NCERT Class 10 Science Chapter 4 Activities — Carbon and its Compounds
NCERT Class 10 Science — Carbon and its Compounds, Activities 4.1 to 4.12 on pages 58 to 76.
Also here. Solutions to the end-of-chapter questions for this chapter and the in-text questions are on their own pages.
This chapter’s activities are spread thinly at the start and then come in a rush. Activity 4.1 opens the chapter and Activity 4.2 is a paper exercise nine pages later; everything else is packed into the last ten pages, where the chapter turns from structure to reactions — combustion, oxidation, the reactions of ethanol and ethanoic acid, and finally soap.
That shape is worth knowing when planning. The first half of the chapter is taught from diagrams and structures, and there is nothing to demonstrate until combustion.
Three ideas carry nearly all twelve:
- Carbon compounds burn, and how they burn — sooty or clean — depends on the oxygen supply, not on the fuel alone.
- Ethanol and ethanoic acid have distinct functional groups, and each behaves in a way the group predicts: ethanol gives hydrogen with sodium, ethanoic acid behaves as a weak acid.
- Soap works because its molecule has two incompatible ends, and everything about washing — including why it fails in hard water — follows from that.
Key insight. Activities 4.10, 4.11 and 4.12 are one experiment in three parts, and they are much stronger done together than spread over three lessons. The first shows soap working, the second shows it failing in hard water, and the third shows detergent working where soap failed. Three tubes, one after another, make the argument that no amount of explanation does.
Activity 4.1
List ten things you have used or consumed since the morning. Pool the list with your classmates’ and sort the items into the table in the book, putting anything made of more than one material into both columns.
What you should observe. Almost everything ends up in the carbon column.
Nearly every item on the list — food, clothing, fuel, plastics, paper, soap, medicines — is made of carbon compounds. Very few things are not.
Why the chapter opens with it. It is not really an experiment; it is a way of earning the claim the chapter is about to make, that carbon is a versatile element forming “the basis for all living organisms and many of the things we use”. A student who has just been told that nods; a student who has sorted their own list of ten things has seen it.
What goes wrong. Listing only obviously chemical things. The value is in noticing that a cotton shirt, a chapati and a plastic bottle are all carbon compounds, so push for ordinary objects rather than laboratory ones.
Activity 4.2
Work out the difference in formula and in molecular mass between (a) CH₃OH and C₂H₅OH, (b) C₂H₅OH and C₃H₇OH, and (c) C₃H₇OH and C₄H₉OH. Is the difference the same each time? Arrange them by increasing number of carbon atoms and ask whether this is a homologous series. Then generate the series for other functional groups.
What you should observe. The same gap every time, in both formula and mass.
Each successive pair differs by CH₂, a mass difference of 14 u every time. That constant difference is what makes them a homologous series.
| Pair | Difference | Mass difference |
|---|---|---|
| CH₃OH → C₂H₅OH | CH₂ | 14 u |
| C₂H₅OH → C₃H₇OH | CH₂ | 14 u |
| C₃H₇OH → C₄H₉OH | CH₂ | 14 u |
Why. A homologous series is a family sharing one functional group, where consecutive members differ by a fixed CH₂ unit. Because the functional group is the same throughout, chemical properties stay similar down the series while physical properties — melting point, boiling point — change gradually with chain length.
What goes wrong. Computing the mass difference as 12 (carbon only) and forgetting the two hydrogens. It is CH₂ = 12 + 2 = 14, and the number is worth having by heart because examination questions hinge on it.
Activity 4.3
Take naphthalene, camphor and alcohol in turn on a spatula and burn them. Note the nature of the flame and whether smoke is produced. Hold a metal plate above each flame and look for a deposit. The book asks for the teacher’s assistance.
What you should observe. They do not burn alike at all.
Naphthalene and camphor burn with a yellow, sooty flame and leave a black deposit on the plate. Alcohol burns with a clean blue flame and leaves little or none.
$$\mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2} + \text{heat and light}$$ $$\mathrm{CH_4} + 2\mathrm{O_2} \rightarrow \mathrm{CO_2} + 2\mathrm{H_2O} + \text{heat and light}$$ $$\mathrm{CH_3CH_2OH} + 3\mathrm{O_2} \rightarrow 2\mathrm{CO_2} + 3\mathrm{H_2O} + \text{heat and light}$$
Why the difference. Naphthalene and camphor are aromatic and have a high proportion of carbon relative to hydrogen. There is not enough oxygen reaching the flame to burn all that carbon to carbon dioxide, so some is released as unburnt carbon — soot. Ethanol already carries an oxygen atom of its own and has far less carbon per molecule, so it burns completely.
