NCERT Class 10 Science Chapter 3: Metals and Non-metals - A Complete Guide

Welcome to our comprehensive guide on Chapter 3 of the NCERT Class 10 Science syllabus, 'Metals and Non-metals'. This chapter is fundamental to understanding chemistry, as it classifies elements based on their properties and explores their real-world applications, from the iron in our buildings to the gold in jewellery. We will delve deep into the physical and chemical characteristics that distinguish metals from non-metals, understand their reactions, and explore the fascinating process of extracting metals from the Earth. This guide is designed to provide a clear, step-by-step explanation of every concept, making it an essential resource for your exam preparation.

Introduction to Metals and Non-metals

In our world, we are surrounded by various materials. The elements that make up these materials can be broadly classified into three categories: metals, non-metals, and metalloids. There are over 118 known elements, and a vast majority of them are metals. Familiar examples include iron, aluminium, copper, gold, and silver. Non-metals are fewer in number but equally important, including elements like oxygen, hydrogen, carbon, and sulfur. Metalloids, such as silicon and germanium, exhibit properties intermediate between metals and non-metals.

This chapter focuses primarily on the first two categories. We will start by examining their distinct physical properties like lustre, hardness, and conductivity. Following this, we will explore their chemical behaviours, such as how they react with oxygen, water, acids, and other substances. This study will lead us to the concept of the reactivity series, a crucial tool for predicting chemical reactions. We will then learn about the formation of chemical bonds between metals and non-metals, leading to ionic compounds. Finally, we will cover the practical science of metallurgy—the process of extracting pure metals from their natural sources (ores)—and discuss the problem of corrosion and how we can prevent it.

Physical Properties of Metals and Non-metals

The most straightforward way to differentiate between metals and non-metals is by observing their physical properties. While there are exceptions to every rule, these general characteristics hold true for most elements in each category.

Physical Properties of Metals

Metals are known for a distinct set of physical characteristics that make them incredibly useful in various industries.

• Lustre: Metals, in their pure state, have a characteristic shining surface. This property is called metallic lustre. This is why metals like gold, silver, and platinum are used to make jewellery and decorative items.
• Hardness: Most metals are hard. The hardness varies from one metal to another. For instance, iron is very hard and strong, which is why it's used in construction. Exceptions: Alkali metals like sodium (Na) and potassium (K) are so soft that they can be easily cut with a knife.
• State: Metals are solids at room temperature. Their atoms are tightly packed in a crystalline structure. Exception: Mercury (Hg) is the only metal that exists as a liquid at room temperature. Gallium (Ga) and Caesium (Cs) have very low melting points and will melt if kept on your palm.
• Malleability: This is the property of metals that allows them to be beaten into thin sheets. Gold and silver are among the most malleable metals. This is why aluminium foil is used for wrapping food and silver foil for decorating sweets.
• Ductility: This is the ability of metals to be drawn into thin wires. Gold is the most ductile metal; a single gram of gold can be drawn into a wire about 2 kilometres long. Copper and aluminium are also highly ductile and are used for making electrical wires.
• Conductivity of Heat and Electricity: Metals are excellent conductors of heat and electricity. This is because they have free electrons that can move and transfer energy. Silver (Ag) is the best conductor, followed closely by copper (Cu). This is why cooking utensils are made of metals like copper and aluminium, and electrical wires are made of copper. Exceptions: Lead (Pb) and mercury (Hg) are comparatively poor conductors of heat.
• Density: Most metals have a high density and are heavy. Exceptions: Sodium and potassium have low densities and can float on water.
• Sonority: Metals are sonorous, which means they produce a ringing sound when struck. This property is why metals are used to make bells and musical instruments like cymbals.
• Melting and Boiling Points: Metals generally have high melting and boiling points due to the strong metallic bonds between their atoms. For example, iron melts at 1538°C. Exceptions: As mentioned, gallium and caesium have very low melting points.

Physical Properties of Non-metals

Non-metals exhibit properties that are generally opposite to those of metals.

