NCERT Solutions Class 10 Science Chapter 3 Metals and Non-metals

Why these are exceptions. Each answer here goes against the general picture of metals, so it helps to know the reason rather than just the name. The general rule is that metals are solid, hard and good conductors, so you should be able to explain why each of these four does not fit that rule.

A liquid metal exists because mercury has very weak forces between its atoms, so it cannot stay solid at room temperature; this is why thermometers and old barometers used it. Sodium and potassium are cut easily because they are alkali metals with only one loosely held outer electron, which makes their metallic bonding weak and the metal soft. Silver conducts heat best because its free electrons move through the lattice with very little resistance, which is also why silver is the best electrical conductor.

Lead and mercury sit at the poor end of the conductivity list. Their electron clouds carry heat less freely than silver or copper, so cooking pots are never made of lead. Knowing the reason lets you answer a twisted version of the question, such as “name a metal used in thermometers” (mercury) or “why is silver not used for everyday wiring” (it is costly, so copper is used instead).

Answer: Mercury (liquid), sodium (knife-soft), silver (best heat conductor), lead (poor heat conductor).

The atom-level picture. The reason metals are malleable and ductile is the same in both cases, so it is worth understanding once. In a metal the positive ions sit in a regular pattern inside a “sea” of free electrons. When you hammer or pull the metal, layers of ions slide past one another, but the free electrons keep flowing around them and hold the whole structure together. So the shape changes without the metal cracking.

This is very different from an ionic or brittle solid, where pushing the layers brings like charges next to each other; they repel and the crystal shatters. That is why chalk, glass and most non-metals break instead of bending. The practical pay-off is huge: malleability lets us roll steel into car body panels and beat gold into leaf thinner than paper, while ductility lets us pull copper into the kilometres of wire that carry electricity to your home. Both properties are decided by the metallic bonding, so any factor that weakens that bonding (such as impurities) reduces malleability and ductility.

Answer: Malleable = can be beaten into sheets; ductile = can be drawn into wires; both follow from sliding layers in metallic bonding.

Reactivity-first reasoning. The cleanest way to answer any “why is this metal stored under X” question is to first place the metal on the reactivity series and then ask what it would react with if left open. Sodium and potassium sit right at the top of the series, so they react with both the oxygen and the water vapour present in ordinary air.

If a piece of sodium is left in the open, a fresh cut surface dulls almost at once as it forms sodium oxide, and on a humid day enough heat can build up from the reaction with moisture to set the hydrogen on fire. Kerosene is a non-polar hydrocarbon that does not react with sodium and is also denser handling-wise convenient, so the metal stays fully covered and air-free. Note that potassium and lithium are stored the same way for the same reason. This is a one-mark question that students lose marks on by writing only “it is reactive” without naming what it reacts with (oxygen and moisture) and why kerosene works (it keeps both away).

Answer: Sodium reacts violently with air and moisture; kerosene keeps both away, preventing fire and oxidation.

Read it off the reactivity series. The single idea behind all three equations is position in the reactivity series, so train yourself to predict the product before balancing. Potassium is near the top, so it tears water apart even when cold and the reaction is fast enough to ignite the hydrogen. Calcium is a little lower, so it still reacts with cold water but calmly, and you actually see the calcium float as hydrogen bubbles stick to it.

Iron is well below both, so cold water has no effect and even hot water barely reacts; only high-temperature steam pushes the reaction, and the product is the mixed oxide Fe₃O₄ rather than a hydroxide. A neat balancing check for the iron equation: 3 Fe on the left match 3 Fe in Fe₃O₄, and 4 H₂O supply 4 O and 8 H, which match 4 O in Fe₃O₄ and 8 H in 4 H₂. Getting the product right first, then balancing, stops the most common error of writing a hydroxide for the steam reaction.

Answer: 3Fe + 4H₂O → Fe₃O₄ + 4H₂, Ca + 2H₂O → Ca(OH)₂ + H₂, 2K + 2H₂O → 2KOH + H₂ (violent).

Logic-table method. Treat the table as a set of yes/no clues and turn each “displacement” into a ranking statement, then combine them. B displaces iron, which means B sits above iron in the series; A cannot displace iron but can displace copper, so A sits between iron and copper; C displaces only the least reactive metal (silver), so C sits just above silver; and D displaces nothing, so it is the lowest of the four.