What goes wrong. Holding the plate in the flame rather than above it. In the flame it simply gets hot and the deposit burns off again; held a little above the tip, the soot collects and is unmistakable.
Activity 4.4
Light a bunsen burner and adjust the air hole at its base to get different flames. When do you get a yellow sooty flame? When do you get a blue one?
What you should observe. The same gas gives two completely different flames depending only on the air hole.
Air hole closed → yellow, sooty, luminous flame. Air hole open → clean blue flame.
Why this is the companion to 4.3. Activity 4.3 varied the fuel and kept the air the same; this varies the air and keeps the fuel the same. Together they show that sootiness is about the oxygen supply, not just the substance. With the hole open, the gas is premixed with air and burns completely.
The practical consequence the chapter draws. “If you observe the bottoms of cooking vessels getting blackened, it means that the air holes are blocked and fuel is getting wasted.” That is a genuinely useful thing to be able to tell a class — the soot is unburnt fuel, so a sooty flame is an expensive one.
Why some things glow instead of flaming. A flame needs a gaseous substance to burn. Wood and charcoal flame at first, while their volatile compounds vaporise; once those are gone, the charcoal glows red without a flame.
What goes wrong. Opening the air hole fully at once, which can make the flame strike back and burn at the jet. Open it gradually.
Activity 4.5
Take about 3 mL of ethanol in a test tube and warm it gently in a water bath. Add 5% alkaline potassium permanganate drop by drop. Does the purple colour persist at first? Why does it stop disappearing once excess is added?
What you should observe. The purple vanishes as each early drop goes in, and then suddenly stops vanishing.
The purple colour disappears at first, because the permanganate is being used up oxidising ethanol to ethanoic acid. Once all the ethanol has been oxidised, there is nothing left to react with, so further permanganate stays purple.
$$\mathrm{CH_3CH_2OH} \xrightarrow[\text{alkaline } \mathrm{KMnO_4}]{\ \text{oxidation}\ } \mathrm{CH_3COOH}$$
Why. Alkaline potassium permanganate adds oxygen to the ethanol, converting an alcohol into a carboxylic acid. The chapter defines oxidising agents on exactly this basis: substances “capable of adding oxygen to others”, with alkaline permanganate and acidified potassium dichromate as its two examples.
Why the persisting colour is the real result. The point at which the purple stops disappearing marks the moment the ethanol runs out — so the colour change is doing the work of an end point, exactly as phenolphthalein did in the neutralisation of Chapter 2.
What goes wrong. Adding the permanganate too quickly, which floods the tube and leaves it purple from the start, with nothing seen. Drop by drop, with swirling, is the whole method.
Activity 4.6
Drop a piece of sodium the size of a couple of grains of rice into absolute alcohol. What do you observe, and how would you test the gas? The book marks this as a teacher’s demonstration.
What you should observe. Brisk fizzing, and the sodium dissolves away.
Bubbles of hydrogen are given off, testable by the pop with a burning splint. The other product is sodium ethoxide.
$$2\mathrm{Na} + 2\mathrm{CH_3CH_2OH} \rightarrow 2\mathrm{CH_3CH_2O^-Na^+} + \mathrm{H_2}$$
Why it is worth the risk of using sodium. The chapter asks the follow-up itself: “Can you recall which other substances produce hydrogen on reacting with metals?” Acids do — Chapter 2, Activity 2.3. Getting hydrogen from an alcohol as well shows that the —OH group has acidic character of a sort, which is the property the ethoxide is evidence for.
What goes wrong. Using ordinary alcohol rather than absolute. Rectified spirit contains water, and the sodium reacts violently with the water instead — a different, far more vigorous reaction that can ignite. This is a demonstration with a very small piece of sodium, behind a screen.
Activity 4.7
Compare the pH of dilute acetic acid and dilute hydrochloric acid using both litmus paper and universal indicator. Do both show as acids on litmus? Does the universal indicator show them as equally strong?
What you should observe. Litmus cannot tell them apart. The universal indicator can.
Litmus: both turn blue litmus red — both are acids. Universal indicator: hydrochloric acid gives a much lower pH than acetic acid of the same concentration. Acetic acid is a weak acid.
Why. In the chapter’s words, “unlike mineral acids like HCl, which are completely ionised, carboxylic acids are weak acids”. Both release $\mathrm{H^+}$ in water, but hydrochloric acid releases essentially all of it while acetic acid releases only a small fraction — so at the same concentration there are far fewer hydrogen ions, and the pH is higher.
