Metals and Non-metals

1. Classifying elements: where this chapter sits

In Class 9 you met many elements and learnt they split into two big families — metals and non-metals. This chapter answers three questions: how do their properties differ (physical and chemical), why do they behave that way (electron transfer), and how do we get metals out of the ground and protect them (extraction and corrosion).

There are far fewer non-metals than metals. Common non-metals: carbon, sulphur, iodine, oxygen, hydrogen, nitrogen, chlorine. Non-metals are usually solids or gases — the only liquid non-metal is bromine.

2. Physical properties of metals (3.1.1)

From Activities 3.1-3.6, metals share a set of physical properties:

• Metallic lustre — pure metals have a shining surface (Activity 3.1: rubbing with sandpaper restores the shine).
• Hardness — generally hard, though hardness varies (sodium and other alkali metals are so soft they can be cut with a knife — Activity 3.2).
• Malleability — can be beaten into thin sheets (Activity 3.3). Gold and silver are the most malleable metals.
• Ductility — can be drawn into thin wires (Activity 3.4). Gold is the most ductile — about 2 km of wire can be drawn from just 1 g of gold.
• Conduction of heat — good conductors with high melting points (Activity 3.5). Silver and copper are the best heat conductors; lead and mercury are poor conductors.
• Conduction of electricity — good conductors (Activity 3.6), which is why house wiring uses metal coated with PVC/rubber insulation.
• Sonorous — produce a ringing sound when struck (so school bells are made of metal).
• State — all metals are solids at room temperature except mercury (liquid). Gallium and caesium have such low melting points they melt in your palm.

3. Physical properties of non-metals — and the exceptions (3.1.2)

Non-metals are broadly the opposite: not lustrous, not sonorous, neither malleable nor ductile (they are brittle and shatter), and poor conductors of heat and electricity. But the chapter stresses you cannot classify on physical properties alone — there are striking exceptions:

• Iodine is a non-metal but is lustrous.
• Carbon shows allotropy (different forms): diamond is the hardest natural substance with a very high melting/boiling point, while graphite is a conductor of electricity (the one non-metal that conducts).
• Alkali metals (Li, Na, K) are soft, low density, low melting point — unusual for metals.
• Mercury is a liquid metal; bromine is a liquid non-metal.

Because of these exceptions, the clean dividing line comes from chemical properties, studied next.

4. Chemical properties — metals react with oxygen (3.2.1)

Almost all metals combine with oxygen to give a metal oxide:

Most metal oxides are basic. Soluble oxides dissolve in water to give alkalis:

Amphoteric oxides react with both acids and bases to give salt and water. Examples: aluminium oxide and zinc oxide.

Different metals react at different rates (Activity 3.9): Na and K react so fast they catch fire — so they are stored under kerosene. Mg, Al, Zn, Pb form a thin protective oxide layer that stops further oxidation. Iron does not burn but iron filings burn vigorously when sprinkled in a flame. Copper does not burn but coats with black CuO. Silver and gold do not react with oxygen even at high temperature.

5. Metals react with water (3.2.2)

The general reactions:

But not all metals react, and the violence depends on reactivity:

• K and Na react violently with cold water — so exothermic the hydrogen catches fire.
  2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy
  2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy
• Calcium reacts less violently (hydrogen does not catch fire) and floats because H₂ bubbles stick to it.
  Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)
• Magnesium does not react with cold water; it reacts with hot water to form Mg(OH)₂ and H₂, and also floats.
• Al, Fe, Zn react only with steam, giving the oxide + hydrogen:
  2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)
  3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
• Pb, Cu, Ag, Au do not react with water at all.

6. Metals react with acids (3.2.3)

Metals react with dilute acids to give a salt and hydrogen gas:

From Activity 3.11, the rate of bubbling (and the temperature rise) ranks reactivity: Mg > Al > Zn > Fe. Copper does not react with dilute HCl — no bubbles, no temperature change.

Nitric acid is special: usually no hydrogen is evolved with HNO₃ because it is a strong oxidising agent — it oxidises the H₂ to water and is itself reduced to nitrogen oxides (N₂O, NO, NO₂). The exceptions are Mg and Mn, which release H₂ with very dilute HNO₃.

7. Metals react with salt solutions — displacement (3.2.4)

A more reactive metal displaces a less reactive metal from its salt solution:

In Activity 3.12, an iron nail in copper sulphate turns the blue solution pale green and gets a brown copper coat — so iron is more reactive than copper. A copper wire in iron sulphate shows no change.

