CBSE Class 10 Science Metals and non metals Notes

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CHAPTER 3: METALS AND NON-METALS

Metals occupy the bulk of the periodic table, while non-metallic elements can only be found on the right-hand side of the periodic table. A diagonal line, drawn from boron (B) to polonium (Po), separates the metals from the non-metals. Most elements on this line are metalloids, sometimes called semiconductors, because these elements exhibit electrical properties intermediate to both conductors and insulators. Elements to the lower left of this dividing line are called metals, while elements to the upper right of the dividing line are called non-metals.

On the basis of their general physical and chemical properties, every element in the periodic table can be termed either a metal or a non-metal.

PHYSICAL PROPERTIES OF METALS

Physical state: Metals are solids at room temperature, e.g. sodium, aluminium, potassium, magnesium. There are exceptions to this: mercury and gallium are metals but they are in the liquid state at room temperature.

Lustre: Metals have a shining surface, called lustre, when freshly prepared. They have the quality of reflecting light from their surface and they can be polished, e.g. metals like gold, silver and copper show this property.

Malleability: Metals can be beaten into thin sheets. This property is called malleability. Due to this property, metals can be rolled into sheets, e.g. aluminium, copper and zinc can be beaten into sheets.

Ductility: Metals can be drawn into thin wires. This property is called ductility. For example, 100 grams of silver can be drawn into a thin wire about 200 metres long.

Hardness: Metals are generally hard, e.g. iron, cobalt, nickel. There are a few exceptions to this: sodium and potassium are soft and can be cut with a knife.

Sound: Metals produce a ringing sound, so metals are called sonorous. The sound of metals is also known as metallic sound. This is the reason metal wires are used in making musical instruments.

Conduction: Generally, metals are good conductors of heat and electricity because they have free electrons. Silver and copper are the two best conductors. Relatively, lead and bismuth are poor conductors of heat and electricity.

Density: Metals generally have high density and they are heavy. Iridium and osmium have the highest densities, while lithium has the lowest density.

Melting and boiling points: Metals usually have high melting and boiling points. For example, iron, cobalt and nickel have high melting and boiling points. Tungsten has the highest melting point. There are some exceptions to this; for example, most of the alkali metals have low melting and boiling points.

Strength: Most of the metals are strong and have high tensile strength. Because of this, big structures are made using metals such as copper and iron.

Colour: Most of the metals are grey in colour. But gold and copper are exceptions.

METALS (mind map): remaining labels

• Solid state at room temp: Except Mercury (Liquid)
• High density: Atoms are closely packed
• Good conductors of electricity: Due to the presence of mobile ions
• Good heat conductors: Due to the presence of mobile ions
• Ductile and malleable: Can be bent and stretched without breaking
• Strong and tough
• High Melting and Boiling point: Due to strong forces (metallic) except Mercury and alkali metals
• Shiny appearance

INTEXT QUESTIONS (PAGE 40)

Give an example of a metal which (i) is a liquid at room temperature, (ii) can be easily cut with a knife, (iii) is the best conductor of heat, (iv) is a poor conductor of heat.

Answer: (i) Metal that exists in the liquid state at room temperature: mercury.
(ii) Metal that can be easily cut with a knife: sodium.
(iii) Metal that is the best conductor of heat: silver.
(iv) Metals that are poor conductors of heat: mercury and lead.

Explain the meanings of malleable and ductile.

Answer: Malleable: Substances that can be beaten into thin sheets are called malleable. For example, most of the metals are malleable.
Ductile: Substances that can be drawn into thin wires are called ductile. For example, most of the metals are ductile.

CHEMICAL PROPERTIES OF METALS

REACTION WITH OXYGEN

Most of the metals form the respective metal oxides when they react with oxygen.

Metal + Oxygen → Metal oxide

Examples:

Magnesium: Magnesium gives magnesium oxide when it reacts with oxygen. Magnesium burns with a dazzling light in air and produces a lot of heat.

2Mg + O₂ → 2MgO

Aluminium: Aluminium does not react with oxygen at room temperature, but it gives aluminium oxide when burnt in air.

4Al + 3O₂ → 2Al₂O₃

Potassium: Potassium forms potassium oxide when it reacts with oxygen.

4K + O₂ → 2K₂O

Sodium: Sodium forms sodium oxide when it reacts with oxygen.

4Na + O₂ → 2Na₂O

Lithium, potassium, sodium, etc. are known as alkali metals. Alkali metals react vigorously with oxygen.

Iron: Iron does not react with oxygen at room temperature. But when iron is heated strongly in air, it gives iron oxide.

3Fe + 2O₂ → Fe₃O₄

Iron filings give sparks in a flame when burnt.

Copper: Copper does not react with oxygen at room temperature, but when burnt in air, it gives copper oxide.

2Cu + O₂ → 2CuO

Zinc: Zinc does not react with oxygen at room temperature. But it gives zinc oxide when heated strongly in air.

2Zn + O₂ → 2ZnO

REACTION OF METALS WITH WATER

Metals form the respective metal hydroxide and hydrogen gas when they react with water.

Metal + Water → Metal hydroxide + Hydrogen

Most of the metals do not react with water. However, alkali metals react vigorously with water.

Examples:

Sodium: Sodium forms sodium hydroxide and liberates hydrogen gas along with a lot of heat when it reacts with water.

2Na + 2H₂O → 2NaOH + H₂

Aluminium: The reaction of aluminium with cold water is too slow to come into notice. But when steam is passed over aluminium, aluminium oxide and hydrogen gas are produced.

2Al + 3H₂O → Al₂O₃ + 3H₂

Zinc: Zinc produces zinc oxide and hydrogen gas when steam is passed over it. Zinc does not react with cold water.

Zn + H₂O → ZnO + H₂

Iron: The reaction of iron with cold water is very slow and comes into notice after a long time. Iron forms rust (hydrated iron oxide) when it reacts with moisture present in the atmosphere. Iron oxide and hydrogen gas are formed by passing steam over iron.