Why two indicators. This activity exists to show the limits of an indicator. Litmus answers a yes/no question and cannot do more; the universal indicator answers a question of degree. Choosing the right tool for the question is the lesson.
What goes wrong. Comparing acids of different concentrations, which makes the result meaningless. Both must be at the same dilution for the comparison to say anything about strength.
Activity 4.8
Take 1 mL of absolute ethanol and 1 mL of glacial acetic acid with a few drops of concentrated sulphuric acid in a test tube. Warm in a water bath for at least five minutes. Pour into a beaker of 20–50 mL of water and smell it.
What you should observe. A sweet, fruity smell, quite unlike either starting material.
A sweet, fruity smell — an ester, ethyl ethanoate.
$$\mathrm{CH_3COOH} + \mathrm{CH_3CH_2OH} \xrightarrow[\text{catalyst}]{\ \text{acid}\ } \mathrm{CH_3COOCH_2CH_3} + \mathrm{H_2O}$$
Why each ingredient is there. The acetic acid must be glacial and the ethanol absolute because the reaction produces water, and starting with water present pushes the equilibrium the wrong way. The concentrated sulphuric acid is a catalyst, not a reactant.
Why pour it into water before smelling. Two reasons, and both matter. It dilutes the sharp smells of the unreacted acid and alcohol, which would otherwise mask the ester. And it is far safer than sniffing a hot tube containing concentrated sulphuric acid. Waft the smell towards you; never inhale over the beaker.
What it sets up. Esters are what give fruits their smells and are used in perfumes. Treated with a base they go back to alcohol and the sodium salt of the acid — saponification, the reaction that makes soap:
$$\mathrm{CH_3COOC_2H_5} + \mathrm{NaOH} \rightarrow \mathrm{C_2H_5OH} + \mathrm{CH_3COONa}$$
What goes wrong. Not warming for long enough. Five minutes in the water bath is a minimum, and a tube warmed for one minute gives a smell of vinegar and nothing else.
Activity 4.9
Set up the apparatus from Chapter 2’s Activity 2.5. Put a spatula full of sodium carbonate in a test tube and add 2 mL of dilute ethanoic acid. Pass the gas through freshly prepared lime water. Repeat with sodium hydrogencarbonate.
What you should observe. Brisk fizzing, and lime water that turns milky — exactly as with a mineral acid.
Effervescence of carbon dioxide, turning lime water milky. Yes, the gas can be identified this way.
$$2\mathrm{CH_3COOH} + \mathrm{Na_2CO_3} \rightarrow 2\mathrm{CH_3COONa} + \mathrm{H_2O} + \mathrm{CO_2}$$ $$\mathrm{CH_3COOH} + \mathrm{NaHCO_3} \rightarrow \mathrm{CH_3COONa} + \mathrm{H_2O} + \mathrm{CO_2}$$
Why it matters that this is the same as Chapter 2. Ethanoic acid is a weak acid, but it is genuinely an acid — it does everything a mineral acid does, only less vigorously. It reacts with a base to give a salt and water, and with carbonates to give carbon dioxide. The salt here, sodium ethanoate, is the same sodium acetate met as a salt in Chapter 2’s Activity 2.14.
What goes wrong. Using lime water that has stood open. The chapter says “freshly prepared” for a reason: lime water left in an open bottle has already absorbed carbon dioxide from the air and is spent, so the test fails on perfectly good gas.
Activity 4.10
Put about 10 mL of water in each of two test tubes. Add a drop of cooking oil to both, label them A and B, and add a few drops of soap solution to B only. Shake both vigorously for the same length of time and let them stand.
What you should observe. Both look cloudy while shaking. On standing, the difference appears.
In tube A the oil and water separate into layers again. In tube B with soap they stay mixed as a stable, cloudy emulsion.
Why. A soap molecule is the sodium or potassium salt of a long-chain carboxylic acid, and its two ends want opposite things — the ionic head is hydrophilic and interacts with water, the hydrocarbon tail is hydrophobic and interacts with oil. Surrounded by water, the molecules cluster into micelles with their tails tucked into the oil droplet and their ionic heads facing out. The droplet is then wrapped in a charged shell that keeps it suspended, so it cannot coalesce and separate.
Why this is the whole explanation of washing. Most dirt is oily, and oil does not dissolve in water. Soap does not dissolve the oil either — it surrounds it and carries it away in suspension, which is why clothes need rinsing rather than just soaking.
