Class 10 Science Chapter 3 Notes: Metals and Non-metals

How to tell metals and non-metals apart by their physical properties (lustre, hardness, malleability, ductility), and why a handful of exceptions mean this test alone is never fully reliable
How metals react with oxygen, water, acids and other metals’ salt solutions, and how these reactions build the reactivity series
How an ionic bond forms, and why ionic compounds have the properties they do
Every step from ore to pure metal: enrichment, roasting and calcination, reduction, and refining
The conditions under which corrosion happens, how it is prevented, and how alloying changes a metal’s properties

metallic lustremalleabilityductilityamphoteric oxides
reactivity seriesdisplacement reactionionic compoundsminerals and ores
roasting and calcinationthermit reactionelectrolytic refining
corrosiongalvanisationalloys

1 Telling Metals and Non-metals Apart by Physical Properties

Look around your own home: a spoon, a nail, a wire, a steel vessel: all of these are shiny and strong. Coal, sulphur and iodine look nothing like them. This difference is where the split of elements into two big groups, metals and non-metals, begins.

1.1 Physical properties of metals

Activities 3.1 to 3.6

Samples of iron, copper, aluminium, magnesium and others are cleaned, cut, hammered, drawn into wire, heated and connected in a circuit. Every test brings out the same kind of behaviour in metals.

Metallic lustre: a metal’s surface is shiny in its pure state
Hardness: metals are generally hard, though hardness varies a good deal from one metal to another
Malleability: can be hammered into thin sheets: gold and silver are the most malleable metals
Ductility: can be drawn into thin wire: gold is the most ductile, about 2 km of wire can be drawn from just 1 gram of it
Thermal conductivity: silver and copper are the best conductors of heat; lead and mercury conduct comparatively poorly
Electrical conductivity: connecting a metal into a circuit lights the bulb
Sonority: gives out a ringing sound when struck on a hard surface, which is why a school bell is made of metal

Did you know?

It is malleability and ductility together that let metals be shaped into exactly what we need: vessels, wires, sheets and jewellery all rely on these two properties.

1.2 Physical properties of non-metals

Non-metals are far fewer in number than metals: carbon, sulphur, iodine, oxygen and hydrogen are the main examples. Most non-metals are solids or gases; only bromine is found in the liquid state.

Physically, non-metals are almost the opposite of metals: they are not malleable and not ductile (most are brittle solids or gases), and they are poor conductors of heat and electricity: the one well-known exception is graphite, a non-metal that does conduct electricity.

1.3 The limits of physical properties: classification cannot rest on these alone

If metals and non-metals were classified purely by the properties above, four cases would go wrong:

The mercury exception

Every metal is solid at room temperature except mercury, which is liquid. Gallium and caesium melt at such low temperatures that they start melting if simply held in the palm of the hand.

The iodine exception

Iodine is a non-metal, yet it looks lustrous.

Allotropes of carbon

Diamond and graphite are both allotropes of carbon, a non-metal. Diamond is the hardest known natural substance, while graphite conducts electricity.

The alkali metals

Lithium, sodium and potassium are so soft they can be cut with a knife. Both their density and melting point are low.

Exam Tip

For “Why can metals and non-metals not be classified on the basis of physical properties alone?”, always give at least two of the four exceptions above as your examples.

Activity 3.8

Magnesium is burnt and its ash is dissolved in water; the resulting solution turns red litmus paper blue (it is basic). Sulphur is burnt separately and the fumes given off are collected and dissolved in water; the resulting solution turns blue litmus paper red (it is acidic).

Learn by heartDefinition 1

Most oxides of non-metals are acidic and most oxides of metals are basic. Elements can be divided into metals and non-metals far more reliably on the basis of their chemical properties.

2 Chemical Properties of Metals

2.1 Reaction with oxygen: amphoteric oxides

Activity 3.9

Different metals are burnt in a flame and observed: which one catches fire easily, what colour the flame turns, and what the surface looks like after burning.