3Fe + 4H₂O → Fe₃O₄ + 4H₂

Potassium: Potassium forms potassium hydroxide and liberates hydrogen gas along with a lot of heat when it reacts with water.

2K + 2H₂O → 2KOH + H₂

Calcium: Calcium forms calcium hydroxide along with hydrogen gas and heat when it reacts with water.

Ca + 2H₂O → Ca(OH)₂ + H₂

Magnesium: Magnesium reacts with water slowly and forms magnesium hydroxide and hydrogen gas.

Mg + 2H₂O → Mg(OH)₂ + H₂

When steam is passed over magnesium, magnesium oxide and hydrogen gas are formed.

Mg + H₂O → MgO + H₂

REACTION OF METALS WITH DILUTE ACID

Metals form the respective salts when they react with dilute acid.

Metal + dil. acid → Metal salt + Hydrogen

Examples:

Aluminium with dilute hydrochloric acid: Aluminium chloride and hydrogen gas are formed.

2Al + 6HCl → 2AlCl₃ + 3H₂

Zinc with dilute sulphuric acid: Zinc sulphate and hydrogen gas are formed. This method is used in the laboratory to produce hydrogen gas.

Zn + H₂SO₄ → ZnSO₄ + H₂

Sodium with dilute hydrochloric acid: Sodium chloride and hydrogen gas are formed.

2Na + 2HCl → 2NaCl + H₂

Potassium with dilute sulphuric acid: Potassium sulphate and hydrogen gas are formed.

2K + H₂SO₄ → K₂SO₄ + H₂

Magnesium with dilute hydrochloric acid: Magnesium chloride and hydrogen gas are formed.

Mg + 2HCl → MgCl₂ + H₂

Copper, gold and silver are known as noble metals. These do not react with water or dilute acids.

METAL OXIDES: CHEMICAL PROPERTIES

Metal oxides are basic in nature. An aqueous solution of a metal oxide turns red litmus blue.

REACTION OF METAL OXIDES WITH WATER

Most of the metal oxides are insoluble in water. Alkali metal oxides are soluble in water, and they give a strong base when dissolved in water.

Examples:

Sodium oxide gives sodium hydroxide when it reacts with water.

Na₂O + H₂O → 2NaOH

Magnesium oxide gives magnesium hydroxide with water.

MgO + H₂O → Mg(OH)₂

Potassium oxide gives potassium hydroxide when it reacts with water.

K₂O + H₂O → 2KOH

Zinc oxide and aluminium oxide: Aluminium oxide and zinc oxide are insoluble in water. They are amphoteric in nature. An amphoteric substance shows both acidic and basic character. It reacts with a base like an acid and reacts with an acid like a base.

When zinc oxide reacts with sodium hydroxide, it behaves like an acid. In this reaction, sodium zincate and water are formed.

ZnO + 2NaOH → Na₂ZnO₂ + H₂O

Zinc oxide behaves like a base when it reacts with an acid. Zinc oxide gives zinc chloride and water on reaction with hydrochloric acid.

ZnO + 2HCl → ZnCl₂ + H₂O

In a similar way, aluminium oxide behaves like a base when it reacts with an acid and behaves like an acid when it reacts with a base. Aluminium oxide gives sodium aluminate along with water when it reacts with sodium hydroxide.

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Aluminium oxide gives aluminium chloride along with water when it reacts with hydrochloric acid.

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

REACTIVITY SERIES OF METALS

A series of metallic elements arranged in the increasing or decreasing order of their reactivity is called a reactivity series of metals. In the reactivity series, copper, gold and silver are at the bottom and hence least reactive. These metals are known as noble metals.

The most active metal, potassium, is at the top of the list and the least reactive metal, gold, is at the bottom of the list. Although hydrogen is a non-metal, it is included in the activity series because it behaves like a metal in most chemical reactions, i.e. the hydrogen ion has a positive charge (H⁺) like other metals.

The following points become evident from the activity series of metals:

• The higher the metal in the series, the more reactive it is, i.e. its reaction is fast and more exothermic.
• This also implies that the more reactive a metal, the more difficult it is to extract it from its ore (the reverse reaction becomes more difficult). The metal is also more susceptible to corrosion with oxygen and water.
• The reactivity series can be established by observation of the reaction of metals with water, oxygen or acids.

Within the general reactivity or activity series, there are some periodic table trends:

1. Down Group 1 (I), the “alkali metals”, the activity increases: Cs > Rb > K > Na > Li.
2. Down Group 2 (II), the activity increases, e.g. Ca > Mg.
3. In the same period, the Group 1 metal is more reactive than the Group 2 metal, and the Group 2 metal is more reactive than the Group 3 metal, and all three are more reactive than the “transition metals”, e.g. Na > Mg > Al (in Period 3) and K > Ca > Ga > Fe/Cu/Zn etc. (in Period 4).

REACTION OF METALS WITH SOLUTIONS OF OTHER METAL SALTS

The reaction of a metal with the solution of another metal salt is a displacement reaction. In this reaction, the more reactive metal displaces the less reactive metal from its salt.

Metal A + Salt of metal B → Salt of metal A + Metal B

Examples: Iron displaces copper from copper sulphate solution.

Fe + CuSO₄ → FeSO₄ + Cu

Similarly, aluminium and zinc displace copper from the solution of copper sulphate.

2Al + 3CuSO₄ → Al₂(SO₄)₃ + 3Cu

Zn + CuSO₄ → ZnSO₄ + Cu

In all the above examples, iron, aluminium and zinc are more reactive than copper. That is why they displace copper from its salt solution.

When copper is dipped in the solution of silver nitrate, it displaces silver and forms copper nitrate.

Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag

In this reaction copper is more reactive than silver and hence displaces silver from silver nitrate solution, forming copper nitrate.

