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

Get exam-ready revision notes for Class 10 Science Chapter 3: Metals and Non-Metals. These notes cover the physical and chemical properties of elements, amphoteric oxides, the reactivity series, ionic bond formation with dot structures, metallurgy, ore extraction, electrolytic refining, and methods of corrosion prevention.

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NCERT Solutions on Class 10 Science Chapter 3: Metals and Non-Metals
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1. Introduction: Classification of Elements

Elements are classified into metals, non-metals, and metalloids on the basis of their properties.

2. Physical Properties

2.1 Physical Properties of Metals

• Metallic lustre: Metals in their pure state have a shining surface. This property is called metallic lustre.

• Hardness: Metals are generally hard. The hardness varies from metal to metal.

• Malleability: Metals can be beaten into thin sheets. This property is called malleability. Gold and silver are the most malleable metals.

• Ductility: The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal.

• Conduction of heat: Metals are good conductors of heat and have high melting points. Silver and copper are the best conductors. Lead and mercury are comparatively poor conductors.

• Conduction of electricity: Metals are good conductors of electricity. Electric wires are coated with PVC (polyvinylchloride) or rubber-like material because these are insulators.

• Sonority: Metals produce a ringing sound when they strike a hard surface. So school bells are made of metal.

• State: All metals are solids at room temperature, except mercury, which is a liquid.

2.2 Physical Properties of Non-Metals

• Non-metals are either solids or gases, except bromine, which is a liquid.

• They are generally soft, dull, and neither malleable nor ductile.

• Solid non-metals are brittle, so they break into pieces when hammered.

• They are poor conductors of heat and electricity.

• They are not sonorous.

• They generally have low melting and boiling points.

Comparison of Physical Properties

2.3 Exceptions to Remember

• Mercury is a metal but exists as a liquid at room temperature.

• Bromine is the only liquid non-metal at room temperature.

• Gallium and caesium are metals with very low melting points. They melt on your palm.

• Alkali metals (lithium, sodium, potassium) are so soft that a knife can cut them. They have low density and low melting points.

• Iodine is a non-metal, but it is lustrous.

• Diamond, an allotrope of carbon, is the hardest natural substance. It has very high melting and boiling points.

• Graphite is another allotrope of carbon, which conducts electricity.

3. Chemical Properties of Metals

3.1 Reaction with Oxygen

Almost all metals combine with oxygen to form metal oxides.

• $2\text{Cu} + \text{O}_2 \rightarrow 2\text{CuO}$ (copper(II) oxide, black)

• $4\text{Al} + 3\text{O}_2 \rightarrow 2\text{Al}_2\text{O}_3$ (aluminium oxide)

• $2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$ (burns with a dazzling white flame)

Reactivity of Metals with Oxygen

• Potassium and sodium react so vigorously that they catch fire in open air. We keep them immersed in kerosene oil.

• Magnesium, aluminium, zinc and lead get a thin oxide layer at ordinary temperature. The protective oxide layer prevents the metal from further oxidation.

• Iron does not burn on heating, but iron filings burn vigorously in a flame.

• Copper does not burn, but the hot metal gets a black coating of copper(II) oxide.

• Silver and gold do not react with oxygen even at high temperatures.

Nature of Metal Oxides

• Most metal oxides are basic. They turn red litmus blue.

• Most metal oxides are insoluble in water. Sodium oxide and potassium oxide dissolve in water and form alkalis:

  $$\text{Na}_2\text{O(s)} + \text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)} \text{[cite: 1]}$$
  $$\text{K}_2\text{O(s)} + \text{H}_2\text{O(l)} \rightarrow 2\text{KOH(aq)} \text{[cite: 1]}$$

Amphoteric Oxides

Metal oxides that react with both acids as well as bases to produce salt and water are called amphoteric oxides. Examples: aluminium oxide ($\text{Al}_2\text{O}_3$) and zinc oxide ($\text{ZnO}$).

Anodising

Anodising is a process of forming a thick oxide layer on aluminium. A clean aluminium article is made the anode and is electrolysed with dilute sulphuric acid. The oxygen released at the anode reacts with aluminium and makes a thick, protective oxide layer. This oxide layer can be dyed to give an attractive finish.

3.2 Reaction with Water

Not all metals react with water. Their reaction depends on their reactivity.