• Lustre: Non-metals do not have a shiny surface; they are typically dull. Exception: Iodine (I) is a non-metal but has a lustrous appearance.
• Hardness: Non-metals are generally soft. Exception: Diamond, which is an allotrope (a different physical form) of the non-metal carbon (C), is the hardest natural substance known.
• State: Non-metals can exist in all three states at room temperature. For example, carbon and sulfur are solids, bromine (Br) is a liquid, and oxygen, nitrogen, and chlorine are gases.
• Malleability and Ductility: Non-metals are brittle. They cannot be beaten into sheets or drawn into wires. If a force is applied, they break or shatter.
• Conductivity of Heat and Electricity: Non-metals are poor conductors of heat and electricity. They act as insulators. Exception: Graphite, another allotrope of carbon, is a good conductor of electricity and is used to make electrodes in batteries and electrolytic cells.
• Density: Non-metals generally have low densities.
• Sonority: Non-metals are not sonorous; they do not produce a ringing sound when struck.
• Melting and Boiling Points: Non-metals generally have low melting and boiling points compared to metals. Exception: Diamond and graphite have very high melting points.

Chemical Properties of Metals

The chemical properties of metals are determined by their tendency to lose electrons and form positive ions (cations). Let's explore how they react with different substances.

Reaction of Metals with Air (Oxygen)

Almost all metals combine with oxygen to form metal oxides. The general reaction is:

Metal + Oxygen → Metal Oxide

For example, when copper is heated in the air, it combines with oxygen to form copper(II) oxide, a black substance.

2Cu(s) + O₂(g) → 2CuO(s) (Copper(II) oxide)

Similarly, aluminium forms a layer of aluminium oxide:

4Al(s) + 3O₂(g) → 2Al₂O₃(s) (Aluminium oxide)

Metal oxides are generally basic in nature. This means they react with acids to form salt and water. For example, copper oxide reacts with hydrochloric acid:

CuO + 2HCl → CuCl₂ + H₂O

However, some metal oxides, such as aluminium oxide and zinc oxide, show both acidic and basic behaviour. Such oxides are known as amphoteric oxides. They react with both acids and bases.

• Reaction with acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Reaction with base: Al₂O₃ + 2NaOH → 2NaAlO₂ (Sodium aluminate) + H₂O

The reactivity with oxygen varies significantly among metals. Metals like potassium (K) and sodium (Na) react so vigorously that they catch fire if kept in the open. To prevent this, they are stored under kerosene oil. At ordinary temperatures, metals like magnesium (Mg), aluminium (Al), zinc (Zn), and lead (Pb) are covered with a thin, protective layer of oxide that prevents further oxidation. Iron (Fe) does not burn on heating but iron filings burn vigorously. Copper (Cu) does not burn but forms a black oxide layer. Silver (Ag) and gold (Au) do not react with oxygen even at high temperatures.

Reaction of Metals with Water

Metals react with water to produce a metal oxide or metal hydroxide and hydrogen gas. However, not all metals react with water, and the conditions for reaction differ based on their reactivity.

Metal + Water → Metal oxide / Metal hydroxide + Hydrogen

• Reaction with Cold Water: Highly reactive metals like potassium and sodium react violently with cold water. The reaction is so exothermic that the evolved hydrogen immediately catches fire.

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy

Calcium (Ca) reacts less violently with cold water. The heat evolved is not sufficient for the hydrogen to catch fire. Calcium starts floating because the bubbles of hydrogen gas formed stick to its surface.

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

• Reaction with Hot Water: Magnesium (Mg) does not react with cold water but reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to hydrogen bubbles.

Mg(s) + 2H₂O(l) (hot) → Mg(OH)₂(aq) + H₂(g)

• Reaction with Steam: Metals like aluminium, zinc, and iron do not react with cold or hot water but react with steam to form the metal oxide and hydrogen.

2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) (Ferric oxide) + 4H₂(g)

• No Reaction with Water: Metals such as lead, copper, silver, and gold do not react with water or steam at all.

Reaction of Metals with Acids (Dilute)

Metals generally react with dilute acids to form a salt and hydrogen gas. The reactivity depends on the metal's position in the reactivity series. Metals above hydrogen will displace hydrogen from dilute acids.

Metal + Dilute acid → Salt + Hydrogen gas

For example:

Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)

The rate of formation of hydrogen bubbles indicates the reactivity. Magnesium reacts most vigorously, followed by aluminium, zinc, and then iron. Copper, gold, and silver do not react with dilute acids like HCl or H₂SO₄.