8. The reactivity (activity) series (3.2.5)

Combining all the oxygen, water, acid and displacement experiments gives the reactivity series — metals in order of decreasing reactivity (most reactive at top):

Memory aid: "Please Stop Calling Me A Zebra; I Like People Having Cute Small Goats" — but the standard symbols K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au are what you write in answers. Note hydrogen [H] is placed in the series even though it is not a metal — metals above H displace hydrogen from dilute acids; metals below H (Cu, Ag, Au) do not.

9. Why metals and non-metals react — ionic bonding (3.3)

Noble gases have a completely filled valence shell and are unreactive. Other atoms react to attain a full outer shell (octet). Metals do this by losing their few valence electrons; non-metals by gaining electrons.

Sodium chloride (NaCl): sodium (2,8,1) loses 1 electron to become Na⁺ (2,8); chlorine (2,8,7) gains it to become Cl⁻ (2,8,8):

Magnesium chloride (MgCl₂): Mg (2,8,2) loses 2 electrons to give Mg²⁺; two Cl atoms each take one:

Compounds formed by such transfer of electrons from a metal to a non-metal are ionic (electrovalent) compounds. They do not exist as molecules but as huge aggregates of oppositely charged ions held by strong electrostatic force.

10. Properties of ionic compounds (3.3.1)

From Activity 3.13 (testing NaCl, KI, BaCl₂):

• Physical nature: hard, crystalline solids; brittle (break on applying pressure) because of strong inter-ionic forces.
• High melting and boiling points: a lot of energy is needed to break the strong ionic attraction (e.g. NaCl melts at 1074 K, CaO at 2850 K).
• Solubility: generally soluble in water, insoluble in solvents like kerosene and petrol.
• Conduction of electricity: conduct in molten state or in aqueous solution (ions are free to move), but not in the solid state (ions are locked in a rigid lattice).

11. Occurrence and the start of extraction (3.4, 3.4.1)

The earth's crust is the main source of metals; seawater holds soluble salts (NaCl, MgCl₂). Key terms:

• Mineral: any naturally occurring element/compound in the crust.
• Ore: a mineral with enough metal to be profitably extracted.
• Gangue: impurities (soil, sand) mixed with the ore.

Reactivity decides how a metal occurs and is extracted:

• Low reactivity (Au, Ag, Pt, Cu) — often found free (native).
• Medium reactivity (Zn, Fe, Pb, Cu) — as oxides, sulphides or carbonates.
• High reactivity (K, Na, Ca, Mg, Al) — never free; very stable compounds; need electrolysis.

Metallurgy = extracting a metal from its ore and refining it. The steps: Ore → concentration (enrichment) → conversion to oxide → reduction to metal → refining. Enrichment (3.4.2) removes gangue using differences in physical/chemical properties of ore and gangue.

12. Extracting metals low in the series (3.4.3)

These are so unreactive that their oxides decompose to the metal on heating alone. Mercury from cinnabar (HgS):

Copper from Cu₂S by heating in air:

13. Extracting metals in the middle — roasting, calcination, reduction (3.4.4)

Zn, Fe, Pb, Cu occur as sulphides or carbonates. Since it is easier to reduce an oxide, first convert the ore to oxide:

• Roasting — heat the sulphide ore strongly in excess air → oxide.
  2ZnS(s) + 3O₂(g)  (Heat)→ 2ZnO(s) + 2SO₂(g)
• Calcination — heat the carbonate ore strongly in limited air → oxide.
  ZnCO₃(s)  (Heat)→ ZnO(s) + CO₂(g)

The oxide is then reduced using a reducing agent such as carbon (coke):

Displacement reduction (thermit): highly reactive metals (Na, Ca, Al) can act as reducing agents. Reducing manganese dioxide with aluminium:

The thermit reaction (iron(III) oxide + aluminium) is so exothermic the iron forms molten — used to weld railway tracks and cracked machine parts:

14. Extracting metals at the top — electrolytic reduction (3.4.5)

K, Na, Ca, Mg, Al are too reactive for carbon to reduce (they hold oxygen more strongly than carbon does). They are obtained by electrolysis of their molten chlorides/oxides. For sodium chloride:

Metal deposits at the cathode; chlorine is liberated at the anode. Aluminium is similarly obtained by electrolytic reduction of aluminium oxide.

15. Refining of metals — electrolytic refining (3.4.6)

Reduced metals are impure. The commonest purification is electrolytic refining (used for Cu, Zn, Sn, Ni, Ag, Au):

• Anode (+): the impure metal.
• Cathode (−): a thin strip of pure metal.
• Electrolyte: a solution of a salt of the same metal (e.g. acidified copper sulphate for copper).

On passing current, pure metal from the anode dissolves into the electrolyte and an equal amount of pure metal deposits on the cathode. Soluble impurities go into solution; insoluble impurities settle below the anode as anode mud.