Lining these up gives B above A above C above D, which is exactly the decreasing order asked for. For part (ii), the prediction follows from B being above copper: B kicks copper out of copper sulphate, so the tell-tale signs are the fading of the blue colour (because Cu↑2+ ions leave the solution) and a reddish-brown copper coating on B. A good habit on these questions is to name the colour change, because that is what an examiner uses to check you truly understand displacement rather than just memorising an order. The same reasoning predicts that A would also displace copper, but C and D would not.

Answer: B > A > C > D; B displaces copper from CuSO₄ (blue fades, reddish copper deposits).

Why iron gives the +2 salt. The reaction of metals with dilute acids is really a displacement of hydrogen, so it is worth knowing both the gas test and the salt produced. With dilute H₂SO₄, iron forms iron(II) sulphate, FeSO₄, where iron carries a +2 charge; dilute sulphuric acid is not a strong enough oxidiser to push iron to the +3 state, so you should never write Fe₂(SO₄)₃ for this dilute reaction.

The hydrogen gas itself is confirmed by the pop test, in which the gas burns explosively with oxygen in a small sample. A second point examiners check is that very unreactive metals such as copper, silver and gold do not react with dilute HCl or dilute H₂SO₄ at all, because they lie below hydrogen in the reactivity series and so cannot displace it. So the phrase “a reactive metal” in the question is doing real work: only metals above hydrogen liberate H₂ from dilute acids.

Answer: Hydrogen gas; Fe + H₂SO₄ → FeSO₄ + H₂ (iron forms the +2 salt).

Spot the colour clue. Displacement questions are graded as much on the observation as on the equation, so always state the colour change. Here zinc sits above iron in the activity series, so zinc goes into solution as colourless Zn↑2+ while iron comes out as the metal. As the green Fe↑2+ ions leave the solution, the characteristic pale-green colour of iron(II) sulphate fades towards colourless.

At the same time a dull grey layer of iron settles on the zinc piece, which is the visible proof that the metal has been displaced. A handy way to be sure of the direction is to ask “is the added metal higher in the series than the metal in the salt?” If yes, displacement happens; if no (for example, copper added to zinc sulphate), nothing happens. This single check stops students from writing impossible reactions in which a less reactive metal displaces a more reactive one.

Answer: Zn + FeSO₄ → ZnSO₄ + Fe; green colour fades, grey iron deposits on zinc.

Octet-driven thinking. Every ionic-formation question is really about each atom reaching a stable octet, so reason from electron count rather than memorising formulas. Sodium has the configuration 2, 8, 1, so losing one electron leaves the stable 2, 8 of neon and gives Na↑+. Magnesium is 2, 8, 2, so it loses two electrons to reach the same 2, 8 and becomes Mg↑2+. Oxygen is 2, 6 and needs two more electrons to reach the 2, 8 of neon, so it becomes O↑2-.

Now the formula is just charge balancing. One oxide ion carries -2, so it needs +2 of metal charge to neutralise it. Two sodium ions (2 × +1) do that, giving Na₂O; a single magnesium ion (+2) does it alone, giving MgO. This is why the two oxides have different formulas even though both contain oxygen. A useful self-check on the dot diagrams is that after transfer every ion should show a complete octet of eight dots, and the total positive charge must equal the total negative charge. Doing it this way means you can write the formula of any metal oxide without rote learning.

Answer: Na↑+, Mg↑2+ and O↑2- ions; charge balance gives Na₂O and MgO.

Lattice-energy angle. The high melting point of an ionic solid comes straight from the strength of its lattice, so it helps to connect the idea to “lattice energy”. Lattice energy is the energy released when gaseous ions come together to form one mole of the solid, and it is large because the + and - ions attract strongly and pack closely. To melt the solid you must supply enough thermal energy to overcome a big chunk of that lattice energy, which only happens at a high temperature.

