NCERT Class 10 Science Chapter 3 Activities — Metals and Non-metals
NCERT Class 10 Science — Metals and Non-metals, Activities 3.1 to 3.14 on pages 37 to 53.
Also here. Solutions to the end-of-chapter questions for this chapter and the in-text questions are on their own pages.
The fourteen activities in this chapter fall into three clean groups, and it is worth saying which is which before starting, because they are doing quite different jobs.
- Activities 3.1 to 3.7 are about physical properties — lustre, hardness, malleability, ductility, conductivity, sonority — and they end by testing non-metals the same way to see how the two groups differ.
- Activities 3.8 to 3.12 are about chemical properties, and they build one thing: the reactivity series. Each activity ranks the metals against a different reagent — oxygen, water, acid, each other’s salts — and the series is what survives when all four rankings agree.
- Activities 3.13 and 3.14 are separate: the properties of ionic compounds, and the conditions under which iron rusts.
Key insight. The chapter says on page 45 why Activity 3.12 matters more than the three before it. Oxygen, water and acid each fail to rank all the metals, because some metals do not react with them at all — silver and gold do nothing with any of the three. Displacement works on every pair, so it is the only test that orders the whole list. If time is short, 3.12 is the one to keep.
Activity 3.1
Take samples of iron, copper, aluminium and magnesium and note the appearance of each. Then clean each surface with sandpaper and look again.
What you should observe. Dull before, bright after.
Freshly cleaned metal surfaces are shiny — metals have lustre. The dullness beforehand is a layer of oxide, not the metal itself.
Why it comes first. This is the observation the rest of the chapter keeps coming back to. The tarnish is a reaction product, and by page 42 the chapter is explaining that magnesium, aluminium, zinc and lead all carry a thin oxide layer that protects the metal underneath from further attack.
What goes wrong. Sandpaper that is too fine, or not enough rubbing. Aluminium in particular has a tough oxide film, and a token rub leaves it looking much as it did.
Activity 3.2
Try to cut pieces of iron, copper, aluminium and magnesium with a sharp knife. Then hold a piece of sodium with tongs, dry it between filter paper folds, put it on a watch-glass and try to cut that. The book cautions to handle sodium with care.
What you should observe. The knife makes no impression on the first four. Sodium cuts like firm cheese, and the fresh face is bright and shiny for a moment before it dulls.
Most metals are too hard to cut with a knife. Sodium is the exception — the alkali metals are soft enough to cut, and have low densities and low melting points.
Why the sodium is dried first. It is kept under kerosene because it reacts with the moisture and oxygen in air. Filter paper takes off the kerosene so the cut surface can actually be seen — and the speed with which that bright surface dulls is itself the point.
What goes wrong. This is the most hazardous activity in the chapter and is a demonstration, not a class exercise. Never touch sodium with bare hands, keep the piece small, and keep water well away from the bench.
Activity 3.3
Place a piece of iron, zinc, lead or copper on an iron block and strike it four or five times with a hammer. Repeat with the others and record how the shape changes.
What you should observe. Each flattens rather than shattering, and they flatten by different amounts.
The metals are beaten into thin sheets rather than breaking — this is malleability. Lead flattens most easily of the four; gold and silver are the most malleable metals of all.
Why. Metallic bonding lets layers of atoms slide over one another without the structure falling apart, so a blow changes the shape instead of cracking it. Compare this with Activity 3.13, where ionic solids shatter under the same treatment — that contrast is one of the chapter’s main points about the two kinds of bonding.
What goes wrong. Using a cast iron block that chips, or striking hard enough to send fragments flying. Firm, controlled blows are enough; this does not need force.
Activity 3.4
List the metals whose wires you have seen in daily life.
Copper and aluminium in electrical wiring, steel in cables and fencing, gold and silver in jewellery and fine electronics, tungsten in old filament lamps. Drawing a metal into a wire is ductility, and gold is the most ductile — a wire about 2 km long can be drawn from a single gram.
Why it is an activity and not just a sentence. It asks students to notice that the same two metals turn up almost everywhere a wire is needed, which sets up the question the chapter asks next: why copper and aluminium in particular? The answer is conductivity, tested in Activity 3.6.
A related question worth asking here. Household wires are coated in PVC or rubber. The chapter raises it on page 39 and does not answer it — the coating is an insulator, keeping the current in the wire and away from anyone touching it.
Activity 3.5
Clamp an aluminium or copper wire on a stand as in Fig. 3.1. Fix a pin to the free end with wax. Heat the wire near the clamp with a spirit lamp, candle or burner, and watch.
What you should observe. After a little while the wax softens and the pin drops, though the flame never went anywhere near it.
The wax melts and the pin falls, showing heat has travelled along the wire. The metal itself does not melt — metals are good conductors of heat and have high melting points.