Silver does not react with copper sulphate solution, because silver is less reactive than copper and is not able to displace copper from its salt solution.

Ag + CuSO₄ → No reaction

Similarly, when gold is dipped in the solution of copper sulphate, no reaction takes place, because copper is more reactive than gold.

Au + CuSO₄ → No reaction

In a similar way, no reaction takes place when copper is dipped in the solution of aluminium nitrate, because copper is less reactive than aluminium.

Cu + Al(NO₃)₃ → No reaction

PHYSICAL PROPERTIES OF NON-METALS

Physical state: Non-metals can exist in the solid, liquid or gaseous state at room temperature. For example, carbon, sulphur, phosphorus and iodine are solids, bromine is a liquid, while oxygen, nitrogen and chlorine are gases at room temperature.

Lustre: Non-metals do not have lustre. They do not reflect light from their surface (exceptions: diamond and iodine). Non-metals have a dull appearance. For example, sulphur, phosphorus and carbon show this property.

Malleability: Non-metals are non-malleable. If solids, they are brittle, i.e. they break or shatter on hammering. For example, coal, sulphur and phosphorus are brittle.

Ductility: Non-metals cannot be drawn into thin wires. So they are not ductile.

Hardness: Non-metals are usually not hard. They are soft. For example, coal, sulphur and phosphorus are soft. Diamond is an exception to this: it is the hardest substance known.

Sonority: Non-metals are not sonorous, i.e. they do not produce a typical sound on being hit.

Conduction: Non-metals are usually poor conductors of heat and electricity. However, carbon in the form of gas carbon and graphite is an exception to this. These forms of carbon are good conductors of electricity.

Density: Non-metals which are gases have low density. Solid non-metals have low to moderate density. For example, phosphorus, sulphur and boron have densities 1.82, 2.07 and 2.34 g/cm³ respectively. However, diamond has a high density, about 3.5 g/cm³.

Melting and boiling points: Non-metals usually have low melting and boiling points. For example, phosphorus, sulphur and iodine have melting points 44 °C, 115 °C and 114 °C respectively and boiling points 280 °C, 445 °C and 184 °C respectively. However, carbon, silicon and boron possess very high melting and boiling points.

Tensile strength: Non-metals have low tensile strength, i.e. they have no tenacity.

Colour: Non-metals are of many colours.

CHEMICAL PROPERTIES OF NON-METALS

REACTION OF NON-METALS WITH OXYGEN

Non-metals form the respective oxide when they react with oxygen.

Non-metal + Oxygen → Non-metal oxide

Examples: When carbon reacts with oxygen, carbon dioxide is formed along with the production of heat.

C + O₂ → CO₂ + Heat

When carbon is burnt in an insufficient supply of air, it forms carbon monoxide. Carbon monoxide is a toxic substance. Inhaling carbon monoxide may prove fatal.

2C + O₂ → 2CO + Heat

Sulphur gives sulphur dioxide when it reacts with oxygen. Sulphur catches fire when exposed to air.

S + O₂ → SO₂

When hydrogen reacts with oxygen it gives water.

2H₂ + O₂ → 2H₂O

NON-METAL OXIDES

Non-metal oxides are acidic in nature. A solution of a non-metal oxide turns blue litmus red.

Examples: Carbon dioxide gives carbonic acid when dissolved in water.

CO₂ + H₂O → H₂CO₃

Sulphur dioxide gives sulphurous acid when dissolved in water.

SO₂ + H₂O → H₂SO₃

Sulphur dioxide gives sulphur trioxide when it reacts with oxygen (in the presence of a catalyst).

2SO₂ + O₂ → 2SO₃

Sulphur trioxide gives sulphuric acid when dissolved in water.

SO₃ + H₂O → H₂SO₄

REACTION OF NON-METALS WITH CHLORINE

Non-metals give the respective chloride when they react with chlorine gas.

Non-metal + Chlorine → Non-metal chloride

Examples: Hydrogen gives hydrogen chloride and phosphorus gives phosphorus trichloride when they react with chlorine.

H₂ + Cl₂ → 2HCl

P₄ + 6Cl₂ → 4PCl₃

REACTION OF METALS AND NON-METALS

Many metals form ionic bonds when they react with non-metals. Compounds so formed are known as ionic compounds.

Ions: Positively or negatively charged atoms are known as ions. Ions are formed because of loss or gain of electrons. Atoms form ions to obtain the electronic configuration of the nearest noble gas, that is, to obtain a stable configuration.

Positive ion: A positive ion is formed because of loss of electrons by an atom. Following are some examples of positive ions.

Sodium forms a sodium ion because of loss of one electron. Because of the loss of one electron, one positive charge comes over sodium.

Na → Na⁺ + e⁻

Similarly, potassium gets one positive charge by loss of one electron.

K → K⁺ + e⁻

Magnesium forms a positive ion because of loss of two electrons. Two positive charges come over magnesium because of the loss of two electrons.

Mg → Mg²⁺ + 2e⁻

Similarly, calcium gets two positive charges over it by loss of two electrons.

Ca → Ca²⁺ + 2e⁻

Negative ion: A negative ion is formed because of gain of electrons. Some examples are given below.

Chlorine gains one electron in order to achieve a stable configuration. After the gain of one electron, chlorine gets one negative charge over it, forming the chloride ion.

Cl + e⁻ → Cl⁻

Similarly, fluorine gets one negative charge over it by gain of one electron, forming the fluoride ion, in order to achieve a stable configuration.

F + e⁻ → F⁻

Oxygen gets two negative charges over it by gain of two electrons, forming the oxide ion, in order to obtain a stable configuration.

O + 2e⁻ → O²⁻

USES OF METALS

Metals find a number of applications. Some of them are given below.