What goes wrong. Shaking the two tubes for different lengths of time, which is the one thing the instructions guard against. And letting them stand too briefly: tube A needs a minute or two for the layers to reappear, and the comparison only exists once they have.
Activity 4.11
Take about 10 mL of distilled water (or rain water) and 10 mL of hard water from a tubewell or hand-pump in separate test tubes. Add a couple of drops of soap solution to each, shake for the same length of time, and compare the foam.
What you should observe. One lathers; the other produces a nasty white solid instead.
More foam in the distilled water. In the hard water the soap gives a white curdy precipitate — scum — and little lather.
Why. Hard water contains calcium and magnesium ions. Soap is a sodium or potassium salt of a long-chain acid, and the calcium and magnesium salts of that same acid are insoluble — so the soap is thrown out of solution as scum before it can do any cleaning. Soap is wasted twice over: it fails to lather, and it leaves a deposit on the cloth.
If you have no hard water to hand. The book’s own note to the teacher: prepare some by dissolving hydrogencarbonates, sulphates or chlorides of calcium or magnesium in water.
What goes wrong. Using tap water as the “soft” sample. In much of India tap water is already hard, so both tubes scum and the comparison shows nothing. Distilled or rain water is specified for exactly this reason.
Activity 4.12
Take about 10 mL of hard water in each of two test tubes. Add five drops of soap solution to one and five drops of detergent solution to the other. Shake both for the same period.
What you should observe. The detergent lathers in water where the soap could not.
The detergent foams well; the soap does not. The curdy solid forms only in the soap tube.
Why. Detergents are ammonium or sulphonate salts of long-chain hydrocarbons rather than carboxylate salts. Their calcium and magnesium salts are soluble, so there is nothing to precipitate and the detergent keeps working in hard water. This is why detergents are used in shampoos and washing powders.
What it completes. Read 4.10, 4.11 and 4.12 as one argument: soap works by forming micelles; hard water destroys soap by precipitating it; detergents keep the micelle-forming structure while avoiding the precipitation. Each activity answers the problem the one before it raised.
The chapter’s own follow-up question, worth asking aloud. “Would you be able to check if water is hard by using a detergent?” No — and that is precisely the point. The detergent foams either way, so it cannot distinguish hard water from soft. Only the soap can, which makes the failing reagent the useful test.
What goes wrong. Using different amounts, or shaking unevenly. Five drops in each, shaken for the same time, is the only way the foam heights mean anything.
Common mistakes
- Taking the mass gap in Activity 4.2 as 12. The homologous difference is CH₂, so 14 u — carbon plus two hydrogens.
- Blaming the fuel alone for the soot in Activity 4.3. Activity 4.4 gets a sooty flame and a clean one out of the same gas by changing only the air hole. Incomplete combustion is about the oxygen supply.
- Reporting that the permanganate “did not work” in Activity 4.5 because the colour stayed. The colour staying is the end of the reaction, not a failure of it — it marks the point where the ethanol ran out.
- Using rectified spirit instead of absolute alcohol in Activity 4.6. The sodium then reacts with the water in it, which is both a different reaction and a much more dangerous one.
- Concluding from Activity 4.7 that acetic acid is not really an acid. It is — litmus proves it. It is a weak acid, meaning only partly ionised, and that is a different claim from not being acidic.
- Smelling Activity 4.8 straight from the hot tube. Pour into water first. It dilutes the sharp unreacted acid, and it avoids inhaling over concentrated sulphuric acid.
- Saying soap dissolves the grease in Activity 4.10. It does not. It surrounds the oil in micelles and holds it in suspension — which is why rinsing is needed to carry it away.
- Testing for hardness with detergent. Detergent lathers in hard and soft water alike, so it cannot tell them apart. Only soap can, because only soap fails.
Practise next
- The end-of-chapter Exercises for this chapter, where the combustion of Activities 4.3 and 4.4 returns as questions on complete and incomplete burning, and the soap of 4.10 to 4.12 as questions on micelles and hard water.
- The in-text questions of this chapter, including the two the activities above quote directly — whether a detergent can test for hardness, and how ethanoic acid reacts with carbonates.
- The activities of Chapter 2, where Activity 2.5’s lime water test is the same one used here in Activity 4.9, and Activity 2.8’s conductivity apparatus makes the same argument about ions that Activity 4.7 makes about strength.
- Chapter 3, Metals and Non-metals, where the hydrogen given off by sodium in Activity 4.6 is the same reaction seen with metals and acids in Activity 3.11.
Doubts are answered by Shiwam, usually within a day. Ask about this question specifically — a general question about the chapter is better asked in class.