16. Corrosion and its prevention (3.5)

Corrosion is the slow eating-up of a metal surface by the action of air and moisture:

• Silver turns black — reacts with sulphur in air to form silver sulphide.
• Copper gains a green coat — reacts with moist CO₂ to form basic copper carbonate.
• Iron forms a brown flaky coating called rust.

Activity 3.14 proves iron rusts only when both air (oxygen) and water are present: nails rust in tube A (air + water), but not in tube B (boiled water + oil, no air) or tube C (dry air with calcium chloride).

Prevention (3.5.1): painting, oiling, greasing, galvanising (coating with zinc — protects even if the coat is scratched, because zinc is more reactive than iron and corrodes first), chrome plating, anodising, and alloying.

17. Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal, made by melting the main metal and dissolving the others in fixed proportions, then cooling. Alloying changes properties to suit our needs.

• Steel = iron + a little carbon (~0.05%): hard and strong (pure iron is soft and stretches).
• Stainless steel = iron + nickel + chromium: hard and rust-free.
• Brass = copper + zinc; Bronze = copper + tin — both poor conductors compared with copper.
• Solder = lead + tin: low melting point, used to weld electrical wires.
• Amalgam = any alloy containing mercury.
• Gold: 24-carat gold is pure but very soft; jewellery uses 22-carat gold (alloyed with copper or silver) for hardness.

An alloy has lower electrical conductivity and a lower melting point than its pure metals. (The rust-resistant iron pillar near Qutub Minar, Delhi — over 1600 years old, 8 m tall, 6 tonnes — shows ancient Indian metallurgy.)

18. Non-metals: chemical behaviour (recap)

• Non-metals gain electrons to form negative ions when reacting with metals.
• They form oxides that are acidic or neutral (Activity 3.8: burning sulphur gives an acidic oxide that turns blue litmus red; burning magnesium gives a basic oxide).
• They do not displace hydrogen from dilute acids; they react with hydrogen to form hydrides.

19. In-text QUESTIONS — fully answered

Page 40, Q1. Example of a metal which (i) liquid at room temperature: mercury; (ii) easily cut with a knife: sodium (or potassium); (iii) best conductor of heat: silver; (iv) poor conductor of heat: lead (or mercury).

Page 40, Q2. Malleable = can be beaten into thin sheets. Ductile = can be drawn into thin wires.

Page 46, Q1. Sodium is kept in kerosene because it is so reactive it catches fire on contact with air/moisture; kerosene cuts off air and prevents accidental fires.

Page 46, Q2. (i) Iron with steam: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g). (ii) Calcium with water: Ca + 2H₂O → Ca(OH)₂ + H₂; Potassium with water: 2K + 2H₂O → 2KOH + H₂ + heat.

Page 46, Q3 (A, B, C, D table). (i) Most reactive = B (it displaces Fe and others). (ii) If B is added to copper(II) sulphate, a displacement occurs — copper is deposited and the blue colour fades. (iii) Decreasing reactivity: B > A > C > D.

Page 46, Q4. Hydrogen gas is produced when a reactive metal reacts with dilute acid. Iron with dilute H₂SO₄: Fe + H₂SO₄ → FeSO₄ + H₂.

Page 46, Q5. Zinc is more reactive than iron, so it displaces iron: Zn + FeSO₄ → ZnSO₄ + Fe. The pale-green FeSO₄ solution becomes colourless and a grey deposit of iron forms.

Page 49, Q1. (i) Electron-dot structures — Na: one dot in outer shell; oxygen: six dots (needs 2 more); magnesium: two dots. (ii) Na₂O: two Na each give 1 electron to O → 2Na⁺ + O²⁻. MgO: Mg gives 2 electrons to O → Mg²⁺ + O²⁻. (iii) Ions present: Na₂O has Na⁺ and O²⁻; MgO has Mg²⁺ and O²⁻.

Page 49, Q2. Ionic compounds have high melting points because a large amount of energy is needed to overcome the strong electrostatic force between the oppositely charged ions.

Page 53, Q1. Mineral = naturally occurring element/compound in the crust. Ore = mineral with enough metal to extract profitably. Gangue = impurities (sand, soil) in the ore.

Page 53, Q2. Two metals found free (native): gold and silver (also platinum, copper).

Page 53, Q3. A metal is obtained from its oxide by reduction (using carbon, a more reactive metal, or electrolysis).

Page 55, Q1 (Zn/Mg/Cu oxides). Displacement occurs when a more reactive metal meets a less reactive metal's oxide. Reactivity Mg > Zn > Cu, so: zinc oxide is displaced by magnesium; magnesium oxide is displaced by none (Mg most reactive); copper oxide is displaced by both zinc and magnesium.

