CBSE Notes Class 10  Chapter 3 - Metals and Non Metals

Understanding the properties and applications of metals and non-metals is a fundamental part of the CBSE Class 10 Science curriculum, as outlined in the textbook of NCERT Solutions. This comprehensive guide offers detailed CBSE notes on Chapter named Metals and Non-Metals, designed to help students grasp key concepts and prepare effectively for their exams.

1.0Download CBSE Notes for Class 10 Science Chapter 3: Metals and Non Metals - Free PDF!!

Hey there! Looking to ace the "Metals and Non-metals" chapter? These free CBSE Class 10 notes will give you a solid understanding of the properties and reactions of these important elements.

2.0What are Metals and Non-Metals

The periodic table broadly categorizes elements into metals and non-metals, each possessing distinct properties. Metals like aluminum and copper are known for their high thermal and electrical conductivity and are widely used in construction, electronics, and transportation. On the other hand, non-metals such as sulfur and phosphorus are insulators and play essential roles in various biological and chemical processes. Understanding whether an element is a metal or non-metal helps in determining its applications and behavior in different environments.

3.0Physical Properties of Metals and Non-Metals

4.0Chemical Properties of Metals and Non-Metals

Here are some examples of Chemical Reactions:

Reaction with Oxygen

• Metals: 2Mg(s) + O₂(g) → 2MgO(s) 
• Non-Metals: C(s) + O₂(g) → CO₂(g)

Reaction with Water

• Metals: 2Na(s) + 2H₂O(l) → 2NaOH (aq) + H₂(g)
• Non-Metals: CO₂(g) + H₂O(l) → H₂CO₃ (aq)

Reaction with Acids

• Metals: Zn(s) + 2HCl (aq) →ZnCl₂(aq) +H₂ (g)
• Non-Metals: Non-metals generally do not react with acids to produce hydrogen gas.

Reaction with Bases (mostly applicable to non-metals)

• Non-Metals: Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)

5.0Reactivity Series and its Significance

The reactivity series is a list of metals arranged in order of their reactivity from highest to lowest. It is used to predict the products of displacement reactions and the method used for extracting metals.

Here is a simplified version of the reactivity series from the most reactive to the least:

• Potassium (K)
• Barium (Ba)
• Sodium (Na)
• Calcium (Ca)
• Magnesium (Mg)
• Aluminum (Al)
• Zinc (Zn)
• Iron (Fe)
• Nickel (Ni)
• Tin (Sn)
• Lead (Pb)
• Hydrogen (H)
• Copper (Cu)
• Mercury (Hg)
• Silver (Ag)
• Gold (Au)
• Platinum (Pt)

6.0Occurrence of Metals

Metals are found in the Earth’s crust either as native metals (pure form) or more commonly as minerals within ores. Ores require processing to extract the metal.

7.0Extraction of Metals

The extraction process involves several steps to convert raw ore into purified metal, considering the metal’s position in the reactivity series.

• Enrichment of Ores: This is the process of increasing the concentration of metals in ores by removing impurities and other non-metallic elements, often through physical or chemical means such as flotation and leaching.
• Extracting Metals Low in the Activity Series: Metals low in the activity series (like copper and silver) are often found in a free state and can be extracted simply by heating their oxides with carbon or carbon monoxide.
• Extracting Metals from the Middle of the Activity Series: Metals, which are present  in the middle of the activity series (like zinc and iron) are extracted by reducing their oxides with carbon or carbon monoxide at high temperatures.
• Refining of Metals: Refining is the final step in metal extraction, where impurities are removed from the crude metal to achieve high purity. Techniques include electrolytic refining, distillation, and chemical methods.

• Corrosion: Corrosion is the deterioration of metals caused by chemical reactions with their surrounding environment, typically involving oxygen (rusting of iron). Preventative measures include galvanization, painting, and using anti-corrosive alloys.

8.0Uses of Metals and Non-Metals

Metals:

• Iron: Used in construction and manufacturing machinery.
• Copper: Used in electrical wiring due to its excellent electrical conductivity.
• Aluminum: Used in packaging, transport, and construction because of its low density.

Non-Metals:

• Oxygen: Essential for respiration and used in medical treatments.
• Carbon: Used in filters, pencils, and as a fuel in its various forms.
• Sulfur: Used in the manufacture of fertilizers and gunpowder.