• Zinc is used for galvanizing iron, in anti-corrosion material, in medicinal fields and in alloys.
• Iron is used as a construction material in bridges, houses, ships, etc. Iron, in the form of steel, is used for making domestic utensils.
• Tin is used for soldering, for preparing foils, for metal coatings to prevent chemical action and corrosion, for panel lighting, etc.
• Lead is used in making water pipes, in pigments, batteries, in alloys, etc.
• Titanium finds extensive use in the aircraft industry.
• Pure metals which display zero resistance to electrical current are called superconductors. Hg and Nb are examples of superconductors. They become superconductors below critical temperatures of 4.2 K and 9.2 K respectively. Superconductors have many applications in research and industry.
• Almost all metals, including Zr and Ti, find wide applications in atomic and space programmes and experiments.
• Mercury is used in thermometers.
• Silver, gold and platinum are precious metals and they are used in making ornaments.
• Radioactive metals like uranium and plutonium are used in nuclear power plants to produce atomic energy via nuclear fission.

USES OF NON-METALS

Non-metals find a number of applications. Some of them are given below.

• Sulphur is used in making compounds like sulpha drugs, sulphuric acid, in matches, in gunpowder, for vulcanization of rubber, etc.
• Boron, in the form of the compound borax, is used in making skin ointments.
• Phosphorus is used in making crackers.
• Oxygen is used for respiration.
• Chlorine, in the form of bleaching powder, is used for purification of water.
• Carbon is used as a fuel, as electrodes (graphite) and as a reducing agent in metallurgy.
• Oxygen, hydrogen and nitrogen are used by all living things; they are the “building blocks” of life.
• Iodine is used to prevent thyroid problems.
• Bromine is used in the preparation of dyes.
• Some compounds of fluorine (such as sodium fluoride and stannous fluoride) are added to toothpastes to prevent dental decay or formation of cavities.

INTEXT QUESTIONS (PAGE 46)

Why is sodium kept immersed in kerosene oil?

Answer: Sodium and potassium are very reactive metals and combine explosively with air as well as water. Hence, they catch fire if kept in the open. Therefore, to prevent accidental fires and accidents, sodium is stored in kerosene oil.

Write equations for the reactions of (i) iron with steam, (ii) calcium and potassium with water.

Answer: (i) 3Fe + 4H₂O → Fe₃O₄ + 4H₂ (iron + steam gives iron(II,III) oxide + hydrogen)
(ii) Ca + 2H₂O → Ca(OH)₂ + H₂ + heat (calcium + water gives calcium hydroxide + hydrogen)
    2K + 2H₂O → 2KOH + H₂ + heat (potassium + water gives potassium hydroxide + hydrogen)

Samples of four metals A, B, C and D were taken and added to the following solutions one by one. The results obtained have been tabulated as follows.

Use the table above to answer the following questions about metals A, B, C and D. (i) Which is the most reactive metal? (ii) What would you observe if B is added to a solution of copper(II) sulphate? (iii) Arrange the metals A, B, C and D in the order of decreasing reactivity.

Answer: Explanation:
A + FeSO₄: no reaction, i.e. A is less reactive than iron.
A + CuSO₄: displacement, i.e. A is more reactive than copper.
B + FeSO₄: displacement, i.e. B is more reactive than iron.
B + ZnSO₄: no reaction, i.e. B is less reactive than zinc.
C + FeSO₄: no reaction, i.e. C is less reactive than iron.
C + CuSO₄: no reaction, i.e. C is less reactive than copper.
C + ZnSO₄: no reaction, i.e. C is less reactive than zinc.
C + AgNO₃: displacement, i.e. C is more reactive than silver.
D + FeSO₄/CuSO₄/ZnSO₄/AgNO₃: no reaction, i.e. D is less reactive than iron, copper, zinc and silver.
From the above, we obtain:
(i) B is the most reactive metal.
(ii) If B is added to a solution of copper(II) sulphate, it would displace copper (B + CuSO₄: displacement).
(iii) The order of decreasing reactivity is: B > A > C > D.

The reactivity order from the table is: Zn > B > Fe > A > Cu > C > Ag > D (most reactive at the top).

Which gas is produced when dilute hydrochloric acid is added to a reactive metal? Write the chemical reaction when iron reacts with dilute H₂SO₄.

Answer: Hydrogen gas is evolved when dilute hydrochloric acid is added to a reactive metal. When iron reacts with dilute H₂SO₄, iron(II) sulphate is formed with the evolution of hydrogen gas.
Fe + H₂SO₄ → FeSO₄ + H₂

What would you observe when zinc is added to a solution of iron(II) sulphate? Write the chemical reaction that takes place.

Answer: Zinc is more reactive than iron. Therefore, if zinc is added to a solution of iron(II) sulphate, it would displace iron from the solution (the green colour of the solution fades and iron is deposited).
Zn + FeSO₄ → ZnSO₄ + Fe

IONIC BONDS

Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the non-metal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a metal, and an electron acceptor, a non-metal.

Ionic bonding is observed because metals have few electrons in their outermost orbits. By losing those electrons, these metals can achieve a noble gas configuration and satisfy the octet rule. Similarly, non-metals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve a noble gas configuration. In ionic bonding, more than one electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the number of electrons donated or received. In ionic bonds, the net charge of the compound must be zero.

FORMATION OF IONIC BOND

The positive ions (cations) and negative ions (anions) that are formed experience electrostatic forces and get attracted to form a chemical bond. As this bond is between charged particles known as ions, it is called an ionic bond. Because the forces are electrostatic, the bond is also called the electrostatic bond. As the valence concept has been explained in terms of electrons, it is also called the electrovalent bond.

Thus, we can define an ionic bond as follows: the electrostatic attractive force that keeps a cation and an anion together to form a new electrically neutral entity is called an ionic bond.

EXAMPLES

FORMATION OF SODIUM CHLORIDE (NaCl)

In sodium chloride, sodium is a metal (alkali metal) and chlorine is a non-metal.