An important exception is nitric acid (HNO₃). When a metal reacts with nitric acid, hydrogen gas is generally not evolved. This is because HNO₃ is a strong oxidising agent. It oxidises the H₂ produced to water (H₂O) and is itself reduced to any of the nitrogen oxides (N₂O, NO, NO₂). However, magnesium (Mg) and manganese (Mn) are exceptions; they react with very dilute HNO₃ to evolve H₂ gas.

Reaction of Metals with Solutions of Other Metal Salts

A more reactive metal can displace a less reactive metal from its salt solution. This is a type of displacement reaction.

Metal A + Salt Solution of B → Salt Solution of A + Metal B (if A is more reactive than B)

For example, if an iron nail is placed in a blue copper sulphate solution, the iron displaces the copper. The blue colour of the solution fades, and it turns light green due to the formation of iron(II) sulphate, while a reddish-brown coating of copper metal is deposited on the nail.

Fe(s) + CuSO₄(aq) (Blue) → FeSO₄(aq) (Light Green) + Cu(s) (Red-brown)

This type of reaction helps us build the Reactivity Series, which is a list of metals arranged in order of their decreasing reactivity.

The Reactivity Series (or Activity Series):

• K (Potassium) - Most reactive
• Na (Sodium)
• Ca (Calcium)
• Mg (Magnesium)
• Al (Aluminium)
• Zn (Zinc)
• Fe (Iron)
• Pb (Lead)
• [H] (Hydrogen) - Reference point
• Cu (Copper)
• Hg (Mercury)
• Ag (Silver)
• Au (Gold) - Least reactive

How do Metals and Non-metals React?

The reactions of elements are governed by their tendency to achieve a completely filled valence shell, which is a stable electronic configuration, similar to noble gases. Metals achieve this by losing electrons, while non-metals achieve it by gaining electrons.

Formation of Ionic Compounds

When a metal reacts with a non-metal, there is a transfer of electrons from the metal atom to the non-metal atom. This results in the formation of ions: positively charged cations (from metals) and negatively charged anions (from non-metals). The electrostatic force of attraction between these oppositely charged ions holds them together, forming an ionic compound with an ionic bond.

Let's take the example of the formation of sodium chloride (NaCl):

• A sodium atom (Na) has an electronic configuration of 2, 8, 1. It has one electron in its outermost shell. By losing this electron, it forms a sodium ion (Na⁺) with a stable configuration of 2, 8.

Na → Na⁺ + e⁻

• A chlorine atom (Cl) has an electronic configuration of 2, 8, 7. It needs one electron to complete its octet. It gains the electron lost by sodium to form a chloride ion (Cl⁻) with a stable configuration of 2, 8, 8.

Cl + e⁻ → Cl⁻

• The positively charged Na⁺ ion and the negatively charged Cl⁻ ion are attracted to each other, forming sodium chloride (NaCl).

Properties of Ionic Compounds

Ionic compounds have a unique set of properties due to the strong ionic bonds between their constituent ions.

• Physical Nature: They are hard, crystalline solids. This is because the positive and negative ions are held in a fixed, ordered arrangement called a crystal lattice by strong electrostatic forces. They are also brittle and break into pieces when pressure is applied.
• Melting and Boiling Points: Ionic compounds have very high melting and boiling points. This is because a significant amount of energy is required to overcome the strong inter-ionic forces of attraction.
• Solubility: They are generally soluble in water but insoluble in organic solvents like petrol and kerosene. Water molecules can surround the ions and weaken the ionic bond, causing the compound to dissolve.
• Conduction of Electricity: In the solid state, ionic compounds do not conduct electricity because the ions are not free to move. However, in the molten state or when dissolved in water (aqueous solution), the ions become mobile and can move to the electrodes, allowing the compound to conduct electricity.

Occurrence of Metals (Metallurgy)

Metals are sourced from the Earth's crust. Some, like gold and platinum, are so unreactive that they are found in their native or free state. Most metals, however, are found combined with other elements as compounds. The naturally occurring chemical substances in the Earth's crust containing metals are called minerals. Minerals from which metals can be extracted profitably and conveniently are called ores. Therefore, all ores are minerals, but not all minerals are ores. The unwanted impurities like soil, sand, and rock present in the ore are called gangue.