Almost all metals combine with oxygen to form the corresponding metal oxide:

General rule

metal + oxygen → metal oxide

2\,\ce{Cu + O2 -> 2CuO}

(black copper(II) oxide, where the “(II)” shows copper carries a +2 charge here) ·

4\,\ce{Al + 3O2 -> 2Al2O3}

Learn by heartDefinition 2

An amphoteric oxide is a metal oxide that reacts with both acids and bases, giving a salt and water in each case. The book names two examples: aluminium oxide (Al₂O₃) and zinc oxide (ZnO) (it says “some metal oxides, such as” these two, so they are examples, not a complete list).

\ce{Al2O3 + 6HCl -> 2AlCl3 + 3H2O}

(with an acid)

\ce{Al2O3 + 2NaOH -> 2NaAlO2 + H2O}

(with a base, forming sodium aluminate)

Some metal oxides also dissolve in water to form alkalis (bases that dissolve in water):

\ce{Na2O(s) + H2O(l) -> 2NaOH(aq)}

\ce{K2O(s) + H2O(l) -> 2KOH(aq)}

Not every metal combines with oxygen at the same rate:

Potassium and sodium react so fast that they catch fire the moment they are left in the open, so they are stored under kerosene oil
Magnesium, aluminium, zinc and lead form a thin oxide layer on the surface, which protects the metal from further oxidation
Iron does not burn directly when heated, but iron filings sprinkled into a flame burn vigorously
Copper does not burn, but a black layer of copper(II) oxide forms on the heated metal
Silver and gold do not react with oxygen even at very high temperatures

The most reactive metals hold onto oxygen so strongly that ordinary heating cannot pull them apart from their compounds at all. Getting those metals out, and purifying some others, needs a different method altogether: electrolysis, defined here so it is ready before Section 4 uses it for extraction and refining.

Learn by heartDefinition 3

Electrolysis is the process of passing an electric current through a molten (melted) or dissolved compound to break it apart into its elements. The two rods or plates that carry the current into the liquid are called electrodes, and the liquid itself is the electrolyte. The cathode is the negatively charged electrode; the anode is the positively charged electrode.

Did you know?

Anodising: the process of building up a thick protective oxide layer on aluminium. A clean aluminium article is made the anode and electrolysed in dilute sulphuric acid. The resulting thick oxide layer can be dyed to make it attractive.

2.2 Reaction with water

Activity 3.10

Metal samples are tested first with cold water, then hot water, and finally by passing steam directly over the heated metal.

General rule

metal + water → metal oxide + hydrogen

If the metal oxide formed is soluble in water, it dissolves further to give the metal hydroxide: metal oxide + water → metal hydroxide

Metal	With cold water	With hot water	With steam
Potassium, sodium	Violent, exothermic (releases heat), H₂ catches fire immediately	×	×
Calcium	Slower reaction, starts floating	×	×
Magnesium	No reaction	Forms Mg(OH)₂ + H₂, floats	×
Aluminium, iron, zinc	No reaction	No reaction	Forms metal oxide + H₂
Lead, copper, silver, gold	Do not react with water at all

× = not tested at that stage, because the metal had already reacted at a milder stage (e.g. potassium already reacts violently with cold water, so there is no need to test it with hot water or steam)

2\,\ce{K(s) + 2H2O(l) -> 2KOH(aq) + H2(g)}

+ heat energy

2\,\ce{Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)}

+ heat energy

\ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)}

2\,\ce{Al(s) + 3H2O(g) -> Al2O3(s) + 3H2(g)}

3\,\ce{Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g)}

Common Mistake

Calcium and magnesium float on water, not because they are light but because bubbles of the hydrogen gas produced stick to the metal’s surface.

2.3 Reaction with dilute acids

Activity 3.11

All metals except sodium and potassium are dropped into dilute hydrochloric acid, and the rate of bubbling and the change in temperature are noted.

General rule

metal + dilute acid → salt + hydrogen

The order of reactivity found experimentally: Mg > Al > Zn > Fe. Copper shows no bubbling and no change in temperature at all: copper does not react with dilute HCl.