Atomic number of sodium = 11; electronic configuration of sodium: 2, 8, 1; number of electrons in the outermost orbit = 1; valence electrons = 1.
Atomic number of chlorine = 17; electronic configuration of chlorine: 2, 8, 7; electrons in the outermost orbit = 7; therefore, valence electrons = 7.

Sodium has one valence electron and chlorine has seven valence electrons. Sodium requires losing one electron to obtain a stable configuration and chlorine requires gaining one electron in order to obtain a stable electronic configuration. Thus, in order to obtain a stable configuration, sodium transfers one electron to chlorine.

After the loss of one electron, sodium gets one positive charge (+) and chlorine gets one negative charge after the gain of one electron. Sodium chloride is formed because of the transfer of electrons. Thus, an ionic bond is formed between sodium and chlorine. Since sodium chloride is formed because of an ionic bond, it is called an ionic compound. In a similar way, potassium chloride (KCl) is formed.

FORMATION OF MAGNESIUM CHLORIDE (MgCl₂)

The atomic number of magnesium is 12; electronic configuration of magnesium: 2, 8, 2; number of electrons in the outermost orbit = 2; valence electrons = 2.
Atomic number of chlorine = 17; electronic configuration of chlorine: 2, 8, 7; electrons in the outermost orbit = 7; therefore, valence electrons = 7.

Magnesium loses two electrons in order to obtain a stable electronic configuration. Each of the two chlorine atoms gains one electron lost by magnesium to obtain a stable electronic configuration. The bonds so formed between magnesium and chlorine are ionic bonds and the compound (magnesium chloride) is an ionic compound.

FORMATION OF CALCIUM CHLORIDE (CaCl₂)

Atomic number of calcium is 20; electronic configuration of calcium: 2, 8, 8, 2; number of electrons in the outermost orbit = 2; valence electrons = 2. Valence electrons of chlorine = 7.

Calcium loses two electrons in order to achieve a stable electronic configuration. Each of the two chlorine atoms, on the other hand, gains one electron lost by calcium to get stability. By losing two electrons, calcium gets two positive charges over it. Each of the chlorine atoms gets one negative charge over it.

The bonds formed in calcium chloride are ionic bonds and the compound (calcium chloride) is an ionic compound. In a similar way, barium chloride is formed.

FORMATION OF CALCIUM OXIDE (CaO)

Valence electrons of calcium = 2. Atomic number of oxygen is 8; electronic configuration of oxygen is 2, 6; number of electrons in the outermost orbit = 6; valence electrons = 6.

Calcium loses two electrons and gets two positive charges over it in order to get stability. Oxygen gains the two electrons lost by calcium and thus gets two negative charges over it. The bond formed in calcium oxide is an ionic bond. Calcium oxide is an ionic compound. In a similar way, magnesium oxide is formed.

PROPERTIES OF IONIC COMPOUNDS

Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.

Melting and boiling points: Ionic compounds have high melting and boiling points. This is because a considerable amount of energy is required to break the strong inter-ionic attraction.

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

Conduction of electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

INTEXT QUESTIONS (PAGE 49)

(i) Write the electron-dot structures for sodium, oxygen and magnesium. (ii) Show the formation of Na₂O and MgO by the transfer of electrons. (iii) What are the ions present in these compounds?

Answer: (i) The representation of elements with valence electrons as dots around the elements is referred to as the electron-dot structure of the elements.

(iii) The ions present in Na₂O are Na⁺ and O²⁻ ions, and in MgO are Mg²⁺ and O²⁻ ions.

Why do ionic compounds have high melting points?

Answer: Ionic compounds have strong electrostatic forces of attraction between the ions. Therefore, it requires a lot of energy to overcome these forces. That is why ionic compounds have high melting points.

OCCURRENCE AND EXTRACTION OF METALS

Metals occur in nature in free as well as combined form. Metals having low reactivity show little affinity for air, moisture, carbon dioxide or other non-metals present in nature. Such metals may remain in elemental or native (free) state in nature. Such metals are called “noble metals” as they show the least chemical reactivity. For example, gold, silver, mercury and platinum occur in the free state.

On the other hand, most of the metals are active and combine with air, moisture, carbon dioxide and non-metals like oxygen, sulphur, halogens, etc. to form their compounds, like oxides, sulphides, carbonates, halides and silicates, i.e. they occur in nature in a combined state.

A naturally occurring material in which a metal or its compound occurs is called a mineral. A mineral from which a metal can be extracted economically is called an ore.

An ore is a mineral in which a metal is present in appreciable quantities and from which the metal can be extracted economically.

Metals found at the bottom of the reactivity series are least reactive and are often found in nature in the free state, such as gold, silver, copper, etc. Copper and silver are also found in the form of sulphide and oxide ores.

Metals found in the middle of the reactivity series, such as Zn, Fe, Pb, etc., are usually found in the form of oxides, sulphides or carbonates.

Metals found at the top of the reactivity series are never found in the free state as they are very reactive, e.g. K, Na, Ca, Mg and Al, etc.

Many metals are found in the form of oxides because oxygen is abundant in nature and is very reactive.

TABLE: SOME IMPORTANT ORES

EXTRACTION OF METALS

Metals can be categorised into three parts on the basis of their reactivity: most reactive, medium reactive and least reactive.

CONCENTRATION OF THE ORE

Ores that are mined from the earth are usually contaminated with a large amount of impurities such as soil and sand, etc.

Concentration or dressing means simply getting rid of as much of the unwanted rocky material as possible before the ore is converted into the metal. The impurities like clay are called gangue.

Enrichment of the ore: Physical methods are used to enrich the ore. In many cases, it is possible to separate the metal compound from unwanted rocky material by physical means. A common example of this involves froth flotation.

The physical methods adopted in dressing (or enriching) the ore depend upon the difference between the physical properties of the ore and the gangue.

METHODS USED TO ENRICH THE ORE

Hand picking: If the ore particles and the impurities differ in one of the properties like colour, size, etc., then using that property the ore particles are hand-picked, separating them from the other impurities.

