NCERT Class 10 Science Chapter 2 Activities — Acids, Bases and Salts
NCERT Class 10 Science — Acids, Bases and Salts, Activities 2.1 to 2.15 on pages 18 to 32.
Also here. Solutions to the end-of-chapter questions for this chapter and the in-text questions are on their own pages.
Fifteen activities, more than any other chapter in the book, and they are not a miscellany. The chapter is building one argument across them: that everything acids have in common, and everything bases have in common, comes down to the ions they release in water.
The sequence goes like this. Activities 2.1 and 2.2 establish how you detect an acid or a base at all. Activities 2.3 to 2.7 work through what acids and bases react with — metals, carbonates, each other, metal oxides. Activity 2.8 is the pivot: it shows that acid solutions conduct electricity, which means ions. Activities 2.9 and 2.10 pin that down — no water, no ions, no acidity. Activities 2.11 to 2.15 then take the idea into pH and into salts.
Key insight. Activity 2.8 is the one to protect time for. Everything before it is a collection of reactions that happen to look similar; everything after it is explained by a single cause. A class that does 2.8 properly — and sees the bulb stay dark for glucose — has the whole chapter. A class that skips it is left memorising a list.
Activity 2.1
Collect dilute solutions of hydrochloric, sulphuric, nitric and acetic acids, and of sodium, calcium, potassium, magnesium and ammonium hydroxides. Put a drop of each on a watch-glass and test with red litmus, blue litmus, phenolphthalein and methyl orange, filling in Table 2.1.
What you should observe. Each indicator answers only one of the two questions, which is the point of using four.
| Indicator | In acid | In base |
|---|---|---|
| Red litmus | no change | turns blue |
| Blue litmus | turns red | no change |
| Phenolphthalein | colourless | turns pink |
| Methyl orange | turns red | turns yellow |
Why. Litmus of one colour can only detect its opposite — red litmus stays red in acid, so it tells you nothing there. This is exactly what the chapter’s own in-text question asks on page 18: given three tubes and only red litmus paper, how do you identify distilled water, an acid and a base? The answer depends on noticing that the base is the only thing red litmus can report, and that the acid and the water then have to be told apart by first making a blue strip with the base.
What goes wrong. Using a dropper between solutions without rinsing. One drop of NaOH carried into the HCl watch-glass neutralises it, and the class gets a confident, wrong result with no way of knowing.
Activity 2.2
Chop onions finely, seal them in a plastic bag with clean cloth strips, and leave overnight in the fridge. Check the odour of two strips, then put dilute HCl on one and dilute NaOH on the other. Repeat with dilute vanilla essence and clove oil.
What you should observe. The onion smell survives contact with the acid and is destroyed by the base. Vanilla behaves the same way.
The strip treated with base loses its onion smell; the one treated with acid keeps it. Vanilla and onion work as olfactory indicators; clove oil does not.
Why. An olfactory indicator is a substance whose smell, rather than its colour, changes with acidity. The chapter introduces them precisely because they let a student who cannot see a colour change still do the experiment — and because they make the point that “indicator” is a broader idea than litmus.
What goes wrong. The overnight step is not optional and is the one most often skipped. It takes real time for the cloth to pick up enough of the onion’s compounds to smell strongly, and a bag made up at the start of the lesson smells of nothing by the end of it.
Activity 2.3
Set up the apparatus of Fig. 2.1. Put about 5 mL of dilute sulphuric acid in a test tube with a few zinc granules. Watch the surface of the zinc, pass the gas through soap solution, and bring a burning candle to a gas-filled bubble. Repeat with HCl, HNO₃ and CH₃COOH. The book marks this one as needing the teacher’s assistance.
What you should observe. Bubbles form on the zinc, the soap solution froths with gas-filled bubbles, and each bubble goes off with a pop at the candle.
Hydrogen gas, which burns with a pop. The reaction is the same for every acid tried:
$$\mathrm{Zn}(s) + \mathrm{H_2SO_4}(aq) \rightarrow \mathrm{ZnSO_4}(aq) + \mathrm{H_2}(g)$$
$$\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas}$$
Why. The chapter states the general rule on page 20: the metal displaces hydrogen from the acid as hydrogen gas and forms a salt. The soap solution exists only to trap the gas — hydrogen is far lighter than air and would otherwise be gone before it could be tested.