Common Mistake

Nitric acid (HNO₃) reacting with a metal does not usually release hydrogen gas, because HNO₃ is a strong oxidising agent: it oxidises other substances by supplying them oxygen, and in doing so is itself reduced, meaning it loses oxygen. Here it oxidises the hydrogen formed into water, and is itself reduced to an oxide of nitrogen: you will meet oxidation and reduction properly defined in the extraction section below. Only magnesium and manganese, both reactive metals, give H₂ gas with very dilute HNO₃.

Did you know?

Aqua regia: a freshly prepared 3:1 mixture of concentrated hydrochloric acid and concentrated nitric acid. It can dissolve gold and platinum, even though neither acid can do so on its own.

2.4 Reaction with salt solutions of other metals: displacement

Activity 3.12

A copper wire is placed in an iron sulphate solution, and an iron nail in a copper sulphate solution. After 20 minutes, only the test tube with the iron nail shows a change: a brownish coating of copper deposits on the nail and the blue of the solution fades. The test tube with the copper wire shows no change at all.

Learn by heartDefinition 4

A more reactive metal displaces a less reactive metal from a solution or molten state of its compound: metal (A) + salt solution of (B) → salt solution of (A) + metal (B), where metal A is more reactive than metal B.

The result of Activity 3.12 shows that iron is more reactive than copper: iron displaced copper from its solution, but copper could not displace iron from its solution.

2.5 The reactivity series

Learn by heartDefinition 5

The reactivity (activity) series is a list of metals arranged in order of decreasing reactivity. It has been built up from displacement experiments like the one above, repeated with many pairs of metals.

Metal	Symbol	Reactivity
Potassium	K	Most reactive
Sodium	Na
Calcium	Ca
Magnesium	Mg
Aluminium	Al
Zinc	Zn	Reactivity decreases
Iron	Fe
Lead	Pb
(Hydrogen, for comparison)	[H]	×
Copper	Cu
Mercury	Hg
Silver	Ag
Gold	Au	Least reactive

× = not applicable: hydrogen is not a metal and is placed here only for comparison, to show which metals can displace it from a dilute acid and which cannot. Only potassium, zinc and gold carry a reactivity label here; the metals in between are simply read top to bottom, most to least reactive.

Exam Tip

Being able to write out the entire reactivity series (K to Au) from memory is essential: this same order decides how each metal is extracted, in the next section. Remember where [H] sits too: metals above it can displace hydrogen from a dilute acid, metals below it cannot.

2.6 Chemical properties of non-metals

Non-metals react in almost the opposite way to metals. A non-metal atom gains electrons instead of losing them, to become a negatively charged ion.

Learn by heartDefinition 6

Non-metals combine with oxygen to form acidic or neutral oxides (never basic). Sulphur dioxide (SO₂), formed when sulphur burns, is an acidic oxide. Non-metals do not displace hydrogen from dilute acids; instead, a non-metal can react directly with hydrogen gas to form a compound called a hydride (for example, hydrogen chloride, HCl).

3 How Metals and Non-metals React: Ionic Compounds

You studied in Class 9 that noble gases have a completely filled valence shell, which is why they hardly react at all. The reactivity of every other element can be understood as its tendency to complete that same valence shell: the electronic configurations below (Table 3.3) make this comparison directly.

Type of element	Element	K	L	M	N
Noble gases	Neon (Ne)	2	8	×	×
Noble gases	Argon (Ar)	2	8	8	×
Metals	Sodium (Na)	2	8	1	×
	Magnesium (Mg)	2	8	2	×
	Aluminium (Al)	2	8	3	×
Non-metals	Chlorine (Cl)	2	8	7	×
Non-metals	Oxygen (O)	2	6	×	×

× = this shell does not exist in the atom (the atom has fewer occupied shells than the table has columns)

A sodium atom’s outermost (M) shell holds just one electron. The moment it loses that electron, its L shell, already a stable and fully filled octet (eight electrons, the arrangement every atom “wants”), becomes the outermost shell. A chlorine atom’s outermost shell, on the other hand, holds 7 electrons, one short of a full octet. So the electron sodium gives up is exactly the one chlorine needs:

Figure 1 Formation of sodium chloride (top) and magnesium chloride (bottom) by transfer of electrons. Dots are the metal atom's valence electrons and crosses the chlorine atom's; the orange arrow shows an electron moving. Sodium loses one electron (2,8,1 → Na⁺, 2,8) and chlorine gains it (2,8,7 → Cl⁻, 2,8,8). Magnesium loses two electrons, one to each of two chlorine atoms (2,8,2 → Mg²⁺, 2,8). The oppositely charged ions are held together by strong electrostatic attraction, the pull between opposite electric charges. — **Figure 1** Formation of sodium chloride (top) and magnesium chloride (bottom) by transfer of electrons. Dots are the metal atom’s valence electrons and crosses the chlorine atom’s; the orange arrow shows an electron moving. Sodium loses one electron (2,8,1 → Na⁺, 2,8) and chlorine gains it (2,8,7 → Cl⁻, 2,8,8). Magnesium loses two electrons, one to each of two chlorine atoms (2,8,2 → Mg²⁺, 2,8). The oppositely charged ions are held together by strong electrostatic attraction, the pull between opposite electric charges.

\ce{Na -> Na+ + e-}

(2,8,1 → 2,8)

\ce{Cl + e- -> Cl-}

(2,8,7 → 2,8,8)

Common Mistake

Sodium chloride is not found as molecules: it is an aggregate (a large cluster) of oppositely charged ions.

Magnesium chloride forms the same way, except magnesium has to give up two electrons (so it needs two chlorine atoms):

\ce{Mg -> Mg^2+ + 2e-}

(2,8,2 → 2,8)

2\,\ce{(Cl + e- -> Cl-)}

(2,8,7 → 2,8,8 each)

Sodium oxide and magnesium oxide form in exactly the same way. Oxygen (2,6) needs two more electrons, so it takes one from each of two sodium atoms, or both from one magnesium atom (this is the NCERT in-text question on electron-dot structures):

Figure 2 Electron-dot structures of sodium (1 valence electron), oxygen (6) and magnesium (2), and the formation of sodium oxide and magnesium oxide. Two sodium atoms each give one electron to an oxygen atom, giving 2Na⁺ and O²⁻ (Na₂O); one magnesium atom gives both its electrons to an oxygen atom, giving Mg²⁺ and O²⁻ (MgO). Every ion now has a complete octet (2,8). The ions present are Na⁺ and O²⁻ in Na₂O, and Mg²⁺ and O²⁻ in MgO. — **Figure 2** Electron-dot structures of sodium (1 valence electron), oxygen (6) and magnesium (2), and the formation of sodium oxide and magnesium oxide. Two sodium atoms each give one electron to an oxygen atom, giving 2Na⁺ and O²⁻ (Na₂O); one magnesium atom gives both its electrons to an oxygen atom, giving Mg²⁺ and O²⁻ (MgO). Every ion now has a complete octet (2,8). The ions present are Na⁺ and O²⁻ in Na₂O, and Mg²⁺ and O²⁻ in MgO.

Learn by heartDefinition 7

Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.

3.1 Properties of ionic compounds

Activity 3.13

Salts such as sodium chloride and potassium iodide are heated in a flame, and attempts are made to dissolve them in water, petrol and kerosene. Their solutions are then tested for electrical conductivity using a simple circuit with a battery and a bulb.

Ionic compound	Melting point (K)	Boiling point (K)
NaCl	1074	1686
LiCl	887	1600
CaCl₂	1045	1900
CaO	2850	3120
MgCl₂	981	1685

CaO2850 K

NaCl1074 K

CaCl₂1045 K

MgCl₂981 K

LiCl887 K

Melting points of ionic compounds (Table 3.4 data): CaO has the highest melting point. This is because both Ca²⁺ and O²⁻ carry a double charge, so the pull between them is the strongest of all the compounds in the table.