Two factors make the lattice energy, and therefore the melting point, even higher: larger ionic charges and smaller ionic sizes. That is why magnesium oxide (Mg↑2+, O↑2-, both doubly charged) melts far higher (around 2800 ^°C) than sodium chloride (singly charged ions). The same strong lattice also explains why ionic compounds are hard and brittle, and why they conduct electricity only when molten or dissolved, since the ions become free to move only then. So one idea, strong electrostatic attraction in a lattice, explains a whole family of properties.

Answer: Strong electrostatic forces in the ionic lattice need a large amount of heat to break, so melting points are high (higher still for small, highly charged ions).

Think of it as a sorting funnel. The three terms make sense if you picture how a metal travels from the ground to a factory. At the widest part of the funnel are minerals, the huge variety of naturally occurring compounds in the crust. Most of them are too poor in metal or too hard to process to be useful, so only a fraction make it through the funnel.

The ones that contain enough metal to extract at a profit are the ores; the economics matter, because a low-grade mineral may technically contain a metal yet cost more to process than the metal is worth. When an ore is dug up it never comes clean: it is mixed with sand, clay and rocky matter called gangue, which has to be separated out (by methods such as washing or magnetic separation) before the actual extraction begins. So the chain is mineral → ore (rich enough to be worth it) → remove gangue → extract metal. Keeping this order in mind makes the rest of the metallurgy section much easier to follow.

Answer: Mineral (any natural compound) ⊃ ore (profitable to extract); gangue is the rocky impurity removed before extraction.

Tie it to the reactivity series. The simplest way to answer “which metals are found free” is to look at the very bottom of the reactivity series, because reactivity and the way a metal occurs in nature are directly linked. Metals near the bottom, such as gold, silver and platinum, are so unreactive that they do not combine with oxygen, water or sulphur over geological time, and so they survive as shiny lumps of pure metal, which is exactly why ancient civilisations could find and use gold without any chemistry.

Metals higher up the series are far more reactive and are always locked away as oxides, sulphides or carbonates; you will never find pure sodium or magnesium lying in the ground. Copper is an interesting in-between case: mostly it occurs as compounds, but small amounts of native copper are found because it is only moderately reactive. So the rule to remember is simple, the lower a metal sits in the reactivity series, the more likely it is to be found in the free state.

Answer: Gold and silver (least reactive metals) are found in the free state.

Match the method to the reactivity. The key to extraction questions is that the choice of reducing method depends on where the metal sits in the reactivity series, so always ask “how reactive is this metal” first. For metals in the middle of the series, the oxide is cheap to reduce by heating with carbon, which grabs the oxygen and leaves the metal; this is how iron is made in a blast furnace and how zinc is obtained from zinc oxide.

For metals at the top of the series, carbon is not a strong enough reducing agent, because these metals hold on to oxygen too tightly. Instead the molten oxide or chloride is reduced electrolytically: the metal ions pick up electrons at the cathode and are deposited as the pure metal. Aluminium, sodium, calcium and magnesium are all extracted this way. So the single idea “reduction of the oxide” splits into two practical routes, carbon reduction for moderately reactive metals and electrolytic reduction for highly reactive ones, and naming the right route earns full marks.

Answer: Reduction of the oxide: carbon (coke) reduction for moderately reactive metals, electrolytic reduction for highly reactive metals.

Above-or-below check. This grid question is fast once you fix the order magnesium > zinc > copper and apply one rule: a metal displaces another from its oxide only if it is more reactive. Run through each pairing asking “is the added metal above the oxide's metal in the series?” If yes, write the displacement; if no, write no reaction.

Magnesium is at the top, so it displaces both zinc and copper from their oxides, giving two reactions. Zinc is in the middle, so it displaces only copper (which is below it) but cannot touch magnesium oxide. Copper is at the bottom, so it displaces nothing. A metal never reacts with its own oxide, which removes the three diagonal cells at once. This leaves exactly three displacement reactions, the same three NCERT expects. The grid habit is worth keeping because the board often dresses the same idea up with different metals, and the method does not change.

Answer: Three displacements: ZnO + Mg, CuO + Mg, CuO + Zn; everything else is no reaction.

