Hello, and welcome to today's chemistry lesson. Today, we are diving into practical chemistry. This chapter is all about bringing chemistry to life in the laboratory. We will learn how to recognise and identify common gases, observe what happens when substances are heated, understand the flame test for metal identification, explore the chemistry of hard and soft water, and examine the important topic of water pollution. Let us begin.
First, let us explore how to recognise and identify gases in the laboratory. Knowing how to detect a gas from its appearance, smell, and chemical behaviour is a fundamental skill for any chemist.
We begin with hydrogen, H₂. To prepare hydrogen, add dilute hydrochloric acid or dilute sulphuric acid to a reactive metal such as magnesium, zinc, or iron. The metal reacts vigorously, producing a colourless, odourless gas that is neutral to litmus. When pure hydrogen burns, it produces a pale blue flame. However, when hydrogen is mixed with air, it burns with a characteristic pop sound. This pop test is a classic confirmation for hydrogen.
Next, oxygen, O₂. Oxygen is prepared by heating heavy metal oxides such as lead oxide or mercury oxide. The gas is colourless, odourless, and neutral to litmus. Its most distinctive property is that it rekindles a glowing wooden splinter, causing it to burst into flame. Oxygen is also absorbed by an alkaline solution of pyrogallol, turning it dark brown.
Carbon dioxide, CO₂, is prepared by heating metallic carbonates except sodium and potassium carbonates, or by adding dilute acid to any carbonate or hydrogen carbonate. The gas is colourless and odourless, but it turns moist blue litmus faintly red, showing its acidic nature. When passed through lime water, carbon dioxide turns it milky due to the formation of insoluble calcium carbonate. Interestingly, if you continue passing carbon dioxide, the milkiness disappears as the precipitate dissolves to form soluble calcium hydrogen carbonate. Carbon dioxide has no effect on acidified potassium dichromate or potassium permanganate solutions.
Chlorine, Cl₂, is prepared by adding concentrated hydrochloric acid to oxidising agents like manganese dioxide, lead dioxide, or red lead, followed by gentle heating. Chlorine is immediately recognisable by its greenish-yellow colour and sharp, pungent, choking odour. It turns moist blue litmus red, then bleaches it completely. This bleaching action occurs because chlorine reacts with water to form hypochlorous acid, which releases oxygen that destroys colour. Chlorine also turns moist starch-iodide paper blue-black, due to the liberation of iodine. When passed through silver nitrate solution, chlorine produces a white precipitate of silver chloride.
Hydrogen chloride, HCl, is prepared by heating a metal chloride like sodium chloride or potassium chloride with concentrated sulphuric acid. It is a colourless gas with a pungent, choking odour that turns moist blue litmus red. When a glass rod dipped in ammonia solution is brought near hydrogen chloride 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, SO₂, is prepared by adding dilute acid to a metallic sulphite. It is a colourless gas with the suffocating smell of burning sulphur. It extinguishes a burning splinter and turns moist blue litmus red before bleaching it temporarily. Sulphur dioxide turns lime water milky, similar to carbon dioxide, but the milkiness also disappears with excess gas. The key distinguishing tests are with acidified potassium permanganate and acidified potassium dichromate. Sulphur dioxide decolourises pink potassium permanganate and turns orange potassium dichromate green. Carbon dioxide does neither of these. Additionally, sulphur dioxide has no effect on lead acetate paper, unlike hydrogen sulphide.
Hydrogen sulphide, H₂S, is prepared by adding dilute acid to metallic sulphides like iron sulphide or zinc sulphide. This gas has a foul, unmistakable odour of rotten eggs. It turns moist blue litmus red and, most characteristically, turns lead acetate paper black due to the formation of lead sulphide. It also blackens lead nitrate solution.
Ammonia, NH₃, is prepared by heating an ammonium salt with an alkali. It is a colourless gas with a sharp, pungent smell. Unlike the acidic gases we have discussed, ammonia turns moist red litmus blue, showing its basic nature. When a rod dipped in concentrated hydrochloric acid is brought near ammonia, dense white fumes of ammonium chloride appear. Ammonia turns Nessler's reagent brown and, when passed through copper sulphate solution, first forms a pale blue precipitate that dissolves in excess ammonia to give a deep blue solution.