3.3 Reaction with Acids

• $\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$

• $2\text{Al} + 6\text{HCl} \rightarrow 2\text{AlCl}_3 + 3\text{H}_2$

• $\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$

• $\text{Fe} + 2\text{HCl} \rightarrow \text{FeCl}_2 + \text{H}_2$

• $\text{Fe} + \text{H}_2\text{SO}_4 \rightarrow \text{FeSO}_4 + \text{H}_2$

It is observed from the above reaction the rate of formation of bubbles was the fastest in the case of magnesium. The order of reactivity with dilute acids is $\text{Mg} > \text{Al} > \text{Zn} > \text{Fe}$.

Copper does not react with dilute $\text{HCl}$. No bubbles form and the temperature stays the same.

• Nitric acid: Hydrogen gas is not released when a metal reacts with $\text{HNO}_3$. $\text{HNO}_3$ is a strong oxidising agent. It oxidises the hydrogen to water and itself gets reduced to nitrogen oxides ($\text{N}_2\text{O, NO, NO}_2$). Only magnesium ($\text{Mg}$) and manganese ($\text{Mn}$) react with very dilute $\text{HNO}_3$ to evolve $\text{H}_2$ gas.

• Aqua regia: Aqua regia means “royal water”. It is a freshly prepared mixture of concentrated $\text{HCl}$ and concentrated $\text{HNO}_3$ in the ratio $3 : 1$. Aqua regia is a highly corrosive, fuming liquid. It can dissolve gold and platinum.

3.4 Reaction with Salt Solutions (Displacement)

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

• $\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$

• $\text{Zn} + \text{FeSO}_4 \rightarrow \text{ZnSO}_4 + \text{Fe}$

• $\text{Cu} + 2\text{AgNO}_3 \rightarrow \text{Cu(NO}_3)_2 + 2\text{Ag}$

Activity 3.12: An iron nail dipped in copper sulphate solution shows a reaction. The blue colour of the solution fades to pale green, and a reddish-brown layer of copper deposits on the nail. So iron is more reactive than copper. Copper wire in iron sulphate solution shows no reaction. This is a displacement reaction. It gives better evidence of reactivity than the reactions with oxygen, water or acids.

3.5 The Reactivity Series

The reactivity series is a list of metals arranged in the decreasing order of their reactivity.

Metals above hydrogen displace hydrogen from dilute acids. $\text{Cu, Hg, Ag}$ and $\text{Au}$ do not.

4. Chemical Properties of Non-Metals

• Oxides: Non-metals form oxides that are acidic or neutral. Non-metal oxides turn blue litmus red when dissolved in water.

  $$\text{S} + \text{O}_2 \rightarrow \text{SO}_2 \quad;\quad \text{SO}_2 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_3\text{ (sulphurous acid)} \text{[cite: 1]}$$
  $$\text{C} + \text{O}_2 \rightarrow \text{CO}_2 \quad;\quad \text{CO}_2 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{CO}_3\text{ (carbonic acid)} \text{[cite: 1]}$$

  Carbon monoxide ($\text{CO}$) and water ($\text{H}_2\text{O}$) are neutral oxides.

• Hydrogen: Non-metals react with hydrogen to form hydrides. Examples: $\text{H}_2\text{S}$ and $\text{NH}_3$.

• Acids and water: Non-metals do not displace hydrogen from dilute acids. They do not react with water or steam to give hydrogen.

• With metals: Non-metals gain electrons and form negative ions (anions).

5. How Do Metals and Non-Metals React? (Ionic Bonding)

Atoms react to get a completely filled outermost shell, like noble gases. A metal does this by losing electrons. A non-metal does this by gaining electrons.

Formation of Ionic Compounds

• Sodium chloride ($\text{NaCl}$):

  $$\text{Na } (2,8,1) \rightarrow \text{Na}^+ (2,8) + e^- \text{[cite: 1]}$$
  $$\text{Cl } (2,8,7) + e^- \rightarrow \text{Cl}^- (2,8,8) \text{[cite: 1]}$$

  $\text{Na}^+$ and $\text{Cl}^-$ ions being oppositely charged, attract each other and are held by strong electrostatic forces of attraction to exist as sodium chloride ($\text{NaCl}$). Sodium chloride does not exist as molecules. It exists as aggregates of oppositely charged ions.