Noble metals resist corrosion. The metals that resist corrosion are exactly the noble metals at the bottom of the reactivity series, so this is one more property that follows from reactivity. Gold, silver and platinum hardly react with oxygen, water or most acids, so the surface stays metallic and shiny instead of turning into an oxide or sulphide layer; this is why they are prized for ornaments, coins and high-quality electrical contacts.

There is one subtlety to note: silver does slowly tarnish, forming a thin black coat of silver sulphide when it meets traces of sulphur compounds in air, but this is surface tarnishing rather than the deep corrosion that destroys iron. Aluminium is a different case again; it is reactive, yet it appears not to corrode because it quickly forms a thin, tough, sticky layer of aluminium oxide that seals the metal underneath and stops further attack. So the honest answer to “which metals do not corrode easily” is the noble metals, gold, silver and platinum, with aluminium protected by its own oxide coat.

Answer: Gold, silver and platinum (the least reactive metals) do not corrode easily.

Why we bother alloying. Alloys exist because pure metals are often too soft, too reactive or too weak for real use, so understanding the purpose makes the definition stick. Mixing in another element disturbs the regular arrangement of atoms in the metal, which makes it harder for the layers to slide and so makes the alloy harder and stronger than the pure metal; this is why steel is used for tools and buildings rather than soft pure iron.

Alloying can also lower the melting point (solder, an alloy of lead and tin, is used to join wires), improve corrosion resistance (stainless steel resists rust because of added chromium and nickel), or change appearance (brass looks golden). Two special points the board likes: an amalgam is an alloy in which one of the metals is mercury, and the alloy of a metal with a non-metal, such as iron with carbon in steel, still counts as an alloy. So an alloy is best described as a deliberately made homogeneous mixture designed to give a metal better properties.

Answer: An alloy is a homogeneous mixture of metals (or a metal and a non-metal), made to improve hardness, strength or corrosion resistance.

One quick scan. For an MCQ like this you do not need to balance every equation; just place the added metal and the salt's metal on the reactivity series and keep only the pair where the added metal is higher. In (a), (b) and (c) the added metal (copper, aluminium, silver) is below the salt's metal (sodium, magnesium, iron respectively), so those are ruled out instantly.

Only (d) has the added metal above the salt's metal, since copper lies above silver, so copper pushes silver out of silver nitrate and a shiny grey deposit of silver forms on the copper while the solution turns blue from Cu↑2+. That blue colour and silver deposit is the classic demonstration of displacement, and it is the same reaction used to silver-plate objects. So the answer is (d), and the trick of scanning for “added metal higher than salt metal” solves the whole question in seconds.

Answer: (d) AgNO₃ + copper; copper displaces silver, solution turns blue.

Barrier thinking. The cleanest way to read a rust-prevention MCQ is to ask of each option “does it keep oxygen and moisture off the iron?” Grease leaves an oily water-repelling film, paint dries into a sealing skin, and a zinc coating both seals the surface and offers sacrificial protection, so all three answer “yes”. That immediately points to option (d).

It is worth knowing why zinc is special among the three. Even if the zinc layer gets scratched, the iron still does not rust, because zinc is more reactive than iron and so corrodes in its place; this is called sacrificial protection and is the principle behind galvanised sheets and the zinc blocks bolted to ships. Grease and paint, by contrast, stop protecting the metal the moment the film is broken. For a frying pan that is heated, a heat-stable coating matters, but as a general rust-prevention list the NCERT answer is all of the above.

Answer: (d) All of the above; each method blocks air and moisture, and zinc also gives sacrificial protection.

Decode the three clues together. This MCQ packs three clues into one sentence, and only calcium satisfies all of them, so test each option against all three. The compound has a high melting point, which suggests an ionic (metal) oxide; it is soluble in water, which narrows it to the oxides of very reactive metals; and “element reacts with oxygen” implies the element is the metal itself.

Carbon and silicon fail at once because they are non-metals: carbon dioxide is a gas, and silicon dioxide, though high-melting, is insoluble in water. Iron oxide is high-melting but insoluble, so it fails the second clue. Calcium ticks every box: CaO is ionic and high-melting, and it reacts with water to give the alkali Ca(OH)₂, which is the slaking of lime you see on construction sites. So the only element that fits all three clues is calcium, option (a).