These experiments bring out four properties of ionic compounds:

Physical nature

Solid, somewhat hard (because of strong electrostatic attraction), but generally brittle: they break into pieces the moment pressure is applied

Melting and boiling points

Very high, because a great deal of energy is needed to break the strong inter-ionic force of attraction

Solubility

Soluble in water, insoluble in solvents such as kerosene and petrol

Electrical conductivity

No conduction in the solid state (the ions cannot move), but they conduct in the aqueous solution and molten state (the ions are then free to move)

4 Occurrence and Extraction of Metals

The earth’s crust is the main source of metals. Sea water, too, holds some soluble salts such as sodium chloride and magnesium chloride. Metallurgy is the name for the whole process covered in this section: extracting a metal from its ore and then refining it for use.

Learn by heartDefinition 8

Mineral: an element or compound found naturally in the earth’s crust. Ore: a mineral from which a metal can be extracted profitably, because it contains enough of that metal.

4.1 Three classes of metals by reactivity

Low reactivity (around Cu, Ag, Au)

Found in the free state: gold, silver, platinum and copper. Copper and silver are also found as sulphide/oxide ores.

Medium reactivity (Zn, Fe, Pb)

Found in the earth’s crust mostly as oxides, sulphides or carbonates.

High reactivity (K, Na, Ca, Mg, Al)

So reactive that they are never found as free elements, always found combined in compounds.

4.2 Enrichment of ores

Mined ores come mixed with impurities such as soil and sand, called gangue. Before the metal is extracted, the gangue is removed by exploiting differences between the physical or chemical properties of the ore and the gangue.

4.3 Extraction of low-reactivity metals: heating alone

These metals are so unreactive that simply heating their ore in air, in two steps, is enough to obtain the metal. Cinnabar (HgS), the ore of mercury, is first heated in air to convert it into mercuric oxide (HgO); that oxide is then reduced to metallic mercury on further heating:

2\,\ce{HgS(s) + 3O2(g) ->[\Delta] 2HgO(s) + 2SO2(g)}

(cinnabar roasted to mercuric oxide)

2\,\ce{HgO(s) ->[\Delta] 2Hg(l) + O2(g)}

(mercuric oxide decomposes to give mercury)

Copper, found in nature as Cu₂S, follows the same two-step pattern: heating in air first gives copper(I) oxide, which then reacts further with the remaining sulphide to give the metal directly:

2\,\ce{Cu2S + 3O2(g) ->[\Delta] 2Cu2O(s) + 2SO2(g)}

2\,\ce{Cu2O + Cu2S ->[\Delta] 6Cu(s) + SO2(g)}

4.4 Extraction of medium-reactivity metals

These metals are usually found as sulphides or carbonates, which are first converted to oxides:

Roasting

Heating a sulphide ore strongly in excess air
$2\,\ce{ZnS(s) + 3O2(g) ->[\Delta] 2ZnO(s) + 2SO2(g)}$

Calcination

Heating a carbonate ore in limited air
$\ce{ZnCO3(s) ->[\Delta] ZnO(s) + CO2(g)}$

Learn by heartDefinition 9

Reduction is the removal of oxygen from a compound. A reducing agent is the substance that brings this about, by itself combining with the oxygen that is removed: carbon does exactly this when it turns a metal oxide back into the metal.

The metal oxide formed is then reduced with a reducing agent such as carbon:

\ce{ZnO(s) + C(s) -> Zn(s) + CO(g)}

Highly reactive metals (sodium, calcium, aluminium) are also used as reducing agents, because they can displace a less reactive metal from its compound. Here aluminium displaces manganese, another reactive metal, from manganese dioxide:

3\,\ce{MnO2(s) + 4Al(s) -> 3Mn(l) + 2Al2O3(s)}

+ heat

Learn by heartDefinition 10

The thermit reaction: a highly exothermic displacement reaction between iron(III) oxide and aluminium, in which iron is obtained in the molten state. Used for joining railway tracks and machine parts.

\ce{Fe2O3(s) + 2Al(s) -> 2Fe(l) + Al2O3(s)}

+ heat

4.5 Extraction of high-reactivity metals: electrolytic reduction

The metals at the very top of the reactivity series have a stronger affinity (chemical attraction) for oxygen than carbon does, so carbon cannot reduce them. They are obtained instead by electrolysis of their molten compounds:

At the cathode:

\ce{Na+ + e- -> Na}

At the anode:

2\,\ce{Cl- -> Cl2 + 2e-}

Sodium, magnesium and calcium are obtained this way from their molten chlorides; aluminium, from the electrolytic reduction of aluminium oxide.