9.0Detailed CBSE Class 10 Science Chapter 3 – Key Notes

Introduction

Everything around us is made up of different elements. These elements can be classified into metals or non-metals on the basis of their properties. A few elements have properties common to both metals and non-metals. These are called semi-metals or metalloids.

Physical Properties of Metals

1. Sonorous
2. Lustrous & Hard
3. Electrical conductivity
4. High Melting Point & Boiling Point
5. Malleability
6. Ductility
7. Thermal conductivity

Metals conduct electricity

Metals conduct heat very easily

Physical Properties of Non-Metals

1. Brittle
2. Non-conductor
3. Soft
4. Non-lustrous
5. Non-sonorous
6. Low Melting Point & Boiling Point
7. Low density

Some Exceptions

1. All metals except mercury are solid at room temperature. Gallium (Ga) and cesium (Cs) have very low melting points. These two metals will melt if we keep them on our palm.
2. Iodine is a non-metal but it is lustrous. Non-metals are generally bad conductor of heat and electricity. Exception: Graphite is a good conductor of electricity.
3. Alkali metals such as lithium, sodium and potassium are so soft, that they can be easily cut with a knife i.e. they have low densities and low melting points. but it is lustrous.
4. Carbon can exist in different forms. Each form is called an allotrope. Diamond is the hardest natural substance, which has very high melting and boiling point. Graphite is allotrope of carbon which is good conductor of electricity.

Chemical Properties of Metals

1. Reaction of Metals with Oxygen

Metals react with oxygen to form metal oxides which are basic in nature and some of them react with water to form alkaline solutions, which turn red litmus paper blue.

4Na + O₂ → 2Na₂O

Almost all metals combine with oxygen to form metal oxides.
Metal + Oxygen → Metal oxide

Sodium oxide is a basic oxide which reacts with water to form sodium hydroxide.
Na₂O + H₂O → 2NaOH

Nature of metallic oxide
Generally, metallic oxides are basic in nature except aluminium and zinc oxides which are amphoteric in nature. That means these oxides (Al₂O₃, ZnO) react with base as well as acid.

The basic oxide of metals react with acid to give salt.
For example,
CuO(s) + H₂SO₄(l) → CuSO₄(aq) + H₂O(l)

The order of reactivity with oxygen is:
K > Na > Mg > Al > Cu

2. Reaction of Metals with Water

Metal reacts with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve into it to form metal hydroxide. But all metals do not react with water.

Metal + Water → Metal oxide + Hydrogen gas
Metal oxide + Water → Metal hydroxide

2Na + 2H₂O → 2NaOH + H₂
      (Cold water)

Mg + 2H₂O → Mg(OH)₂ + H₂
      (Boiling water)

Zn + H₂O → ZnO + H₂
      (Steam)
3Fe + 4H₂O → Fe₃O₄ + 4H₂
      (Red hot iron / Steam)

Copper do not react with water even under strong conditions.

The reactivity order of these metals with water are:
K > Na > Ca > Mg > Al > Zn > Fe > Cu

3. Reaction of Metals with Acids

The highly reactive metals react with dilute acid to displace hydrogen from acid and give a salt.

Metal + Dilute acid → Salt + Hydrogen

Na + 2HCl(dil.) → 2NaCl(s) + H₂(g)
Mg + H₂SO₄(dil.) → MgSO₄ + H₂

The reaction between Fe and acid is slow.
Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

The order of reactivity with dilute HCl is:
Mg > Al > Zn > Fe > Cu

Note :

Hydrogen gas is not evolved when metals such as Zn, Fe, Cu and Al react with nitric acid. Because HNO₃ is strong oxidising agent. It oxidises H₂ gas to water and is itself reduced to oxides of nitrogen (NO, N₂O and NO₂).

3Fe(s) + 8HNO₃(aq) → 3Fe(NO₃)₂(aq) + 4H₂O(l) + 2NO(g)
3Cu(s) + 8HNO₃(aq) → 3Cu(NO₃)₂(aq) + 4H₂O(l) + 2NO(g)

But copper reacts with hot concentrated sulphuric acid (H₂SO₄) to produce copper sulphate, sulphur dioxide and water.

Cu(s) + 2H₂SO₄(aq) → CuSO₄(aq) + SO₂(g) + 2H₂O(l)

Only Mg and Mn reacts with very dilute HNO₃ to evolve H₂ gas.

Mg(s) + 2HNO₃(aq) → Mg(NO₃)₂(aq) + H₂(g)

Aqua Regia (Royal water):
Aqua regia is a Latin word it means "royal water". It is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3 : 1. It is a highly corrosive, fuming liquid and is used to dissolve gold and platinum.