Answer: (a) Calcium; CaO is high-melting and soluble, giving Ca(OH)₂ in water.

Reactivity, not cost, is the reason. The tempting wrong answer is the cost option, but the real reason is reactivity, so reason it through. Zinc sits above tin in the activity series, which means zinc reacts more readily with the mild acids naturally present in many foods (citric acid in fruit, for example). If a can were lined with zinc, those acids would slowly dissolve zinc ions into the food, and zinc compounds in quantity are harmful, so the food would be unsafe.

Tin is less reactive and far more resistant to these food acids, so a thin tin coating stays put and keeps the iron of the can sealed away from both the food and the air. That is why “tin cans” are tin-plated steel. Note the contrast with iron buckets, which are galvanised (zinc-coated) precisely because zinc's higher reactivity gives sacrificial protection, but that very reactivity is unacceptable for food. So the correct answer is (c), zinc is more reactive than tin.

Answer: (c) Zinc is more reactive than tin and would react with and contaminate food; tin is safe.

Two independent tests, then judge them. A good answer here does two things: describe each test clearly and then weigh how reliable it is. The hammer probes malleability, a mechanical property, while the circuit probes electrical conductivity, an electrical property, so together they sample two very different metal-like behaviours, which makes a combined verdict stronger than either test alone.

The assessment part is what carries the higher marks. The conductivity test is quick and clear but fails for graphite, which conducts despite being a non-metal, and the malleability test can be ambiguous for samples that are small or hard. Because each test has a blind spot, the sensible conclusion is that the tests are useful as a first screen but should be used together, and borderline results should be confirmed by other properties such as lustre, sound on striking, or reaction with acid. Showing this balanced judgement, rather than claiming the tests are perfect, is exactly what an “assess” command word is looking for.

Answer: Hammer test (malleability) plus circuit test (conductivity) identify most metals; both have exceptions (e.g. conducting graphite), so use them together.

The “both ways” idea. The word amphoteric means “both”, so an amphoteric oxide is one that shows both acidic and basic character, and the cleanest way to prove it is with a pair of reactions. Reacting with an acid shows its basic side; reacting with an alkali shows its acidic side. Aluminium oxide and zinc oxide are the textbook pair, both belonging to metals that sit in a borderline region of the periodic table where the oxide is not strongly basic.

It helps to contrast the three oxide types: sodium oxide is purely basic (it only reacts with acids), sulphur dioxide is purely acidic (it only reacts with bases), and Al₂O₃ and ZnO sit in between, reacting with both. When you write the equations, the acid reaction gives the normal salt (such as AlCl₃), while the base reaction gives a special salt called an aluminate or zincate (such as NaAlO₂). Knowing both products, not just the definition, is what turns a two-mark answer into a full-mark one.

Answer: Amphoteric oxides react with both acids and bases; examples are Al₂O₃ and ZnO.

Read the series around hydrogen. The whole answer hinges on where hydrogen sits in the reactivity series, so picture the series with hydrogen marked in the middle. Everything above hydrogen, such as potassium, sodium, calcium, magnesium, aluminium, zinc and iron, is more reactive than hydrogen and can therefore kick it out of a dilute acid as H₂ gas, which you confirm with the pop test.

Everything below hydrogen, namely copper, silver, gold and platinum, is less reactive than hydrogen and so cannot displace it; that is why a copper coin dropped in dilute hydrochloric acid gives no bubbles. This single dividing line explains a lot of chapter facts at once: it is why reactive metals fizz in acids while noble metals do not, and why noble metals are used for jewellery and coins. So to answer the question safely, pick any two metals from above hydrogen (zinc, magnesium) and any two from below it (copper, silver).

Answer: Displace hydrogen: zinc, magnesium (above H). Do not: copper, silver (below H).

Follow the metal's journey. Electrolytic refining makes sense if you trace where the metal goes during electrolysis. The impure metal is made the anode, so when current flows it dissolves: its atoms lose electrons and enter the electrolyte as positive ions. Those same ions are pulled to the cathode, which is a thin sheet of pure metal, where they gain electrons and deposit as pure metal, so the cathode slowly grows.