Washing: Ore particles are crushed and kept on a sloping surface. They are washed with a controlled flow of water. The less dense impurities are carried away by the water flow, leaving the denser ore particles behind.

Froth flotation: This method is mainly useful for sulphide ores, which are not wetted by water, whereas the impurities get wetted. The ore with impurities is finely powdered and kept in water taken in a flotation cell. Air under pressure is blown to produce froth in the water. The froth so produced takes the ore particles to the surface, whereas the impurities settle at the bottom. The froth is separated and washed to get the ore particles.

Magnetic separation: If the ore or the impurity, one of them, is magnetic and the other non-magnetic, they are separated using electromagnets.

EXTRACTION OF CRUDE METAL FROM THE ORE

After concentration and dressing of the ore obtained from the earth, we get a concentrated or enriched ore. To extract the metal from this enriched ore, it is converted into the metal oxide, which is then reduced to get the metal with certain impurities. Extraction of the metal from its ores depends on the reactivity of the metal.

Arranging the metals in decreasing order of their reactivity gives the activity series. The classification of the metals is on the basis of their reactivity.

A) EXTRACTION OF METALS AT THE TOP OF THE ACTIVITY SERIES

(K, Na, Ca, Mg and Al.) Simple chemical reduction methods like heating with C, CO, etc. to reduce the ores of these metals are not feasible. The temperature required for the reduction is too high and expensive. To make the process economical, electrolysis methods have to be adopted. Again, electrolysis of their aqueous solutions is also not feasible because water in the solution would be discharged at the cathode in preference to the metal ions.

The only viable method is to extract these metals by electrolysis of their fused compounds. For example, to extract Na from NaCl, fused NaCl is electrolysed with a steel cathode (−) and a graphite anode (+). The metal (Na) is deposited at the cathode and chlorine is liberated at the anode.

At the cathode: 2Na⁺ + 2e⁻ → 2Na
At the anode: 2Cl⁻ → Cl₂ + 2e⁻

SUMMARY: BEHAVIOUR OF METALS WITH OXYGEN, WATER, STEAM, ACIDS AND CHLORINE

B) EXTRACTION OF METALS IN THE MIDDLE OF THE ACTIVITY SERIES

(Zinc, iron, tin, lead and copper.) The ores of these metals are generally present as sulphides or carbonates in nature. Therefore, prior to reduction, the ores of these metals must be converted into metal oxides. Sulphide ores are converted into oxides by heating them strongly in excess of air. This process is known as roasting. Generally, the sulphide ores are roasted to convert them into oxides before reducing them to the metal. Carbonate ores are converted into oxides by heating strongly in limited air; this is called calcination.

Example:

2PbS + 3O₂ → 2PbO + 2SO₂

The metal oxides are then reduced to the corresponding metal by using a suitable reducing agent such as carbon.

1. Reduction of metal oxides with carbon: The oxides are reduced by coke in a closed furnace, which gives the metal and carbon monoxide (CO). Example: PbO + C → Pb + CO at 1400 °C.
2. Reduction of oxide ores with CO: Example: Fe₂O₃ + 3CO → 2Fe + 3CO₂ in the blast furnace.
3. Auto (self) reduction of sulphide ores: In the extraction of Cu from its sulphide ore, the ore is subjected to partial roasting in air to give its oxide: 2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂. When the supply of air is stopped and the temperature is raised, the rest of the sulphide reacts with the oxide and forms the metal and SO₂: 2Cu₂O + Cu₂S → 6Cu + SO₂
4. Reduction of ores (compounds) by more reactive metals: Examples: TiCl₄ + 2Mg → Ti + 2MgCl₂ at 850 °C; TiCl₄ + 4Na → Ti + 4NaCl at 850 °C.

Thermite process: When highly reactive metals such as sodium, calcium, aluminium, etc. are used as reducing agents, they displace metals of lower reactivity from their compounds. These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals produced are in the molten state. This type of reaction is used in the thermite process. The reaction of iron(III) oxide (Fe₂O₃) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermite reaction.

2Al + Fe₂O₃ → Al₂O₃ + 2Fe + Heat

2Al + Cr₂O₃ → Al₂O₃ + 2Cr + Heat

C) EXTRACTION OF METALS AT THE BOTTOM OF THE ACTIVITY SERIES (Ag, Hg, etc.)

Metals at the bottom of the activity series are often found in the free state. Their reactivity with other atoms is very low. The oxides of these metals can be reduced to metals by heat alone, and sometimes by displacement from their aqueous solutions.

1. When cinnabar (HgS), an ore of mercury, is heated in air, it is first converted into mercury(II) oxide (HgO), which is then reduced to mercury on further heating. 2HgS + 3O₂ → 2HgO + 2SO₂; then 2HgO → 2Hg + O₂ on heating.
2. Displacement from aqueous solutions: Ag₂S + 4CN⁻ → 2[Ag(CN)₂]⁻ + S²⁻; 2[Ag(CN)₂]⁻(aq) + Zn(s) → [Zn(CN)₄]²⁻(aq) + 2Ag(s) Here Ag₂S is dissolved in, say, KCN solution to get dicyanoargentate(I) ions. From these ions Ag is precipitated by treating with zinc dust.