What goes wrong. Acetic acid is the interesting failure. It fizzes so much more slowly than the others that a class short of time concludes nothing is happening. That is not a failed experiment: it is the chapter’s own Exercise question about magnesium in HCl against acetic acid, and it is worth letting the difference be seen rather than explained.
Activity 2.4
Put a few pieces of granulated zinc in a test tube, add 2 mL of sodium hydroxide solution, warm it, and repeat the rest of the steps of Activity 2.3.
What you should observe. Hydrogen again — the same pop — this time from a base.
Hydrogen gas once more, with sodium zincate left in solution.
$$\mathrm{Zn}(s) + 2\mathrm{NaOH}(aq) \rightarrow \mathrm{Na_2ZnO_2}(aq) + \mathrm{H_2}(g)$$
Why it matters more than it looks. Hydrogen coming from a base breaks the pattern a student has just learned from 2.3, and the chapter is careful to add that “such reactions are not possible with all metals”. Zinc is amphoteric — it reacts with both acids and alkalis. This is the activity that stops “metal + acid gives hydrogen” hardening into “only acids give hydrogen”.
What goes wrong. It needs warming, and cold NaOH gives almost nothing. It is also much slower than 2.3; set it going and come back to it.
Activity 2.5
Take about 0.5 g of sodium carbonate in tube A and 0.5 g of sodium hydrogencarbonate in tube B. Add about 2 mL of dilute HCl to each. Pass the gas produced through lime water.
What you should observe. Both fizz briskly, and in both cases the lime water turns milky.
Brisk effervescence of carbon dioxide from both, turning lime water milky.
$$\mathrm{Na_2CO_3}(s) + 2\mathrm{HCl}(aq) \rightarrow 2\mathrm{NaCl}(aq) + \mathrm{H_2O}(l) + \mathrm{CO_2}(g)$$ $$\mathrm{NaHCO_3}(s) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l) + \mathrm{CO_2}(g)$$ $$\mathrm{Ca(OH)_2}(aq) + \mathrm{CO_2}(g) \rightarrow \mathrm{CaCO_3}(s) + \mathrm{H_2O}(l)$$
Why the lime water goes milky. The milkiness is a precipitate of calcium carbonate, insoluble in water. Keep passing gas through and it clears again, because excess carbon dioxide converts the carbonate to the soluble hydrogencarbonate:
$$\mathrm{CaCO_3}(s) + \mathrm{H_2O}(l) + \mathrm{CO_2}(g) \rightarrow \mathrm{Ca(HCO_3)_2}(aq)$$
What goes wrong. Passing gas through for too long, and then reporting that lime water does not detect carbon dioxide. The clearing is a second reaction, not a failure of the first — and it is a good thing to show deliberately rather than have a student discover it by accident and mistrust the test.
Activity 2.6
Take about 2 mL of dilute NaOH in a test tube with two drops of phenolphthalein. Note the colour. Add dilute HCl drop by drop until the colour changes. Then add a few drops of NaOH again.
What you should observe. Pink, then colourless, then pink again — and it can be cycled as often as you like.
Pink → colourless on adding acid → pink again on adding base.
$$\mathrm{NaOH}(aq) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)$$
$$\text{Base} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$$
Why. In the chapter’s words, “the effect of a base is nullified by an acid and vice-versa”. Phenolphthalein is pink only while the solution is basic, so the colour vanishes at the point the acid has used up the base. This is the chapter’s definition of a neutralisation reaction.
What goes wrong. Adding the acid too fast. The whole value of the activity is in creeping up on the end point drop by drop and watching the pink flash and disappear as it is stirred; a splash of acid gets you to colourless with nothing seen.
Activity 2.7
Take a small amount of copper oxide in a beaker and add dilute hydrochloric acid slowly, stirring. Note the colour of the solution.
What you should observe. The black powder dissolves and the solution turns blue-green.
The black copper oxide dissolves; the solution becomes blue-green with copper(II) chloride.
$$\mathrm{CuO}(s) + 2\mathrm{HCl}(aq) \rightarrow \mathrm{CuCl_2}(aq) + \mathrm{H_2O}(l)$$
Why it is placed here. A metal oxide plus an acid gives a salt and water — which is exactly the pattern of a base plus an acid in Activity 2.6. The chapter draws the conclusion directly: metallic oxides “are said to be basic oxides”. Read together with the lime water of 2.5, where a non-metallic oxide behaved like an acid, this pins down both halves of the rule.