4. Reaction of Metals with Other Metal Salts

When a more reactive metal is placed in a salt solution of less reactive metal, then the more reactive metal displaces the less reactive metal from its salt solution. This reaction is also known as Displacement Reaction.

Metal A + Salt solution of Metal B → Salt solution of Metal A + Metal B
(Only if metal A is more reactive than Metal B)

E.g.
CuSO₄(aq) + Zn(s) → ZnSO₄(aq) + Cu(s) (Zn > Cu)
2AgNO₃(aq) + Cu(s) → Cu(NO₃)₂(aq) + 2Ag(s)

Reactivity Series of Metals

The vertical arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
                  (Blue)             (Green)        (Brown)

5. Reaction of Metals with Non-Metals

• Metals are electron donors so they form cations and non-metals are electron acceptors therefore, they form anions to achieve nearest noble gas configuration and become stable.
• When a metal donates an electron and non-metal accepts the same electron, they get bound by strong electrostatic forces of attraction and form ionic bond or electrovalent bond. This can be represented by Lewis dot structure.

Electrovalent bond or ionic bond may be defined as the electrostatic force of attraction which holds the oppositely charged ions together. The number of electrons lost or gained by the atom is called its electrovalency.

Properties of Ionic Compounds

• Physical state: They are solid, relatively hard, held together by strong electrostatic force of attraction and are generally brittle.
• Solubility: As ionic compounds are polar in nature, so they dissolve in polar solvent like water and are insoluble in non-polar solvents like kerosene.
• Melting and Boiling Points: As they are bound by strong electrostatic forces of attraction a large amount of energy is needed to break these bonds.
• Conduction of electricity: Ions do not move in solid state so, ionic compounds conduct electricity in molten or in aqueous solution only.
• Colour to the flame: Most metal salts when brought into the flame, impart characteristic colour to the flame.

Metal – Colour of flame
Sodium – Yellow
Barium salts – Green
Potassium – Lilac/violet

Occurrence of Metals

The earth's crust is the major source of metals. They are present in nature in the free state (Native state) as well as in combined state.

Minerals

The elementary state of the compounds in the form of which the metals occur in nature are called minerals.

Ores

The minerals from where metals can be conveniently and profitably extracted are called ores.

For example, Copper occurs in nature in the form of several mineral like copper pyrites (CuFeS₂), copper glance (Cu₂S) and cuprite (Cu₂O). We obtain copper metal profitably from copper pyrites mineral, so it is called ore of copper.

Note: All ores are minerals but all minerals are not ores.

Extraction of Metals

The process through which a pure metal is extracted from its ores is known as extraction of metals. The series of various processes involved in the extraction of metals from their ores, followed by refining of the metal is known as metallurgy.

Various steps involved in the extraction of metals or metallurgical process:

1. Crushing and grinding of the ore.
2. Concentration of the ore or enrichment of the ore.
3. Extraction of metal from the concentrated ore.
4. Refining or purification of the impure metal.

1. Crushing and grinding of the ore

Most of the ores in nature occur as big rocks. They are broken into small pieces with the help of crushers. These pieces are then reduced to fine powder with the help of a ball mill or a stamp mill. This process is known as pulverization of the ore.

2. Enrichment of ore or concentration of ore

The ores mined from the earth's crust contain a number of impurities, such as soil, sand, etc. called gangue or matrix. The process of removal of impurities (gangue) from the ore is called enrichment of ore or concentration of ore.

3. Extraction of metal from the enriched ore

The method used for extraction of the metal from the concentrated ore depends upon the nature of metal. Based on the reactivity, the metals have been grouped into the following three categories:

(I) Metals of low reactivity.
(II) Metals of medium reactivity.
(III) Metals of high reactivity.

(I) Extraction of metals low in the activity series (Cu, Hg, Ag, Pt, Au)

These metals are either found in native state or in the form of sulphide ores. E.g., Cinnabar (HgS)

These sulphide ores can be converted to oxide ores on heating in the presence of excess of air called Roasting.

$2HgS(s)+3O_2(g)\stackrel{\text{Heat}}{\to }2HgO(s)+2S{O}_2(g)$

This oxide can be reduced to metal by further heating.