• Magnesium chloride ($\text{MgCl}_2$):

  $$\text{Mg } (2,8,2) \rightarrow \text{Mg}^{2+} (2,8) + 2e^- \text{[cite: 1]}$$
  $$2\text{Cl } (2,8,7) + 2e^- \rightarrow 2\text{Cl}^- (2,8,8) \text{[cite: 1]}$$

  The cation in $\text{MgCl}_2$ is $\text{Mg}^{2+}$. The anion is $\text{Cl}^-$.

• Magnesium oxide ($\text{MgO}$) and sodium oxide ($\text{Na}_2\text{O}$):

  $$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \quad,\quad \text{O} + 2e^- \rightarrow \text{O}^{2-} \quad\Longrightarrow\quad \text{Mg}^{2+} + \text{O}^{2-} \rightarrow \text{MgO} \text{[cite: 1]}$$
  $$2\text{Na} \rightarrow 2\text{Na}^+ + 2e^- \quad,\quad \text{O} + 2e^- \rightarrow \text{O}^{2-} \quad\Longrightarrow\quad 2\text{Na}^+ + \text{O}^{2-} \rightarrow \text{Na}_2\text{O} \text{[cite: 1]}$$

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

Properties of Ionic Compounds

• Physical nature: Ionic compounds are solid and somewhat hard because of strong attraction between the ions. They are brittle and break into pieces under pressure.

• Melting and boiling points: They have high melting and boiling points. This is because a large amount of energy is needed to break the strong inter-ionic attraction. ($\text{NaCl}$ melts at $1074\text{ K}$).

• Solubility: They are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.

• Conduction of electricity:

  • Ionic compounds in the solid state do not conduct electricity because the ions cannot move in the rigid structure.

  • In the molten state and in water solution they conduct electricity, because the ions become free to move.

6. Occurrence and Extraction of Metals (Metallurgy)

6.1 Important Terms

Every ore is a mineral, but not every mineral is an ore. The earth’s crust is the major source of metals. Seawater also contains soluble salts such as $\text{NaCl}$ and $\text{MgCl}_2$.

6.2 Occurrence on the Basis of Reactivity

Many ores of many metals are oxides because oxygen is very reactive and very abundant on the earth.

6.3 Steps in Extraction

1. Enrichment of the ore

2. Conversion of the ore into the metal (reduction)

3. Refining of the impure metal

6.4 Enrichment of Ores

Mined ores contain gangue. The impurities must be removed before extraction. The method depends on the difference between the physical or chemical properties of the ore and the gangue. Common methods are hydraulic washing, magnetic separation, froth flotation and leaching.

6.5 Extracting Metals Low in the Activity Series

The oxides of these metals get reduced to metals by heating alone. For example, cinnabar ($\text{HgS}$) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide ($\text{HgO}$). Mercuric oxide is then reduced to mercury on further heating.

• Mercury from cinnabar ($\text{HgS}$):

  $$2\text{HgS(s)} + 3\text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{HgO(s)} + 2\text{SO}_2\text{(g)} \text{[cite: 1]}$$
  $$2\text{HgO(s)} \xrightarrow{\text{Heat}} 2\text{Hg(l)} + \text{O}_2\text{(g)} \text{[cite: 1]}$$

• Copper from copper glance ($\text{Cu}_2\text{S}$):

  $$2\text{Cu}_2\text{S} + 3\text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{Cu}_2\text{O(s)} + 2\text{SO}_2\text{(g)} \text{[cite: 1]}$$
  $$2\text{Cu}_2\text{O} + \text{Cu}_2\text{S} \xrightarrow{\text{Heat}} 6\text{Cu(s)} + \text{SO}_2\text{(g)} \text{[cite: 1]}$$

6.6 Extracting Metals in the Middle of the Activity Series

The metals in the middle of the activity series such as iron, zinc, lead, copper, etc., are usually present as sulphides or carbonates in nature. It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. So we first convert the ore into metal oxide.

• The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting.

• The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination.

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc:

• Displacement reactions to reduce oxides: Highly reactive metals such as $\text{Na, Ca}$ and $\text{Al}$ can displace metals of lower reactivity from their compounds. So they act as reducing agents.

  $$3\text{MnO}_2 + 4\text{Al} \rightarrow 3\text{Mn} + 2\text{Al}_2\text{O}_3 + \text{Heat} \text{[cite: 1]}$$

  In this reaction, $\text{MnO}_2$ is reduced and $\text{Al}$ is oxidised.