4.6 Refining of metals: electrolytic refining

The metal obtained after reduction is not completely pure. The most common purification method is:

Figure 3 Electrolytic refining. The impure metal is the anode, joined to the positive terminal, and a thin strip of pure metal is the cathode, joined to the negative terminal; a solution of a salt of the metal is the electrolyte (acidified copper sulphate for copper). On passing current, metal from the anode dissolves into the electrolyte and an equal amount of pure metal is deposited on the cathode. Soluble impurities go into the solution; insoluble impurities settle at the bottom, below the anode, as anode mud. — **Figure 3** Electrolytic refining. The impure metal is the anode, joined to the positive terminal, and a thin strip of pure metal is the cathode, joined to the negative terminal; a solution of a salt of the metal is the electrolyte (acidified copper sulphate for copper). On passing current, metal from the anode dissolves into the electrolyte and an equal amount of pure metal is deposited on the cathode. Soluble impurities go into the solution; insoluble impurities settle at the bottom, below the anode, as anode mud.

The impure metal is made the anode, a thin strip of the pure metal is made the cathode, and a solution of the metal’s salt is used as the electrolyte. Soluble impurities pass into the solution. Many metals, including copper, zinc, tin, nickel, silver and gold, are refined this way.

5 Corrosion and Alloys

5.1 The conditions under which corrosion happens

Corrosion is the gradual eating away of a metal’s surface, caused by it reacting with substances around it such as air and moisture. Recall from Chapter 1: silver turns black in air (a layer of silver sulphide forms), copper develops a green coating (basic copper carbonate), and iron develops a brownish, flaky coating known as rust. All three are examples of corrosion.

Activity 3.14

Clean iron nails are placed in three test tubes, A, B and C. A holds plain water; B holds boiled distilled water (dissolved air removed) with a layer of oil on top; C holds dry air and anhydrous calcium chloride (which absorbs moisture). After a few days, only the nail in A has rusted.

Learn by heartDefinition 11

Rusting needs both air (oxygen) and water to be present. Test tube A had both, which is why only that nail rusted.

5.2 Preventing corrosion

Painting

Keeps the surface from direct contact with air and water

Applying oil or grease

A common method for tools and machine parts

Galvanisation

Coating iron/steel with a thin layer of zinc

Chromium plating

Gives a shiny, rust-resistant surface

Anodising

Building up a thick protective oxide layer

Alloying

E.g. stainless steel, which does not rust at all

A galvanised object stays protected even if the zinc coating gets scratched, because zinc is more reactive than iron and corrodes first, in iron’s place.

5.3 Alloys

Learn by heartDefinition 12

Alloy: a homogeneous (uniform throughout, with no separate layers) mixture of two or more metals, or of a metal and a non-metal. The pure metal is melted, other elements are mixed in fixed proportions, and the mixture is then cooled to room temperature.

Pure iron is very soft and stretches easily when heated. Adding about 0.05% carbon makes it hard and strong. Adding nickel and chromium gives stainless steel, which is hard and does not rust.

Learn by heartDefinition 13

An alloy made with mercury is called an amalgam. An alloy’s electrical conductivity and melting point are generally lower than those of the pure metal.

Alloy	Composition	Notable property
Brass	Copper + zinc	Poor conductor of electricity (unlike pure copper)
Bronze	Copper + tin	Poor conductor of electricity
Solder	Lead + tin	Very low melting point, used to weld electrical wires
Stainless steel	Iron + nickel + chromium	Hard, does not rust

Did you know?

Pure gold (24 carat) is too soft to make jewellery from, so silver or copper is mixed in to harden it. Jewellery in India is typically 22 carat gold: 22 parts pure gold to 2 parts silver or copper.