$2HgO\stackrel{\text{Heat}}{\to }2Hg(l)+O_2(g)$

2Cu₂S(s) + 3O₂(g) → 2Cu₂O(s) + 2SO₂(g)
2Cu₂O(s) + Cu₂S(s) → 6Cu(s) + SO₂(g)

The reaction in which one of the reactant (Cu₂S) carries the reduction of the product (Cu₂O) is known as autoreduction or self-reduction.

(II) Extraction of metals in the middle of the activity series (Fe, Zn, Pb, etc.)

These metals are found in the form of their oxides, sulphides and carbonates. For easy extraction, sulphide and carbonate ores are first converted into the oxide.

(a) Conversion into metal oxide

(i) Calcination:
For the conversion of carbonate ores into oxides. It is the process of heating the ore strongly in the absence of air.

E.g.,
ZnCO₃(s) → ZnO(s) + CO₂(g)
(Calamine) (Absence of air)

(ii) Roasting:
For the conversion of sulphide ores into oxides. It is the process of heating the ore strongly in the presence of excess of air.

E.g.,
2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)
(Zinc blende) (Presence of excess of air)

(b) Reduction of the metal oxide to metal

For reduction suitable reducing agents are used, like carbon, carbon monoxide, aluminium, sodium or calcium.

(i) Reduction by heating with carbon (coke)
When zinc oxide is heated with carbon, zinc metal is produced.
ZnO(s) + C(s) → Zn(s) + CO(g)

(ii) Reduction with CO
Iron is obtained from ferric oxide by heating with CO.
Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g)

(iii) Reduction by aluminium
Certain metal oxides are reduced by aluminium to metals. This method is known as aluminothermy or thermite process. These displacement reactions are highly exothermic, so metals are produced in the molten state.

For example, Chromium, manganese, vanadium metals are obtained by the reduction of their oxides with Al powder.

3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + Heat

Thermite reaction – is used to weld railway tracks or cracked machine parts.

Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat

The mixture of iron oxide and aluminium powder is called thermite.

(III) Extraction of metals high up in the activity series: (K, Ca, Na, Mg and Al)

The highly reactive metals such as K, Na, Mg have strong affinity for oxygen, so they can not be reduced with the help of carbon. Hence these metals are obtained by electrolysis of their molten or fused oxides or chlorides, this method is called electrolytic reduction.

On electrolysis, metal ions, being positive, are liberated at the cathode (negative electrode) where they gain electrons and convert in the metal atoms.

For examples,

(i) Sodium metal is obtained by electrolysis of molten sodium chloride.

NaCl(s) → Na⁺(ℓ) + Cl⁻(ℓ)

At Cathode:
Na⁺(ℓ) + e⁻ → Na(s) (Reduction)

At Anode:
Cl⁻(ℓ) → Cl(g) + e⁻ (Oxidation)

Thus, sodium metal is obtained at cathode whereas chlorine gas is obtained at the anode.

(4) Refining of impure metals

The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining of impure metals is electrolytic refining.

Process

(i) In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode.

(ii) A solution of the metal salt is used as an electrolyte. On passing the electric current through the electrolyte, the pure metal from the anode dissolves into the electrolyte.

(iii) An equivalent amount of pure metal from the electrolyte gets deposited on the cathode. The soluble impurities go into the solution, leaving the insoluble impurities which settle down at the bottom of the anode.

At Anode:
M(s) → Mⁿ⁺(aq) + ne⁻

At cathode:
Mⁿ⁺(aq) + ne⁻ → M(s)

Anode mud / anode sludge
The soluble impurities present in the impure metal pass into solution whereas insoluble impurities fall below the anode as anode mud.

E.g., Electrolytic refining of copper.

Electrolytic refining of copper. The electrolyte is a solution of acidified copper sulphate. The anode is impure copper, whereas the cathode is a strip of pure copper. On passing electric current, pure copper is deposited on the cathode.

Corrosion

The process of slowly eating up of metals due to their conversion into oxides, carbonates, sulphide, sulphates, etc. by the action of atmospheric gases and moisture is called corrosion.

In case iron is the metal involved in the chemical process, then corrosion is called rusting.

Factors which promote corrosion

(a) Position of metal in the reactivity series: Active metals placed above hydrogen are easily corroded as compared to metals which are placed below hydrogen.

(b) Presence of water vapours and gases like CO₂, SO₂, etc. in the air.

(c) Presence of salts or electrolyte in water promotes corrosion.
e.g. Rusting of iron is faster in sea water than in ordinary or distilled water.