What goes wrong. Impatience. Copper oxide is slow to dissolve in cold dilute acid; warming it gently makes the colour unmistakable inside a lesson.
Activity 2.8
Fix two nails in a cork and stand it in a 100 mL beaker. Wire the nails to a 6 V battery through a bulb and a switch. Pour in dilute HCl and switch on. Repeat with dilute sulphuric acid, then with glucose and alcohol solutions.
What you should observe. The bulb lights for both acids. It does not light for glucose or for alcohol.
The bulb glows for HCl and H₂SO₄ and stays dark for glucose and alcohol. Acid solutions conduct; those two do not.
Why this is the chapter’s turning point. A glowing bulb means charge is moving through the liquid, and in a solution charge is carried by ions. So acids must be producing ions in water. The chapter names them: $\mathrm{H^+}$ as the cation in every case, with $\mathrm{Cl^-}$, $\mathrm{NO_3^-}$, $\mathrm{SO_4^{2-}}$ or $\mathrm{CH_3COO^-}$ as the anion. The hydrogen ion is what all acids have in common — the question the whole section was asked to answer.
Glucose and alcohol matter as much as the acids do. They contain hydrogen, and they still do not conduct, which shows that merely having hydrogen in the formula is not enough. It has to be released as $\mathrm{H^+}$ in water.
What goes wrong. Nails too close together, or barely dipped, give a dim glow that is hard to call. Push them well in and keep them a couple of centimetres apart. And the same nails carried from acid into glucose without rinsing will light the bulb in the glucose — destroying the one comparison the activity exists to make.
Activity 2.9
Take about 1 g of solid NaCl in a clean, dry test tube and set up the apparatus of Fig. 2.4. Add some concentrated sulphuric acid. Test the gas that comes out of the delivery tube with dry blue litmus paper, then with wet blue litmus paper.
What you should observe. A gas comes off steadily. It leaves the dry paper blue and turns the wet paper red.
Dry blue litmus: no change. Wet blue litmus: turns red. So dry HCl gas is not acidic; its solution in water is.
$$\mathrm{NaCl}(s) + \mathrm{H_2SO_4}(l) \rightarrow \mathrm{NaHSO_4}(s) + \mathrm{HCl}(g)$$
Why. This is the cleanest experiment in the chapter and the one that settles the argument. Acidity is not a property of the molecule on its own — HCl only behaves as an acid once water is there to let it release $\mathrm{H^+}$. The chapter asks it back as in-text question 8: “Why does dry HCl gas not show acidic behaviour in the absence of water?”
What goes wrong. Everything must be genuinely dry — the tube, the salt, and above all the “dry” litmus paper, which picks up moisture from a humid room and then turns red, appearing to disprove the whole point. Keep it in a stoppered bottle until the moment it is used, and on a damp day warm it briefly first.
Activity 2.10
Take 10 mL of water in a beaker, add a few drops of concentrated H₂SO₄, swirl slowly, and touch the base of the beaker. Repeat with sodium hydroxide pellets.
What you should observe. The beaker becomes warm in both cases.
A clear rise in temperature. Dissolving an acid or a base in water is exothermic — the chapter calls it “a highly exothermic one”.
Why the order matters, and it really matters. Always add the acid to the water, slowly, with constant stirring — never water to acid. The chapter is explicit about why: if water is added to a concentrated acid, the heat generated may make the mixture splash out and cause burns, and the glass may break from local heating. Of everything in this chapter, this is the instruction to make a class repeat back to you.
What it sets up. Mixing an acid or base with water reduces the concentration of $\mathrm{H_3O^+}$ or $\mathrm{OH^-}$ per unit volume — which is what dilution means, and which leads straight into pH in the next section.
What goes wrong. Using too much acid to make the effect obvious. A few drops in 10 mL is warm enough to feel; more than that is an unnecessary hazard for no extra teaching value.
Activity 2.11
Test the pH of the solutions listed in Table 2.2 and decide the nature of each.
What you should observe. The everyday liquids spread right across the scale rather than clustering.
pH below 7 is acidic, 7 is neutral, above 7 is basic. A lower pH means a higher hydrogen-ion concentration.