PURIFICATION OF THE CRUDE METAL

The metal obtained by the reduction of the ore is usually contaminated with impurities like unchanged ore, other metals present in the ore and non-metals from the anions in the ore. For example, the blister copper obtained from its sulphide ore (copper pyrites, CuFeS₂) is impure; it contains some copper sulphide, iron and sulphur. It is purified by suitable methods including electrolysis. The process of obtaining the pure metal from the impure metal is called refining of the metal. Refining of the metal involves several types of processes. Some refining methods are given below:

1. Distillation
2. Poling
3. Liquation
4. Electrolysis, etc.

The process that has to be adopted for purification of a given metal depends on the nature of the metal and its impurities.

a) Distillation: This method is very useful for purification of low-boiling metals like zinc and mercury which contain high-boiling metals as impurities. The extracted metal in the molten state is distilled to obtain the pure metal as the distillate.

b) Poling: The molten metal is stirred with logs (poles) of green wood. The impurities are removed either as gases or they get oxidised and form scum (slag) over the surface of the molten metal. Blister copper is purified by this method. The reducing gases evolved from the wood prevent the oxidation of copper.

c) Liquation: In this method a low-melting metal like tin can be made to flow on a sloping surface to separate it from high-melting impurities.

d) Electrolytic refining: In this method, the impure metal is made to act as the anode. A strip of the same metal in pure form is used as the cathode. They are put in a suitable electrolytic bath containing a soluble salt of the same metal. The required metal gets deposited on the cathode in the pure form. The metal constituting the impurity goes into the solution or settles as anode mud.

The reactions are:
Anode: M → Mⁿ⁺ + ne⁻
Cathode: Mⁿ⁺ + ne⁻ → M (M = pure metal; n = 1, 2, 3, …)

We use this electrolytic method to refine copper. For this, impure copper is taken as the anode and pure copper strips are taken as the cathode. The electrolyte is an acidified solution of copper sulphate. As a result of electrolysis, copper in pure form is transferred from the anode to the cathode.

Anode: Cu → Cu²⁺ + 2e⁻
Cathode: Cu²⁺ + 2e⁻ → Cu

The soluble impurities go into the solution, whereas the insoluble impurities from the blister copper settle at the bottom of the anode as anode mud, which contains antimony, selenium, tellurium, silver, gold and platinum; the recovery of these elements may meet the cost of refining. Zinc may also be refined this way.

CORROSION

Most of the metals keep on reacting with the atmospheric air. This leads to the formation of a layer over the metal. In the long run, the underlying layers of the metal keep on getting lost due to conversion into oxides, sulphides, carbonates, etc. As a result, the metal gets eaten up. This process is called corrosion.

Rusting of iron: Rusting of iron is the most common form of corrosion. When iron articles like gates, grills, fencing, etc. come in contact with moisture present in air, the upper layer of iron turns into iron oxide. Iron oxide is brown-red in colour and is known as rust. This phenomenon is called rusting of iron. If rusting is not prevented in time, the whole iron article would turn into iron oxide. This is also known as corrosion of iron. Rusting of iron causes huge loss every year.

Prevention of rusting: For rusting, iron must come in contact with oxygen and water. Rusting is prevented by preventing the reaction between atmospheric moisture and the iron article. This can be done by painting, greasing, galvanization, electroplating, etc.

METHODS OF PREVENTION OF CORROSION

There are various methods of preventing corrosion and rusting of iron. Our main concern is to know the various methods to prevent the rusting of iron because iron is a strategic metal as it plays a very important role in the development of a nation. Some of the important methods of prevention of corrosion are as follows:

a) Painting: This is a common method of preventing iron from rusting. You might have observed that your parents paint the iron gate in the garden and the iron grills in your house. This painting prevents rusting by providing a coating over iron objects.

b) Oiling and greasing: Putting a layer of oil and grease on iron objects also prevents them from rusting. Iron parts of various machines and vehicles are oiled and greased to prevent rusting and to minimise friction.

c) Galvanization: In this method we put a layer of zinc metal on iron objects and this process is known as galvanization. This method is used on a large scale for making galvanized iron sheets for making boxes and for roof covering. Galvanized iron sheets are used to make drums, trunks and other iron containers. Galvanized iron sheets are also used for building roofs and manhole covers. In brief, galvanization prevents rusting in a big way.

d) Alloying: This is a very good method for improving the quality of different metals. In this method, a particular metal is mixed with another metal or non-metal in a fixed proportion to improve its quality, such as resistance towards corrosion, strength, hardness, shining and high tensile strength. For example, iron cannot be used for making utensils because it will rust, but when it is mixed with nickel and chromium it becomes stainless steel.

INTEXT QUESTIONS (PAGE 55)

Metallic oxides of zinc, magnesium and copper were heated with the following metals. In which cases will you find displacement reactions taking place?

Answer: Displacement reactions take place in the three cases marked “Displacement” in the table: magnesium with zinc oxide, and zinc and magnesium with copper oxide. (A metal displaces another metal from its oxide only if it is more reactive.)

Which metals do not corrode easily?

Answer: The more reactive a metal is, the more likely it is to be corroded. Therefore, less reactive metals are less likely to get corroded. This is why gold plating provides high resistance to corrosion.

What are alloys?

Answer: Alloys are homogeneous mixtures of two or more elements. The elements could be two metals, or a metal and a non-metal. An alloy is formed by first melting the metal and then dissolving the other elements in it. For example, steel is an alloy of iron and carbon.

EXERCISE QUESTIONS (PAGES 56 AND 57)

Which of the following pairs will give displacement reactions?

1. NaCl solution and copper metal
2. MgCl₂ solution and aluminium metal
3. FeSO₄ solution and silver metal
4. AgNO₃ solution and copper metal

Answer: (d) AgNO₃ solution and copper metal

Which of the following methods is suitable for preventing an iron frying pan from rusting?

1. Applying grease
2. Applying paint
3. Applying a coating of zinc
4. All of the above

Answer: (c) Applying a coating of zinc. (We can also apply grease and paint to prevent iron from rusting. However, in the case of an iron frying pan, grease and paint cannot be applied because when the pan is heated and washed again and again, the coating of grease and paint would get destroyed.)

An element reacts with oxygen to give a compound with a high melting point. This compound is also soluble in water. The element is likely to be

1. calcium
2. carbon
3. silicon
4. iron

Answer: (a) The element is likely to be calcium.