Example of corrosion

(i) When iron is exposed to moist air for a long time, its surface acquires a brown flaky substance called rust and the process is known as rusting. Rust is a mixture of Fe₂O₃ and Fe(OH)₃.

(ii) Copper reacts with CO₂ in the air and slowly loses its shiny brown surface and acquires a green coating of basic copper carbonate in moist air.

(iii) Silver articles becomes black after sometime when exposed to air. [Due to reaction with H₂S in the air to form a black coating of silver sulphide (Ag₂S).]

(iv) Lead or stainless steel lose their lustre due to corrosion.

Prevention of Corrosion

• By painting
• Oiling and greasing
• Self prevention
• Electroplating
• Cathodic protection
• By alloying

Alloys

An alloy is a homogeneous mixture of two or more metals or non-metal. It can be prepared by first melting the metal and then dissolving the other elements (metal or non-metal) in proper proportions. The physical properties of an alloy are different from the constituent metals (from which it is made).

Some of the Common Alloys are

Steel

Iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But, if it is mixed with a small amount of carbon (about 0.05%) it becomes hard and strong.

When iron is mixed with nickel and chromium to form stainless steel which is hard and does not rust.

Amalgam

An alloy of mercury and one or more other metals is known as an amalgam. It may be solid or liquid. A solution of sodium metal in liquid mercury metal is called sodium amalgam, which is used as a reducing agent. Amalgam of silver, tin and zinc is used by dentists for filling in teeth.

Brass

Brass is an alloy of copper (Cu) and Zn. It contains 80% copper and 20% zinc. It is more malleable and more stronger than pure copper. Brass is used for making cooking utensils, condenser sheets, pipe, screws, bolts, wire, scientific instruments, ornaments, etc.

Bronze

It is also the alloy of copper. It contain 90% of copper and 10% tin. It is highly resistant to corrosion and used for making utensils, statues, coins, hardware, etc.

Solder

It is an alloy of lead (50%) and tin (50%). It is used for soldering (or welding) electrical wires together as it melts at a low temperature.

Alloys of Gold

The purity of gold is expressed is terms of ‘carats’. Pure gold is known as 24 carats gold. It is very soft due to which, it is not suitable for making jewellery. It is alloyed with either silver or copper to make it hard and more suitable for making ornaments. In India, gold ornaments are usually made of 22 carats gold. It is an alloy of gold with silver or copper.

24 carat gold is an element. It contains only gold atoms. 14-carat and 18-carat gold are alloys. They contain a mixture of different atoms.

10.0Solved Examples

Example 1.What happens when zinc granules are treated with dilute hydrochloric acid? Write the chemical equation for the reaction and describe the test for the gas evolved.

Solution: When zinc granules react with dilute hydrochloric acid, hydrogen gas is evolved, and zinc chloride is formed.

Reaction: Zn + 2HCl → ZnCl₂ + H₂​

The gas that evolved is hydrogen. To test for hydrogen gas, bring a burning splinter near the mouth of the test tube. If hydrogen gas is present, it will burn with a pop sound.

Example 2.Why do non-metals not conduct electricity? Name one non-metal that conducts electricity and explain why it is an exception.

Solution: Non-metals do not conduct electricity because they lack free electrons or ions that can move to carry an electric charge. However, graphite is an exception. It is a form of carbon (a non-metal) that conducts electricity because it has delocalized electrons within its structure that are free to move, allowing it to conduct electricity.

Example 3.What happens when sodium reacts with water? Write the balanced chemical equation and mention the type of reaction.

Solution: Sodium reacts vigorously with water, producing sodium hydroxide (a strong base) and hydrogen gas. The reaction is highly exothermic, and the hydrogen gas evolved may catch fire.

Reaction: 2Na + 2H₂O → 2NaOH + H₂

This is a displacement reaction because sodium displaces hydrogen from water.

11.0Benefits of CBSE Notes Class 10 Science Chapter 3: Metals and Non-Metals

• Clarity on Chemical Reactions: Notes often present chemical reactions in a clear and balanced format, which is essential for understanding how metals and non-metals interact.
• Understanding Differences: Notes effectively highlight the contrasting properties and behaviours of metals and non-metals, making it easier to differentiate between them.
• Convenient Reference: You can easily refer to your CBSE notes whenever you have doubts or need to recall specific information, making learning a continuous process.
• Personalised Learning: When you create your own notes, the process itself aids in understanding and allows you to tailor the information to your learning style.