Page 55, Q2. Metals that do not corrode easily: gold, silver, platinum (least reactive).

Page 55, Q3. Alloys = homogeneous mixtures of two or more metals, or a metal and a non-metal.

20. Exercise questions — fully answered

Q1. Which pairs give displacement reactions? Answer: (d) AgNO₃ solution and copper metal — copper is more reactive than silver: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. (In (b) Al cannot displace the more reactive Mg; (a) and (c) involve less reactive metals.)

Q2. Method to prevent an iron frying pan from rusting: (d) All of the above (grease, paint, zinc coating all work).

Q3. Element giving a high-melting, water-soluble oxide: (a) calcium (forms basic CaO which dissolves to give Ca(OH)₂).

Q4. Food cans are coated with tin not zinc because: (c) zinc is more reactive than tin — zinc would react with the food.

Q5. (a) Set up a circuit with battery, bulb and wires; place the sample between the terminals — if the bulb glows the sample conducts (metal); if not, non-metal. (b) Limited use: graphite (a non-metal) also conducts, so the test alone is not fully reliable, though it works for most cases.

Q6. Amphoteric oxides react with both acids and bases to give salt and water. Examples: Al₂O₃ and ZnO.

Q7. Two metals that displace hydrogen from dilute acids (above H): magnesium, zinc (also Fe, Al). Two that do not (below H): copper, silver (also gold, mercury).

Q8. Electrolytic refining of metal M: anode = impure M, cathode = thin strip of pure M, electrolyte = a soluble salt solution of M.

Q9. Pratyush burns sulphur. (a) The gas is SO₂ — acidic: (i) no effect on dry litmus; (ii) turns moist blue litmus red. (b) S + O₂ → SO₂; SO₂ + H₂O → H₂SO₃.

Q10. Two ways to prevent rusting: painting/oiling and galvanising (zinc coating).

Q11. Non-metals combine with oxygen to form acidic or neutral oxides.

Q12. (a) Platinum, gold, silver are used for jewellery because they are lustrous, very unreactive (do not corrode/tarnish) and malleable. (b) Na, K, Li are stored under oil because they react violently with air and moisture, even catching fire. (c) Aluminium, though reactive, forms a thin tough oxide layer that protects it from further corrosion, so it is safe for utensils. (d) Carbonate and sulphide ores are first converted to oxides because oxides are easier to reduce to the metal than sulphides/carbonates.

Q13. Tarnished copper vessels are cleaned with lemon/tamarind because the tarnish is basic copper carbonate (or copper oxide); the citric/tartaric acid in these sour juices dissolves it, restoring the shine.

Q14. Metal vs non-metal (chemical): metals lose electrons → form positive ions, give basic oxides, displace H₂ from acids, are reducing agents. Non-metals gain electrons → form negative ions, give acidic/neutral oxides, do not displace H₂, are oxidising agents.

Q15. The fake goldsmith dipped the bangles in aqua regia (conc. HCl + conc. HNO₃, 3:1), which dissolves gold; that is why the bangles shone but lost weight.

Q16. Copper is used for hot-water tanks (not steel/iron) because copper does not react with hot water or steam, while iron reacts with steam (3Fe + 4H₂O → Fe₃O₄ + 4H₂) and would corrode.

21. Common mistakes to avoid

• Confusing roasting (sulphide, excess air) with calcination (carbonate, limited air).
• Writing that metals react with nitric acid to give H₂ — usually they do not (HNO₃ is an oxidiser); only Mg/Mn with very dilute HNO₃.
• Saying ionic compounds conduct in the solid state — they conduct only when molten or dissolved.
• Forgetting that graphite (non-metal) conducts electricity and iodine (non-metal) is lustrous.
• Mixing up cathode/anode in electrolytic refining — impure metal = anode, pure metal = cathode.
• Forgetting to balance — e.g. 4Al + 3O₂ → 2Al₂O₃, not Al + O₂.

22. Quick revision checklist

• Physical: lustre, malleable, ductile, sonorous, conduct heat/electricity, solid (except Hg).
• Metal + O₂ → basic oxide; Al₂O₃ & ZnO are amphoteric.
• Metal + acid → salt + H₂; reactivity Mg > Al > Zn > Fe; Cu unreactive with dilute HCl.
• Reactivity series: K Na Ca Mg Al Zn Fe Pb [H] Cu Hg Ag Au.
• Ionic compound = metal loses, non-metal gains electrons; high m.p., soluble, conduct when molten/aqueous.
• Extraction: roasting (sulphide), calcination (carbonate), reduction (C / thermit / electrolysis), electrolytic refining.
• Corrosion prevented by painting, oiling, galvanising, anodising, alloying.