• Thermit reaction: These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metal is produced in the molten state. The reaction of iron(III) oxide with aluminium is used to join railway tracks and cracked machine parts. This reaction is known as the thermit reaction.

  $$\text{Fe}_2\text{O}_3 + 2\text{Al} \rightarrow 2\text{Fe (molten)} + \text{Al}_2\text{O}_3 + \text{Heat} \text{[cite: 1]}$$

6.7 Extracting Metals at the Top of the Activity Series

These metals cannot be obtained by heating with carbon. Carbon cannot reduce their oxides, because these metals have more affinity for oxygen than carbon. These metals are obtained by electrolytic reduction of their molten salts.

Sodium, magnesium and calcium come from the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas, chlorine is liberated at the anode (the positively charged electrode). The reactions are:

• At cathode: $\text{Na}^+ + e^- \rightarrow \text{Na}$

• At anode: $2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$

Aluminium comes from the electrolytic reduction of molten aluminium oxide.

6.8 Refining of Metals

The metals produced by various reduction processes are not pure. The most widely used method of refining is electrolytic refining. Copper, zinc, tin, nickel, silver and gold are refined this way.

In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte.

When current passes through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode.

Soluble impurities go into the solution. Insoluble impurities settle below the anode as anode mud.

• At anode: $\text{Cu} \rightarrow \text{Cu}^{2+} + 2e^-$

• At cathode: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$

7. Corrosion

Corrosion is the slow damage of the surface of a metal when it reacts with air, moisture or gases around it.

Conditions for Rusting (Activity 3.14)

Three iron nails are kept in three test tubes:

Conclusion: Iron rusts only when it is in contact with both air (oxygen) and water (moisture).

Prevention of Corrosion

• The rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising or making alloys.

• Painting, oiling and greasing keep air and moisture away from the metal.

• Galvanisation is the process of coating iron or steel with a thin layer of zinc to protect it from rusting. The galvanised article stays protected even if the zinc coating breaks. The reason: zinc is more reactive than iron, so zinc reacts first and protects the iron.

• Alloying improves the properties of a metal.

Question: “Why does a galvanised article not rust even when the zinc layer is scratched?”

Answer: Zinc is more reactive than iron. So zinc gets oxidised in place of iron and protects it.

8. Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the main metal and then dissolving the other elements in it in definite proportions. It is then cooled to room temperature.

Why Do We Make Alloys?

• Alloying changes the properties of a metal and gives us the properties we want.

• Pure iron is very soft and stretches easily when hot. Adding about $0.05\%$ carbon makes it hard and strong.

• When iron is mixed with nickel and chromium, we get stainless steel. It is hard and does not rust.

• The electrical conductivity and the melting point of an alloy are less than those of the pure metals.

Common Alloys

• If one of the metals is mercury, then the alloy is known as an amalgam.

• Do you know?

  • Pure gold is 24 carat gold. It is very soft, so it is not suitable for jewellery. We alloy it with silver or copper to make it hard. In India, 22 carat gold is used for ornaments. This means 22 parts of pure gold are alloyed with 2 parts of copper or silver.

  • The iron pillar near the Qutub Minar in Delhi is more than 1600 years old. It is $8\text{ m}$ high and weighs $6\text{ tonnes}$. It shows the great skill of ancient Indian metallurgy, because it has not rusted.

9. Quick Revision

Key Definitions

• Malleability: The property by which metals can be beaten into thin sheets.

• Ductility: The property by which metals can be drawn into thin wires.

• Sonorous: A metal that produces a ringing sound when it strikes a hard surface.

• Amphoteric oxide: A metal oxide that reacts with both acids and bases to give salt and water.

• Reactivity series: A list of metals arranged in decreasing order of reactivity.

• Ionic compound: A compound formed by the transfer of electrons from a metal to a non-metal.

• Roasting: Heating a sulphide ore strongly in excess air to get an oxide.

• Calcination: Heating a carbonate ore strongly in limited air to get an oxide.

• Thermit reaction: The reaction of iron(III) oxide with aluminium to give molten iron.

• Corrosion: The slow damage of a metal by air, moisture or gases around it.

• Galvanisation: Coating iron or steel with a thin layer of zinc.

• Alloy: A homogeneous mixture of two or more metals, or a metal and a non-metal.