The electrolyte must contain ions of the same metal, which is why a soluble salt of M is used; for copper the standard choice is acidified copper sulphate solution. The clever part is what happens to the impurities: more reactive impurities stay dissolved in the solution, while less reactive ones such as gold and silver simply fall to the bottom as anode mud, which is collected and sold. So the rule to remember is impure metal at the anode, pure metal at the cathode, and a salt solution of the metal as the electrolyte.

Answer: Anode: impure M; cathode: pure M; electrolyte: a soluble salt solution of M.

Non-metal oxide = acidic. The backbone of this question is that non-metals form acidic oxides, so identify the gas first and the litmus behaviour follows. Burning sulphur gives sulphur dioxide, a colourless gas with a sharp, choking smell, and being a non-metal oxide it is acidic. The dry versus moist contrast is the real test of understanding: with no water, the gas cannot form an acid, so dry litmus is unaffected, but on moist litmus the SO₂ dissolves to sulphurous acid and turns blue litmus red.

This same pattern is why non-metal oxides such as CO₂, SO₂ and NO₂ make rainwater acidic and contribute to acid rain. Contrast it with a metal oxide such as magnesium oxide, which would turn moist red litmus blue because metal oxides are basic. So from one observation, that the gas reddens moist blue litmus, you can correctly conclude the burning element is a non-metal. The balanced equation S + O₂ → SO₂ is already balanced as written, with one sulphur and two oxygen atoms on each side.

Answer: No change on dry litmus; reddens moist blue litmus (acidic); S + O₂ → SO₂.

Attack the two conditions for rust. Since rusting needs oxygen and water acting together, every prevention method works by removing at least one of them, so frame your two ways around that. The first family is barrier methods: paint, grease, oil or a coating of a less reactive metal physically seal the surface so neither air nor moisture can reach the iron. These are cheap and common, but they stop working the moment the coating is broken.

The second, smarter family is sacrificial protection, of which galvanising is the everyday example. Coating iron with the more reactive metal zinc means that even a scratched surface stays safe, because the zinc corrodes in preference to the iron; the same idea is used by bolting blocks of magnesium or zinc to ships and pipelines. A third route is alloying, turning iron into stainless steel by adding chromium and nickel so the metal itself resists rust. Naming one barrier method and one sacrificial method gives two genuinely different answers.

Answer: Galvanising (sacrificial zinc coat) and painting/oiling/greasing (barrier); both keep oxygen and moisture off the iron.

Metal vs non-metal oxides. The cleanest way to remember oxide character is to split it by metal versus non-metal, so anchor this answer to that contrast. Metal oxides are basic (and a couple, like Al₂O₃ and ZnO, are amphoteric), while non-metal oxides are mostly acidic, because non-metals tend to gain or share electrons and their oxides dissolve in water to give acids.

So when sulphur or carbon burns, the oxide formed reddens moist litmus and is acidic. The detail that earns the extra mark is remembering the small set of neutral non-metal oxides, water, carbon monoxide and nitrous oxide, which do not show acidic or basic behaviour. This acidic character of non-metal oxides is the chemistry behind acid rain, where SO₂ and oxides of nitrogen dissolve in rainwater. Framing the answer as “non-metals give acidic oxides, with a few neutral exceptions” captures the full picture.

Answer: Acidic oxides mainly (e.g. CO₂, SO₂), with a few neutral ones (e.g. CO, H₂O, N₂O).

One idea, four answers. All four parts come from two big themes, reactivity and ease of reduction, so apply these ideas rather than learning four separate facts. Reactivity explains (a), (b) and (c): the noble metals are at the bottom of the series, so they neither corrode nor lose lustre, making them perfect for jewellery; the alkali metals are at the very top, so they must be sealed under oil to keep air and water away; and aluminium, though reactive, is saved by a self-formed oxide skin that is thin, tough and non-porous, so the metal inside is shielded.

Part (d) is about the metallurgy that follows. Reduction of an oxide is far easier and cheaper than reduction of a carbonate or sulphide, so the first job in extraction is to turn the ore into the oxide: carbonate ores are calcined (heated in limited air) and sulphide ores are roasted (heated strongly in air). After that, a reducing agent such as carbon, or electrolysis for very reactive metals, finishes the job. Presenting the four answers under these two linked ideas shows the examiner you understand the chapter as a connected whole rather than as isolated facts.