Food cans are coated with tin and not with zinc because

1. zinc is costlier than tin.
2. zinc has a higher melting point than tin.
3. zinc is more reactive than tin.
4. zinc is less reactive than tin.

Answer: (c) Food cans are coated with tin and not with zinc because zinc is more reactive than tin.

You are given a hammer, a battery, a bulb, wires and a switch. (a) How could you use them to distinguish between samples of metals and non-metals? (b) Assess the usefulness of these tests in distinguishing between metals and non-metals.

Answer: (a) With the hammer, we can beat the sample; if it can be beaten into thin sheets (that is, it is malleable), then it is a metal, otherwise a non-metal. Similarly, we can use the battery, bulb, wires and switch to set up a circuit with the sample. If the sample conducts electricity, then it is a metal, otherwise a non-metal.
(b) The above tests are useful in distinguishing between metals and non-metals as they are based on physical properties. No chemical reactions are involved in these tests. (Some exceptions exist, such as graphite, a non-metal, which conducts electricity.)

What are amphoteric oxides? Give two examples of amphoteric oxides.

Answer: Those oxides that behave as both acidic and basic oxides are called amphoteric oxides. Examples: aluminium oxide (Al₂O₃) and zinc oxide (ZnO).

Name two metals which will displace hydrogen from dilute acids, and two metals which will not.

Answer: Metals that are more reactive than hydrogen displace it from dilute acids, for example sodium and potassium. Metals that are less reactive than hydrogen do not displace it, for example copper and silver.

In the electrolytic refining of a metal M, what would you take as the anode, the cathode and the electrolyte?

Answer: In the electrolytic refining of a metal M:
Anode: impure metal M
Cathode: thin strip of pure metal M
Electrolyte: solution of a salt of the metal M

Pratyush took sulphur powder on a spatula and heated it. He collected the gas evolved by inverting a test tube over it, as shown in the figure below. (a) What will be the action of the gas on (i) dry litmus paper? (ii) moist litmus paper? (b) Write a balanced chemical equation for the reaction taking place.

Answer: (a) (i) There will be no action on dry litmus paper.
(ii) Since the gas is sulphur dioxide (SO₂), it turns moist blue litmus paper red, because sulphur dioxide reacts with moisture to form sulphurous acid.
(b) S(s) + O₂(g) → SO₂(g)
SO₂(g) + H₂O(l) → H₂SO₃(aq) (sulphurous acid)

State two ways to prevent the rusting of iron.

Answer: Two ways to prevent the rusting of iron are:
(i) Oiling, greasing or painting: By applying oil, grease or paint, the surface becomes waterproof and the moisture and oxygen present in the air cannot come into direct contact with iron. Hence, rusting is prevented.
(ii) Galvanisation: An iron article is coated with a layer of zinc metal, which prevents the iron from coming in contact with oxygen and moisture. Hence, rusting is prevented.

What type of oxides are formed when non-metals combine with oxygen?

Answer: Non-metals combine with oxygen to form acidic oxides (a few, such as CO and H₂O, are neutral). For example:
S(s) + O₂(g) → SO₂(g) (acidic in nature)

Give reasons: (a) Platinum, gold and silver are used to make jewellery. (b) Sodium, potassium and lithium are stored under oil. (c) Aluminium is a highly reactive metal, yet it is used to make utensils for cooking. (d) Carbonate and sulphide ores are usually converted into oxides during the process of extraction.

Answer: (a) Platinum, gold and silver are used to make jewellery because they are very lustrous. Also, they are very unreactive and do not corrode easily.
(b) Sodium, potassium and lithium are very reactive metals and react very vigorously with air as well as water. Therefore, they are kept immersed in kerosene oil in order to prevent their contact with air and moisture.
(c) Though aluminium is a highly reactive metal, it is resistant to corrosion. This is because aluminium reacts with oxygen present in air to form a thin layer of aluminium oxide. This oxide layer is very stable and prevents further reaction of aluminium with oxygen. Also, it is light in weight and a good conductor of heat. Hence, it is used to make cooking utensils.
(d) Carbonate and sulphide ores are usually converted into oxides during the process of extraction because metals can be easily extracted from their oxides rather than from their carbonates and sulphides.

You must have seen tarnished copper vessels being cleaned with lemon or tamarind juice. Explain why these sour substances are effective in cleaning the vessels.

Answer: Copper reacts with moist carbon dioxide in air to form copper carbonate and, as a result, the copper vessel loses its shiny brown surface, forming a green layer of basic copper carbonate. The citric acid present in lemon or tamarind neutralises the basic copper carbonate and dissolves the layer. That is why tarnished copper vessels are cleaned with lemon or tamarind juice to give the surface of the copper vessel its characteristic lustre.

Differentiate between metals and non-metals on the basis of their chemical properties.

Answer: Answer in the table below.

A man went door to door posing as a goldsmith. He promised to bring back the glitter of old and dull gold ornaments. An unsuspecting lady gave a set of gold bangles to him which he dipped in a particular solution. The bangles sparkled like new but their weight was reduced drastically. The lady was upset but after a futile argument the man beat a hasty retreat. Can you play the detective to find out the nature of the solution he had used?

Answer: He must have dipped the gold in a solution of aqua regia, a 3:1 mixture of concentrated HCl and concentrated HNO₃. Aqua regia is a fuming, highly corrosive liquid. It dissolves gold. After dipping the gold ornaments in aqua regia, the outer layer of gold gets dissolved and the inner shiny layer appears. That is why the weight of the gold ornament was reduced.

Give reasons why copper is used to make hot water tanks and not steel (an alloy of iron).

Answer: Copper does not react with cold water, hot water or steam. However, iron reacts with steam. If the hot water tanks are made of steel (an alloy of iron), then iron would react vigorously with the steam formed from hot water.
3Fe + 4H₂O → Fe₃O₄ + 4H₂ (iron + steam gives iron(II,III) oxide + hydrogen)
That is why copper is used to make hot water tanks, and not steel.