The science of extracting pure metals from their ores is called metallurgy. The process involves three main steps:

1. Concentration or Enrichment of the Ore: This is the removal of the gangue from the ore.
2. Extraction of Metal from the Concentrated Ore: This involves converting the ore into a form that can be reduced, and then reducing it to get the metal.
3. Refining or Purification of the Metal: This is the final step to obtain the metal in a very pure form.

The specific extraction technique used depends on the metal's reactivity, as indicated by its position in the activity series.

Extracting Metals Low in the Activity Series

Metals like mercury (Hg), silver (Ag), and gold (Au) are very unreactive. Their oxides can be reduced to metals by simple heating. For example, mercury is extracted from its ore, cinnabar (HgS).

• Roasting: The sulphide ore is heated strongly in the presence of air to convert it into an oxide.

2HgS(s) + 3O₂(g) --(Heat)--> 2HgO(s) + 2SO₂(g)

• Reduction: The metal oxide is then heated further to decompose it into the pure metal.

2HgO(s) --(Heat)--> 2Hg(l) + O₂(g)

Extracting Metals in the Middle of the Activity Series

Metals like zinc (Zn), iron (Fe), and lead (Pb) are moderately reactive. They are usually found as sulphide or carbonate ores. It is easier to obtain a metal from its oxide than from its sulphide or carbonate. So, the first step is to convert the ore into the metal oxide.

• Roasting: This process is used for sulphide ores. The ore is heated strongly in the presence of excess air.

2ZnS(s) + 3O₂(g) --(Heat)--> 2ZnO(s) + 2SO₂(g)

• Calcination: This process is used for carbonate ores. The ore is heated strongly in a limited supply or absence of air.

ZnCO₃(s) --(Heat)--> ZnO(s) + CO₂(g)

Once the metal oxide is formed, it is reduced to the corresponding metal using a suitable reducing agent, most commonly carbon (in the form of coke).

ZnO(s) + C(s) --(Heat)--> Zn(s) + CO(g)

Sometimes, more reactive metals are used as reducing agents. The Thermite reaction is a prime example, where aluminium is used to reduce iron(III) oxide. The reaction is highly exothermic, and the iron produced is in a molten state. This is used to join railway tracks or cracked machine parts.

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

Extracting Metals Towards the Top of the Activity Series

Metals like potassium, sodium, calcium, and aluminium are highly reactive. They have a strong affinity for oxygen and cannot be reduced from their oxides using carbon. This is because these metals themselves are stronger reducing agents than carbon. Therefore, they are extracted by electrolytic reduction (electrolysis) of their molten chlorides or oxides.

For example, in the extraction of sodium from molten sodium chloride (NaCl):

• An electric current is passed through the molten salt.
• At the cathode (negative electrode), sodium ions (Na⁺) gain electrons and are deposited as pure sodium metal.

Na⁺ + e⁻ → Na

• At the anode (positive electrode), chloride ions (Cl⁻) lose electrons and are liberated as chlorine gas.

2Cl⁻ → Cl₂(g) + 2e⁻

Refining of Metals

The metals produced by the above methods are often not completely pure. The most widely used method for refining impure metals is electrolytic refining. This is used for metals like copper, zinc, tin, nickel, silver, and gold.

In the electrolytic refining of copper:

• Anode: A thick block of impure copper metal.
• Cathode: A thin strip of pure copper metal.
• Electrolyte: A solution of an acidic salt of the metal to be refined (e.g., acidified copper sulphate solution).

When an electric current is passed through the electrolyte, pure copper from the anode dissolves into the solution as Cu²⁺ ions. An equivalent amount of pure copper from the solution is then deposited onto the cathode. The soluble impurities go into the solution, whereas the insoluble impurities settle at the bottom of the anode as anode mud, which can contain valuable metals like gold and silver.

Corrosion

Corrosion is the process of gradual degradation of metals by the action of air, moisture, or chemicals on their surface. The most common example is the rusting of iron. When iron is exposed to moist air for a long time, it acquires a coating of a brown, flaky substance called rust (hydrated iron(III) oxide, Fe₂O₃·xH₂O). Both oxygen and water are essential for rusting to occur.