Why. The universal indicator is a mixture of several indicators, so instead of a single yes/no it gives a different colour at each hydrogen-ion concentration. That is what turns “is this an acid?” into “how strong an acid is this?”, which is the question Section 2.3 is asking.
What goes wrong. Reading the colour against the wrong lighting, or trusting it to a unit. The book’s own Fig. 2.7 carries the caveat that the colours “are only a rough guide”. Treat a pH paper reading as good to about a whole unit, and do not let a class record 4.2.
Activity 2.12
Put about 2 g of soil in a test tube with 5 mL of water, shake, filter, and check the pH of the filtrate with universal indicator paper.
What you should observe. Most garden and farm soils come out somewhere between about 6 and 8.
The soil’s pH, read from the filtrate. Most plants grow best in soil that is close to neutral; strongly acidic or strongly alkaline soil suppresses growth.
Why it is worth doing locally. This is the one activity in the chapter with no fixed answer, and that is deliberate — the chapter asks what the ideal pH is “for the growth of plants in your region“, and suggests noting which plants actually grow where the sample came from. A class that tests soil from three different places has a real result that no textbook could have printed.
What goes wrong. Testing the mud rather than the filtrate. Soil particles hold colour and smear the paper, and the reading becomes unreadable. Filter properly and test the clear liquid.
Activity 2.13
Write the chemical formulae of: potassium sulphate, sodium sulphate, calcium sulphate, magnesium sulphate, copper sulphate, sodium chloride, sodium nitrate, sodium carbonate and ammonium chloride. Then identify the acid and base each could be made from, and group them into families.
| Salt | Formula | From acid | From base |
|---|---|---|---|
| Potassium sulphate | $\mathrm{K_2SO_4}$ | $\mathrm{H_2SO_4}$ | $\mathrm{KOH}$ |
| Sodium sulphate | $\mathrm{Na_2SO_4}$ | $\mathrm{H_2SO_4}$ | $\mathrm{NaOH}$ |
| Calcium sulphate | $\mathrm{CaSO_4}$ | $\mathrm{H_2SO_4}$ | $\mathrm{Ca(OH)_2}$ |
| Magnesium sulphate | $\mathrm{MgSO_4}$ | $\mathrm{H_2SO_4}$ | $\mathrm{Mg(OH)_2}$ |
| Copper sulphate | $\mathrm{CuSO_4}$ | $\mathrm{H_2SO_4}$ | $\mathrm{Cu(OH)_2}$ |
| Sodium chloride | $\mathrm{NaCl}$ | $\mathrm{HCl}$ | $\mathrm{NaOH}$ |
| Sodium nitrate | $\mathrm{NaNO_3}$ | $\mathrm{HNO_3}$ | $\mathrm{NaOH}$ |
| Sodium carbonate | $\mathrm{Na_2CO_3}$ | $\mathrm{H_2CO_3}$ | $\mathrm{NaOH}$ |
| Ammonium chloride | $\mathrm{NH_4Cl}$ | $\mathrm{HCl}$ | $\mathrm{NH_4OH}$ |
The families. A family is a group sharing one radical, and the chapter gives both directions: $\mathrm{NaCl}$ and $\mathrm{Na_2SO_4}$ are both sodium salts, while $\mathrm{NaCl}$ and $\mathrm{KCl}$ are both chloride salts. The nine above give a sulphate family of five, a sodium family of four, and a chloride family of two — and the same salt can belong to more than one family, which is the part students find surprising.
What goes wrong. Writing $\mathrm{KSO_4}$ or $\mathrm{NaSO_4}$. The sulphate ion carries two negative charges, so it needs two singly-charged cations. Every formula in the table balances charge, and that is the only rule needed.
Activity 2.14
Collect sodium chloride, potassium nitrate, aluminium chloride, zinc sulphate, copper sulphate, sodium acetate, sodium carbonate and sodium hydrogencarbonate. Dissolve each in distilled water, test with litmus, and find the pH with pH paper. Then work out which acid and base each came from.
What you should observe. Salts are not all neutral, and which way a salt goes is decided by its parents.