Water vapour, H₂O, is obtained by heating hydrated salts like washing soda or blue vitriol. The colourless vapours condense on cooler parts of the test tube as droplets of liquid. This liquid is neutral to litmus but turns white anhydrous copper sulphate blue and blue cobalt chloride paper pink.
Finally, nitrogen dioxide, NO₂, is prepared by heating heavy metal nitrates like copper nitrate or lead nitrate. Note that sodium nitrate and potassium nitrate do not produce nitrogen dioxide when heated. It is a brown gas with an irritating, pungent odour. It turns moist blue litmus red and liberates iodine from potassium iodide, producing violet vapours. It also turns green acidified ferrous sulphate solution brown.
Now, let us turn to the action of heat on substances. When you heat an unknown compound, careful observation can reveal its identity.
Copper carbonate, a light green amorphous powder, turns black on heating and gives off carbon dioxide. This gas turns lime water milky and has no effect on acidified potassium dichromate or potassium permanganate. The black residue is copper oxide.
Zinc carbonate, a white amorphous solid, becomes pale yellow when hot and white again when cool. The milkiness it produces in lime water disappears on passing excess gas, confirming carbon dioxide. It too evolves carbon dioxide.
Washing soda, or hydrated sodium carbonate, swells, crumbles and then melts when heated, giving off steamy vapours. These vapours condense as droplets on cooler parts of the test tube, turn anhydrous copper sulphate blue, and turn blue cobalt chloride paper pink. The white residue is anhydrous sodium carbonate.
Blue copper sulphate crystals crumble to a white amorphous powder when heated, releasing water vapour. This vapour turns anhydrous copper sulphate blue and turns blue cobalt chloride paper pink. On stronger heating, the white powder turns black as copper oxide forms, and sulphur dioxide and oxygen are evolved. The sulphur dioxide turns moist blue litmus red and changes acidified potassium dichromate from orange to green.
Zinc nitrate, a white crystalline deliquescent solid, melts to a sticky mass and gives off steamy vapours. These vapours turn anhydrous copper sulphate blue and cobalt chloride paper pink. On strong heating, reddish-brown nitrogen dioxide fumes appear, a glowing splint bursts into flame in the gas showing oxygen is present, and the residue is yellow when hot, white when cold. This is zinc oxide.
Copper nitrate, a bluish-green crystalline solid, behaves similarly. First, it loses water to form a bluish-green mass. Then on further heating, it forms black copper oxide, with nitrogen dioxide and oxygen as gases.
Lead nitrate, a heavy white crystalline solid, decomposes with a crackling sound, producing reddish-brown nitrogen dioxide. The residue is reddish-brown when hot and yellow when cold. It partly fuses with and stains the glass yellow. A glowing splint relights in the gas, showing oxygen is present.
Ammonium chloride sublimes on heating, forming dense white fumes of ammonia and hydrogen chloride. These form a white mass on cooler parts, with no residue left behind.
Iodine, a violet crystalline solid, sublimes to form violet vapours. These condense as violet crystals on cooler parts, with no residue left behind. Starch paper turns blue-black in contact with iodine vapour, and paper dipped in silver nitrate solution turns yellow.
Ammonium dichromate, an orange-red crystalline solid, decomposes violently with flashes of light, swelling up and leaving a greenish-grey residue of chromium oxide. Water vapour and nitrogen gas are evolved.
The flame test provides another powerful tool for identifying metal ions. Clean a platinum wire by dipping it in concentrated hydrochloric acid and heating it in the non-luminous flame until the flame shows no colour. Repeat this process until the wire is ready. Then dip the wire in concentrated hydrochloric acid, then into the substance so a little sticks to it. Hold it in the non-luminous part of the flame.
Sodium salts impart a persistent golden-yellow flame that vanishes when viewed through blue glass. Potassium salts give a lilac or violet flame, seen as violet or pink through blue glass. Calcium compounds produce a brick-red flame that is fugitive, appearing and disappearing, and appears light green through blue glass. Copper salts give a peacock bluish-green flame, and the wire slowly gets corroded.
Let us now examine the action of dilute sulphuric acid on unknown substances.