Answer: (a) Unreactive and lustrous; (b) too reactive, so oil keeps air/water out; (c) protective Al₂O₃ layer; (d) oxides are easier to reduce than carbonates/sulphides.

Acid dissolves a basic coat. This is a neat real-life neutralisation question, so explain it as acid meeting base. When copper is exposed to moist air it slowly forms a dull coating of copper oxide and basic copper carbonate; both are basic in nature. To clean the vessel you need something that reacts with and dissolves this basic layer, and a mild acid does exactly that.

Lemon juice contains citric acid and tamarind contains tartaric acid, so rubbing the vessel with either provides acid that reacts with the basic tarnish, forming a soluble copper salt that wipes away and revealing bright copper beneath. The key insight is that the acid attacks the tarnish, not the underlying copper metal, which is why a brief clean restores shine without eating into the vessel. The same chemistry is why we descale kettles with vinegar or citric acid: a mild acid removes a basic or carbonate deposit by neutralisation.

Answer: The mild acids in lemon/tamarind juice dissolve the basic copper tarnish by neutralisation, restoring the shine.

Sort everything by electrons lost or gained. The single idea that organises every chemical difference between metals and non-metals is electron behaviour, so build the answer around it. Metals are electropositive: they readily lose electrons to form positive ions, and this is why they form basic oxides, why reactive ones displace hydrogen from acids, and why their chlorides are ionic. Non-metals are electronegative: they tend to gain electrons to form negative ions, which is why they form acidic oxides, do not displace hydrogen from acids, and form covalent compounds.

So instead of memorising a long list, remember the cause (electron loss versus gain) and derive each property from it. This also explains the exceptions neatly: amphoteric oxides such as Al₂O₃ sit at the metal-non-metal borderline, and graphite conducts electricity though it is a non-metal. The question asks specifically for chemical properties, so keep to oxide nature, reaction with water and acids, and ion formation, and present them in a clean two-column table for the clearest answer.

Answer: Metals lose electrons (positive ions, basic oxides, react with acids/water); non-metals gain electrons (negative ions, acidic/neutral oxides, no reaction with acids).

Only one solution dissolves gold. The detective work here turns on a single fact: gold is so unreactive that no ordinary acid dissolves it, and the one common reagent that does is aqua regia, so that has to be the answer. Aqua regia is a 3:1 mixture of concentrated hydrochloric and nitric acids, and it works because the nitric acid oxidises the gold while the chloride ions from the hydrochloric acid lock up the gold as a soluble chloride complex, the two acting together where neither could alone.

The con relies on dissolving just the outer skin of gold. Stripping that thin layer exposes fresh, bright gold underneath, so the bangles look newly polished, but the dissolved gold leaves with the man in the solution, and the ornaments are measurably lighter. So the chemistry both restores the shine and explains the lost weight, which is exactly the clue the question hands you. The takeaway worth remembering is that aqua regia is the standard answer whenever a problem mentions dissolving gold or platinum.

Answer: Aqua regia (3 conc. HCl : 1 conc. HNO₃); it dissolved the outer gold layer, so shine returned but weight fell.

Compare reactivity with steam. The choice of copper over steel comes down to how each metal behaves with hot water and steam, so compare them directly. Iron lies above copper in the reactivity series, and crucially iron reacts with steam to form iron oxide and hydrogen gas; in a hot water tank that constant exposure to hot water and steam would slowly corrode a steel tank and weaken it. Copper, being below iron and below hydrogen in the series, does not react with cold water, hot water or steam, so a copper tank stays sound for years.

There is a useful bonus reason: copper is an excellent conductor of heat, which helps a hot water tank heat and hold water efficiently, and copper also resists ordinary corrosion well. Steel would need extra protection such as galvanising or enamelling to survive the same conditions. So the full answer is that copper is chosen because it does not react with hot water or steam and conducts heat well, while steel would react with steam and corrode.

Answer: Copper does not react with hot water or steam and conducts heat well; iron/steel reacts with steam (3Fe + 4H₂O → Fe₃O₄ + 4H₂), so copper is used.