- Strong acid + strong base → neutral, pH 7 (NaCl, KNO₃)
- Strong acid + weak base → acidic, pH below 7 (AlCl₃, ZnSO₄, CuSO₄)
- Weak acid + strong base → basic, pH above 7 (sodium acetate, Na₂CO₃, NaHCO₃)
Why. The weaker parent is the one that loses. A salt of a weak base and a strong acid leaves behind more $\mathrm{H^+}$ than $\mathrm{OH^-}$ in solution, and the solution is acidic — and the mirror image for the other case.
Why distilled water is specified. Tap water is not neutral; in many places it is distinctly alkaline. Use it here and every salt appears basic, which destroys the whole comparison. This is the reason the book says “use distilled water only”, and it is the most common way this activity fails.
What goes wrong. Aluminium chloride is worth warning students about: it is strongly acidic in solution and fumes in moist air. Handle the solid quickly and keep the bottle closed.
Activity 2.15
Heat a few crystals of copper sulphate in a dry boiling tube. Note the colour afterwards and look for water droplets in the tube. Then add 2–3 drops of water to the heated sample.
What you should observe. The blue crystals turn white, water condenses on the cool upper part of the tube, and adding water brings the blue straight back.
Blue crystals turn white; water droplets collect in the tube; adding water restores the blue.
$$\mathrm{CuSO_4}\!\cdot\!5\mathrm{H_2O}(s) \xrightarrow{\ \text{heat}\ } \mathrm{CuSO_4}(s) + 5\mathrm{H_2O}(g)$$
Why. Crystals that look and feel perfectly dry are not. Copper sulphate crystals carry five molecules of water per formula unit, chemically bound in the crystal — the water of crystallisation — and it is that water which makes them blue. Heat drives it off and the anhydrous salt is white. Put the water back and the blue returns, which shows the change is reversible and that the water was genuinely part of the crystal, not spilled on it.
What it sets up. This is the idea the chapter needs for Plaster of Paris a page later, where the odd-looking $\mathrm{CaSO_4}\cdot\frac{1}{2}\mathrm{H_2O}$ has half a water molecule per formula unit — because two formula units share one water molecule between them.
What goes wrong. Two things. Heating too hard scorches the salt and it goes greyish rather than clean white, and the change then looks like decomposition rather than dehydration. And the droplets are easy to miss if the tube is held level — tilt it slightly so the mouth is lower, and they collect where they can be seen.
Common mistakes
- Concluding from Activity 2.1 that red litmus detects acids. It cannot; it is already red. Red litmus reports bases only, and blue litmus reports acids only — which is exactly what the chapter’s in-text question about three unlabelled test tubes is testing.
- Treating Activity 2.4 as a repeat of 2.3. Getting hydrogen from a base is the surprise, and it only happens because zinc is amphoteric. A student who files both under “acids give hydrogen with metals” has taken the wrong lesson from the pair.
- Reporting that lime water failed in Activity 2.5 because it went clear again. It went milky first. The clearing is a second reaction with excess carbon dioxide, and it is worth showing on purpose.
- Skipping the glucose and alcohol in Activity 2.8. The dark bulb is half the experiment. Without it there is nothing to rule out “anything dissolved conducts”, and the conclusion about $\mathrm{H^+}$ does not actually follow.
- Blaming Activity 2.9 on the apparatus when the dry litmus turns red. Nine times out of ten the paper picked up moisture from the room. The experiment is unusually sensitive to this, because a damp “dry” strip reverses the result it exists to demonstrate.
- Adding water to concentrated acid in Activity 2.10. Always acid into water, slowly, stirring. This is the one instruction in the chapter where getting it the wrong way round causes injury rather than a wrong answer.
- Using tap water in Activity 2.14. It is usually alkaline, so every salt reads basic and the three-way classification collapses. The book specifies distilled water for this reason.
Practise next
- The end-of-chapter Exercises for this chapter, where Question 8 asks why dry HCl gas is not acidic — Activity 2.9 exactly — and Question 10 asks about fizzing with HCl against acetic acid, which is the comparison built into Activity 2.3.
- The in-text questions of this chapter, including the three-test-tube problem that Activity 2.1 prepares for and the “metal compound A” question that follows from Activity 2.5.
- The activities of Chapter 1, where Activity 1.10’s precipitation of barium sulphate is the same double displacement that makes lime water turn milky here.
- Chapter 3, Metals and Non-metals, where the metal oxides of Activity 2.7 become a full account of basic and amphoteric oxides.
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.