If a reactive metal like zinc, iron, or magnesium reacts with dilute sulphuric acid, hydrogen gas evolves with effervescence. It burns with a pop sound when a burning splint is brought near it.
If a carbonate reacts, carbon dioxide evolves with brisk effervescence, turning lime water milky. It has no effect on acidified potassium dichromate.
If a sulphide reacts, hydrogen sulphide gas evolves with its characteristic rotten egg smell. It turns lead acetate paper black and also blackens lead nitrate solution.
If a sulphite reacts, sulphur dioxide evolves with a suffocating smell of burning sulphur. It extinguishes a burning splinter, turns lime water milky, and turns acidified potassium dichromate paper green.
Now we turn to hard water and soft water. Hard water contains dissolved calcium and magnesium salts and does not form lather with ordinary soap. Soft water, free from these salts, lathers readily.
Temporary hardness is caused by calcium and magnesium bicarbonates. It is removable by boiling followed by filtration.
Permanent hardness is caused by calcium and magnesium chlorides and sulphates. It is removable by washing soda or caustic soda.
Soap is the sodium or potassium salt of an organic fatty acid. In hard water, soap reacts with calcium and magnesium ions to form insoluble scum, wasting the soap.
Detergents are sodium salts of alkyl sulphonic acids. They contain a sulphonic acid group instead of a carboxylic group. Due to the solubility of their calcium and magnesium salts in water, they do not form scum and can lather even with hard water.
Finally, let us discuss water pollution. Water pollution occurs when substances are introduced that alter the natural quality of water and impair its usefulness for living organisms.
Polluted water may have a foul smell, bad taste, oil and grease on the surface, or excessive algae growth. The main sources include untreated sewage, industrial effluents containing heavy metals, agricultural runoff, and oil spills. Thermal pollution from industrial cooling also raises water temperature and depletes oxygen, harming aquatic life.
Water quality can be assessed through several parameters. The pH value indicates acidity or alkalinity. Non-polluted water has a pH of seven or close to seven. Values much higher or lower than seven indicate polluted and potentially harmful water.
Dissolved oxygen is essential for aquatic life. Flowing water contains more dissolved oxygen than stagnant water. Low dissolved oxygen indicates pollution.
Biological oxygen demand, or BOD, measures the oxygen utilized by microorganisms during oxidation of organic substances. BOD increases in polluted water, with higher values indicating greater pollution.
Turbidity is the amount of particulate matter suspended in water. It measures the scattering effect that suspended solids have on light. High turbidity reduces water's ability to support life.
Bacterial growth in water samples also indicates pollution. Clean water should be free of harmful bacteria.
To control water pollution, sewage must be fully treated before release. Industrial chemicals should be neutralised. Solids can be removed by settlement and screening. Organic matter can be removed by oxidation or precipitation. Pathogens can be destroyed by ultraviolet radiation.
Let us recap the key takeaways from this chapter.
First, you can identify gases by their colour, odour, effect on litmus, and specific chemical tests. These include the pop sound for hydrogen, rekindling of a glowing splint for oxygen, lime water turning milky for carbon dioxide, and acidified potassium dichromate or permanganate for distinguishing sulphur dioxide.
Second, heating substances produces characteristic changes in colour, residue formation, sublimation, and gas evolution that help identify unknown compounds.
Third, the flame test uses characteristic colours to identify sodium, potassium, calcium, and copper ions.
Fourth, dilute sulphuric acid reacts differently with metals, carbonates, sulphides, and sulphites, producing distinct gases that can be identified.
Fifth, hard water contains calcium and magnesium salts that prevent soap from forming lather. Soft water lathers readily. Temporary hardness is removed by boiling followed by filtration. Permanent hardness is removed by washing soda or caustic soda.
Sixth, water pollution can be monitored through pH, dissolved oxygen, BOD, turbidity, and bacterial content. It can be controlled through sewage treatment, neutralization of industrial chemicals, settlement and screening, oxidation or precipitation of organic matter, and ultraviolet radiation for pathogens.
That brings us to the end of our lesson on practical chemistry. Remember, the laboratory is where theory meets reality. Every observation you make, every gas you identify, and every colour change you notice builds your skills as a chemist. Stay curious, observe carefully, and always think about what your results mean. Until next time, keep exploring the fascinating world of chemistry.