Why both halves matter. Heat conducted along the wire is why cooking vessels are made of metal. The wire not melting is why they survive the flame. Two of the chapter’s facts come out of one experiment.
Exceptions the chapter is careful to add. Most metals have high melting points, but gallium and caesium have such low ones that they will melt in the palm of your hand.
What goes wrong. Too much wax. A thick blob insulates the pin and takes so long to soften that the class loses interest; a small dab drops within a minute or two.
Activity 3.6
Set up the circuit of Fig. 3.2 and put the metal to be tested between terminals A and B. Does the bulb glow?
What you should observe. It glows for every metal tried.
The bulb glows — metals are good conductors of electricity.
Why this is the pair to Activity 3.5. Conduction of heat and conduction of electricity have the same cause in a metal: electrons free to move through the structure. That is why the metals which conduct heat best also conduct electricity best.
What it sets up. Run the same circuit with non-metals in Activity 3.7 and the bulb stays dark — except for graphite, which is the exception the chapter wants noticed.
What goes wrong. Testing a tarnished sample. Oxide layers are poor conductors, and an uncleaned piece of aluminium can give a dim glow or none, which looks like a result and is not.
Activity 3.7
Collect carbon (coal or graphite), sulphur and iodine, and carry out Activities 3.1 to 3.4 and 3.6 with them.
What you should observe. They behave the opposite way to the metals on almost every test — and there are two clear exceptions.
Non-metals are dull, brittle, non-malleable, non-ductile and do not conduct electricity. Two exceptions the chapter names: iodine is lustrous, and graphite conducts electricity.
Why the exceptions are the interesting part. They show that “metal” and “non-metal” are useful groupings rather than exact rules, which is precisely the chapter’s reason for moving on to chemical properties as a better basis for classification. Carbon also makes the point about allotropes — diamond and graphite are the same element, yet diamond is the hardest natural substance known and graphite conducts.
What goes wrong. Hammering sulphur or iodine expecting a sheet: they crumble. That is the result, not a failure — brittleness is what distinguishes them from the metals of Activity 3.3.
Activity 3.8
Burn a magnesium ribbon, collect the ash, dissolve it in water and test with red and blue litmus. Then burn sulphur powder, collect the fumes in an inverted test tube, add water, shake, and test that solution with litmus too.
What you should observe. The two solutions behave in opposite ways.
Magnesium oxide solution turns red litmus blue — basic. Sulphur dioxide solution turns blue litmus red — acidic.
$$2\mathrm{Mg}(s) + \mathrm{O_2}(g) \rightarrow 2\mathrm{MgO}(s) \qquad \mathrm{MgO}(s) + \mathrm{H_2O}(l) \rightarrow \mathrm{Mg(OH)_2}(aq)$$ $$\mathrm{S}(s) + \mathrm{O_2}(g) \rightarrow \mathrm{SO_2}(g) \qquad \mathrm{SO_2}(g) + \mathrm{H_2O}(l) \rightarrow \mathrm{H_2SO_3}(aq)$$
Why this is the chapter’s real definition. Lustre and malleability are useful but leaky — iodine is lustrous, graphite conducts. This test is not leaky: metal oxides are basic, non-metal oxides are acidic, and that is a chemical distinction that holds. It also connects straight back to Chapter 2, where copper oxide neutralised hydrochloric acid and lime water absorbed carbon dioxide.
What goes wrong. Losing the sulphur dioxide. It is the fumes that matter, and an inverted tube held loosely over the flame lets most of them escape, leaving a solution too weak to change litmus convincingly. Hold the tube close and steady. Do this in good ventilation — sulphur dioxide is an irritant.
Activity 3.9
Hold each metal sample with tongs and try burning it over a flame. Collect any product, let it cool, and record which metals burn easily, what colour the flame is, how the surface looks afterwards, and whether the products dissolve in water. Then rank the metals by reactivity towards oxygen. The book asks for the teacher’s assistance and eye protection.
What you should observe. They differ enormously, and most do not burn at all.
- Magnesium burns with a dazzling white flame.
- Copper does not burn, but its surface is coated with black copper(II) oxide.
- Iron does not burn on heating, although iron filings burn vigorously when sprinkled into a flame.
- Silver and gold do not react with oxygen even at high temperatures.
$$2\mathrm{Cu}(s) + \mathrm{O_2}(g) \rightarrow 2\mathrm{CuO}(s) \qquad 4\mathrm{Al}(s) + 3\mathrm{O_2}(g) \rightarrow 2\mathrm{Al_2O_3}(s)$$
Why iron filings burn when an iron nail does not. Surface area. The filings expose far more iron to the oxygen at once, and the heat released is enough to keep the reaction going instead of being conducted away into the bulk of the metal.
