphthalein stays colourless in acid but turns pink in alkali.
Chemical tests provide further confirmation. Adding sodium carbonate to an acid produces carbon dioxide, which turns lime water milky. Adding ammonium carbonate to an alkali produces ammonia gas, detected by its smell and its effect on red litmus.
Two common black powders — copper oxide and manganese dioxide — can be distinguished using concentrated hydrochloric acid. When heated with concentrated hydrochloric acid, manganese dioxide evolves greenish-yellow chlorine gas, and the filtrate is brownish. Copper oxide does not evolve chlorine; its filtrate is bluish, and with ammonium hydroxide it forms a pale blue precipitate that dissolves in excess to give an azure blue solution.
The flame test is a quick method to identify certain metal ions. The procedure is specific: clean a platinum wire by dipping in concentrated hydrochloric acid and heating until it imparts no colour to the flame. Then make a paste of the substance with concentrated hydrochloric acid on the wire, and introduce it to the non-luminous flame.
Sodium produces a persistent golden-yellow flame that vanishes when viewed through blue glass. Potassium gives a violet or lilac flame, visible as violet or pink through blue glass. Calcium produces a brick-red flame, appearing light green through blue glass. Copper gives a peacock bluish-green flame, remaining bluish-green through blue glass.
Finally, a brief note on the pH scale and indicators. The pH scale measures hydrogen ion concentration, ranging from zero to fourteen. A pH of seven is neutral, like pure water. Values below seven indicate acidity, with lower numbers meaning stronger acids. Values above seven indicate alkalinity, with higher numbers meaning stronger bases.
Universal indicator provides a broad pH range, changing from red through orange, yellow, green, blue, indigo to violet as pH increases from one to fourteen. This allows you to estimate the pH of unknown solutions by comparing the colour produced with a standard chart.
Let us recap the key takeaways from this chapter.
First, gases are identified by their preparation methods, physical properties, and specific chemical tests — hydrogen by its pop sound, oxygen by rekindling a glowing splint, carbon dioxide by turning lime water milky, chlorine by bleaching litmus and turning starch-iodide paper blue-black, and so on.
Second, the action of heat on substances reveals their identity through colour changes, gases evolved, and residues formed — copper compounds yield black copper oxide, zinc compounds give yellow-hot white-cold zinc oxide, and nitrates typically produce nitrogen dioxide and oxygen.
Third, cations are identified by their characteristic precipitates with sodium hydroxide and ammonium hydroxide, noting which precipitates dissolve in excess reagent.
Fourth, anions are identified by their reactions with dilute and concentrated acids, and by specific tests such as the brown ring test for nitrates and the barium chloride test for sulphates.
Fifth, colourless acids and alkalis are distinguished using indicators and by their reactions with carbonates.
Sixth, the flame test provides rapid identification of sodium, potassium, calcium, and copper ions through their characteristic flame colours.
Practical chemistry is where theory meets reality. These systematic methods of analysis — observing carefully, testing methodically, and reasoning logically — are the foundation of scientific investigation. Master these techniques, and you will approach the laboratory with confidence and precision. Keep practising, stay curious, and remember: every test tube holds a discovery waiting to be made. Until next time, happy experimenting!