The Iron Pillar of Delhi, near the Qutub Minar, built more than 1600 years ago and still not rusted — **Figure 4** The Iron Pillar of Delhi, near the Qutub Minar. Built more than 1600 years ago, it stands 8 m tall and weighs 6 tonnes (6000 kg), yet has not rusted to this day. Scientists across the world have studied this feat of ancient Indian metallurgy. Photo: Mark A. Wilson / Wikimedia Commons / Public Domain

A metal in its pure state is rarely of much use on its own: mix in even a small amount of another element and its properties change completely, and that single fact is the foundation of modern engineering.

This chapter in one line

Check yourself before the exam

Can I name the seven physical properties of metals, and the four exceptions to classifying metals by physical properties alone, with examples?
Can I write out the full reactivity series (K to Au) from memory?
Can I explain how NaCl and MgCl₂ form, in terms of electron transfer?
Can I explain the difference between roasting and calcination?
Can I list both conditions needed for rusting, and the ways to prevent it?

All definitions in one place

Amphoteric oxideA metal oxide that reacts with both acids and bases to give a salt and water (Al₂O₃, ZnO)

HydrideA compound formed when a non-metal reacts directly with hydrogen (e.g. HCl); non-metals do not displace hydrogen from dilute acids

Displacement reactionA more reactive metal removes a less reactive metal from its compound

Reactivity seriesA list of metals arranged in order of decreasing reactivity

Ionic / electrovalent compoundA compound formed by the transfer of electrons from a metal to a non-metal

MineralAn element or compound found naturally in the earth’s crust

OreA mineral from which extracting a metal is profitable

MetallurgyThe process of extracting a metal from its ore and then refining it for use

GangueImpurities such as soil and sand mixed with an ore

RoastingConverting a sulphide ore to an oxide by strong heating in excess air

CalcinationConverting a carbonate ore to an oxide by heating in limited air

Reduction / reducing agentReduction is the removal of oxygen from a compound; a reducing agent (e.g. carbon) brings this about

Thermit reactionThe exothermic displacement reaction between Fe₂O₃ and Al, used to join railway tracks

ElectrolysisPassing an electric current through a molten or dissolved compound to break it apart into its elements

Electrode / anode / cathodeThe rods or plates carrying current into an electrolyte; the anode is positively charged, the cathode is negatively charged

ElectrolyteThe molten or dissolved compound that conducts current during electrolysis

Electrolytic refiningPurifying a metal by making the impure metal the anode and the pure metal the cathode

CorrosionThe gradual destruction of a metal’s surface in the presence of air and water (e.g. rust)

GalvanisationCoating iron/steel with a thin layer of zinc to prevent rusting

AlloyA homogeneous mixture of two or more metals, or a metal and a non-metal

AmalgamAn alloy made with mercury

Quick Revision: read this the night before the exam

Physical properties: lustre, hardness, malleability, ductility, thermal/electrical conductivity, sonority, but mercury, iodine, allotropes of carbon and the alkali metals are four big exceptions
Non-metals: opposite of metals physically (not malleable/ductile, poor conductors except graphite); form acidic or neutral oxides, never basic; do not displace hydrogen from acids, form hydrides instead
Metal + oxygen → metal oxide (basic); amphoteric oxides (react with both acids and bases), e.g. Al₂O₃ and ZnO
Metal + water: K, Na violent; Ca slower; Mg only with hot water; Al, Fe, Zn only with steam; Pb, Cu, Ag, Au never
Metal + dilute acid → salt + H₂; order Mg > Al > Zn > Fe; copper does not react
Reactivity series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au
Ionic compounds: formed by electron transfer, do not conduct in the solid state but conduct well as solutions/melts, high melting and boiling points
Extraction depends on reactivity: low reactivity = heating alone; medium = roasting/calcination then reduction with carbon; high = electrolytic reduction
Rusting needs both air and water; prevention: paint, oil, grease, galvanisation, chromium plating, anodising, alloying
Alloying changes a metal’s properties: remember brass, bronze, solder and stainless steel

More for this chapter

Important Questions MCQ / Objective NCERT Solutions Short Notes