Why sodium and potassium are not in this test. They are so reactive that they catch fire if simply left in the open, which is why the chapter says they are stored under kerosene.
What goes wrong. Concluding “copper does not react” from the absence of a flame. It does — the black coating is the product. Burning and reacting are not the same thing, and this is the activity where that distinction gets made.
Activity 3.10
Put small pieces of the same metal samples separately into beakers half-filled with cold water. Which react? Rank them by reactivity with cold water. Did any catch fire? Did any float? Teacher’s assistance again.
What you should observe. A clear ladder, from violent to nothing at all.
- Potassium and sodium react violently with cold water, and the hydrogen released catches fire.
- Calcium reacts less violently — not enough heat to ignite the hydrogen — and starts floating.
- Magnesium does not react with cold water, only with hot.
- Aluminium, iron and zinc react with steam only.
- Lead, copper, silver and gold do not react with water at all.
$$2\mathrm{Na}(s) + 2\mathrm{H_2O}(l) \rightarrow 2\mathrm{NaOH}(aq) + \mathrm{H_2}(g) + \text{heat}$$ $$\mathrm{Ca}(s) + 2\mathrm{H_2O}(l) \rightarrow \mathrm{Ca(OH)_2}(aq) + \mathrm{H_2}(g)$$ $$3\mathrm{Fe}(s) + 4\mathrm{H_2O}(g) \rightarrow \mathrm{Fe_3O_4}(s) + 4\mathrm{H_2}(g)$$
Why calcium floats — the detail students get wrong. Not because calcium is less dense than water. Bubbles of hydrogen stick to the metal’s surface and carry it up. The chapter says this explicitly, and it is a favourite examination question.
What goes wrong. Sodium in an open beaker is genuinely dangerous — the hydrogen ignites and the piece can skate about and spit hot alkali. Use a small piece behind a screen, and treat this as a demonstration only. The chapter’s own caution in Activity 3.11 is not to take sodium and potassium at all.
Activity 3.11
Clean the metal samples — but not sodium or potassium — and put each in a test tube of dilute hydrochloric acid with a thermometer bulb dipped in the acid. Watch the rate of bubbling and the temperature. Rank them by reactivity with dilute acids.
What you should observe. The same order as in water, measured two ways at once.
Magnesium reacts most vigorously and gives the highest temperature rise, then aluminium, then zinc, then iron. Copper does not react.
$$\text{Metal} + \text{Dilute acid} \rightarrow \text{Salt} + \text{Hydrogen}$$ $$\mathrm{Mg} + 2\mathrm{HCl} \rightarrow \mathrm{MgCl_2} + \mathrm{H_2} \qquad 2\mathrm{Al} + 6\mathrm{HCl} \rightarrow 2\mathrm{AlCl_3} + 3\mathrm{H_2}$$ $$\mathrm{Zn} + 2\mathrm{HCl} \rightarrow \mathrm{ZnCl_2} + \mathrm{H_2} \qquad \mathrm{Fe} + 2\mathrm{HCl} \rightarrow \mathrm{FeCl_2} + \mathrm{H_2}$$
Why the thermometer is in there. It turns a subjective judgement — “which fizzed more?” — into a number. Two groups can disagree about bubbling rates; they cannot disagree about which tube got hottest.
The exception worth stating. Nitric acid does not give hydrogen with metals. It is a strong oxidising agent and oxidises the hydrogen to water while itself being reduced to one of the nitrogen oxides. Magnesium and manganese with very dilute nitric acid are the only exceptions.
What goes wrong. Not cleaning the samples, which the chapter puts first in the instructions. A tarnished metal has an oxide layer that must be dissolved before the metal itself can react, so it appears sluggish and the ranking comes out wrong.
Activity 3.12
Put a clean copper wire into iron sulphate solution, and an iron nail into copper sulphate solution. Look again after 20 minutes.
What you should observe. One changes; the other does nothing at all.
The iron nail in copper sulphate becomes coated with brown copper, and the blue solution fades. The copper wire in iron sulphate is unchanged. So iron is more reactive than copper.
$$\mathrm{Fe}(s) + \mathrm{CuSO_4}(aq) \rightarrow \mathrm{FeSO_4}(aq) + \mathrm{Cu}(s)$$
Why the tube where nothing happens is the important one. A single displacement only tells you that iron can displace copper. It is the failure of the reverse reaction that proves the order, rather than merely suggesting it. The chapter’s rule on page 45: if metal A displaces metal B from its solution, A is more reactive than B.
Why this activity outranks 3.9 to 3.11. Oxygen, water and acid all leave gaps — silver and gold do nothing with any of them, so those three tests cannot place them. Displacement works on every pair of metals, which is why the chapter says displacement reactions “give better evidence about the reactivity of metals” and builds the reactivity series from them.
What goes wrong. Twenty minutes is a genuine minimum, and the nail must be clean. It is also worth keeping both tubes side by side: the unchanged one is easy to forget, and it is half the evidence.
Activity 3.13
Take sodium chloride, potassium iodide, barium chloride or another salt. Note the physical state. Heat a little on a metal spatula directly in a flame. Try dissolving the samples in water, in petrol and in kerosene. Then put electrodes into a solution of each and see whether current flows.
What you should observe. Four properties, all pointing the same way.
- Physical nature: hard solids, but brittle — they break when pressure is applied.
- Melting point: high; they do not melt in a spatula flame.
- Solubility: soluble in water, insoluble in petrol and kerosene.
- Conductivity: their solutions conduct; the solids do not.
Why they all follow from one cause. Ionic compounds are held together by strong attraction between oppositely charged ions. That strength is why the melting points are high and why the solids are hard. The brittleness is the same structure seen from the other side — force the layers to slide and like charges come together, and the crystal splits. And conduction needs ions free to move, which they are in solution or in the melt, and are not in a rigid solid.
What goes wrong. Reporting that ionic compounds do not conduct, full stop, after testing only the solid. Testing solid and solution, and getting different answers, is the whole design of this activity.
Activity 3.14
Put clean iron nails in three test tubes, A, B and C. Pour ordinary water into A and cork it. Pour boiled distilled water into B, add about 1 mL of oil to float on top, and cork it. Put anhydrous calcium chloride into C and cork it. Leave them a few days.
What you should observe. Only one nail rusts.
Tube A rusts. Tubes B and C do not. Iron needs both air (oxygen) and water to rust — either one alone is not enough.
Why each tube is built the way it is. This is a controlled experiment and every detail is doing a job, which the chapter’s own caption for Fig. 3.13 spells out:
- A has both air and water — the control where rusting is expected.
- B has water but no air. The water is boiled to drive out dissolved air, and the oil floats on top to stop more dissolving back in. Without the oil, the boiling would be undone within hours.
- C has air but no water. Anhydrous calcium chloride absorbs any moisture, leaving the air inside genuinely dry.
What goes wrong. Three things, each of which quietly ruins the result. Using unboiled water in B leaves dissolved oxygen and B rusts too. Forgetting the oil layer does the same thing more slowly. And corking any tube loosely lets moist air into C. If all three tubes rust, the experiment did not disprove the conclusion — it failed to exclude the variables.
What it sets up. Everything the chapter then says about preventing corrosion — painting, oiling, greasing, galvanising, anodising, alloying — works by removing one or other of the two things tube A has.
Common mistakes
- Treating the dullness in Activity 3.1 as dirt. It is an oxide layer, formed by reaction with the air, and the chapter later relies on that layer to explain why aluminium resists corrosion.
- Concluding from Activity 3.7 that all non-metals are dull and non-conducting. Iodine is lustrous and graphite conducts. The chapter names both exceptions on purpose, and they are why it moves to chemical properties for a reliable classification.
- Saying copper “does not react” in Activity 3.9 because it does not burn. It reacts — the black coating is copper(II) oxide. Not burning and not reacting are different claims.
- Explaining the floating calcium of Activity 3.10 by density. It floats because hydrogen bubbles stick to its surface, which the chapter states in as many words.
- Expecting hydrogen from nitric acid in Activity 3.11. It is an oxidising agent and gives nitrogen oxides instead, except with very dilute acid on magnesium or manganese.
- Ignoring the tube where nothing happened in Activity 3.12. The unchanged copper wire in iron sulphate is what makes the ordering conclusive rather than merely suggestive.
- Testing only solid salt in Activity 3.13 and concluding ionic compounds do not conduct. The solid does not; the solution does. The difference is the result.
- Letting tubes B and C rust in Activity 3.14 and reporting it as the answer. Unboiled water, a missing oil layer or a loose cork will each do it. When all three rust, the controls failed, not the theory.
Practise next
- The end-of-chapter Exercises for this chapter, where the reactivity series built in Activities 3.9 to 3.12 is used to predict displacement reactions, and the rusting conditions of Activity 3.14 return as questions on galvanisation.
- The in-text questions of this chapter, which ask why sodium is stored in kerosene and why copper vessels lose their shine — both answered directly by Activities 3.2 and 3.9.
- The activities of Chapter 1, where Activity 1.9 is this chapter’s Activity 3.12 met for the first time, as an example of displacement before the reactivity series existed to explain it.
- Chapter 4, Carbon and its Compounds, which takes the one non-metal the chapter kept flagging as exceptional and gives it a chapter of its own.
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.