Corrosion causes enormous damage to buildings, bridges, ships, and all objects made of metal, especially iron. It is a major economic concern worldwide.

Prevention of Corrosion

The rusting of iron can be prevented by cutting off its contact with oxygen and moisture. This can be done by several methods:

• Painting, Oiling, or Greasing: Applying a coat of paint, oil, or grease on the surface of the metal creates a protective barrier.
• Galvanization: This is the process of coating iron or steel with a thin layer of zinc. Zinc is more reactive than iron, so it corrodes first, protecting the iron underneath. This is an example of sacrificial protection.
• Chrome Plating and Anodising: Coating the metal with a layer of another, less reactive metal like chromium or tin. Anodising involves creating a thick, protective oxide layer on metals like aluminium.
• Alloying: This is one of the best methods. When iron is mixed with other elements, its properties change. For example, stainless steel is an alloy of iron with nickel and chromium, which is highly resistant to rust.

Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by melting the primary metal and then dissolving the other elements in it in definite proportions, followed by cooling.

Alloying is done to enhance the properties of metals. For example:

• Hardness and Strength: Pure iron is soft and stretches easily when hot. Adding a small amount of carbon makes it hard and strong, creating steel.
• Corrosion Resistance: As mentioned, stainless steel is rust-resistant.
• Lower Melting Point: Solder, an alloy of lead and tin, has a lower melting point than either of its constituents, making it ideal for welding electrical wires.
• Different Electrical Conductivity: Brass (copper and zinc) and bronze (copper and tin) are not good conductors of electricity, whereas copper is an excellent one.

An alloy where one of the metals is mercury is called an amalgam.

Important Questions and Answers

Q1: Why is sodium kept immersed in kerosene oil?

A: Sodium is a highly reactive metal belonging to the alkali metal group. It reacts vigorously and exothermically with both the oxygen and moisture present in the air. The reaction is so intense that it can catch fire. To prevent this accidental fire and to protect the metal from reacting, sodium is stored under an inert liquid like kerosene oil, which does not contain oxygen and prevents the metal's contact with air and moisture.

Q2: Write equations for the reaction of (i) iron with steam and (ii) calcium with water.

A: (i) Iron is a moderately reactive metal that does not react with cold or hot water. However, it reacts with steam (gaseous water) to form iron(II,III) oxide (ferrosoferric oxide) and hydrogen gas. The balanced chemical equation is:
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

(ii) Calcium is a reactive metal that reacts with cold water to form calcium hydroxide and hydrogen gas. The reaction is less violent than that of sodium. The balanced chemical equation is:
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Q3: What are amphoteric oxides? Give two examples of amphoteric oxides with reactions.

A: Amphoteric oxides are metal oxides that exhibit both acidic and basic properties. This means they can react with both acids and bases to form salt and water. This dual behaviour is characteristic of oxides of metals like aluminium and zinc.
Examples:
1. Aluminium Oxide (Al₂O₃)

• Reaction with an acid (HCl): Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)
• Reaction with a base (NaOH): Al₂O₃(s) + 2NaOH(aq) → 2NaAlO₂(aq) (Sodium Aluminate) + H₂O(l)

• Reaction with an acid (HCl): ZnO(s) + 2HCl(aq) → ZnCl₂(aq) + H₂O(l)
• Reaction with a base (NaOH): ZnO(s) + 2NaOH(aq) → Na₂ZnO₂(aq) (Sodium Zincate) + H₂O(l)

Q4: 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.

A: (a) We can use the given items to test for two key physical properties: malleability and electrical conductivity.

• Test for Malleability: Place the sample on a hard surface and hit it with the hammer. If the sample flattens into a sheet without breaking, it is a metal (malleable). If it shatters or breaks into pieces, it is a non-metal (brittle).
• Test for Electrical Conductivity: Set up a simple circuit using the battery, bulb, wires, and switch. Leave a gap in the circuit where the sample can be inserted. Connect the sample into the gap and close the switch. If the bulb lights up, the sample is a conductor of electricity, and therefore a metal. If the bulb does not light up, the sample is a poor conductor, and therefore a non-metal.

• The malleability test is quite reliable, although some metals like zinc can be brittle.
• The conductivity test is also very effective, but one must be aware of exceptions like graphite (a non-metal) which is a good conductor of electricity, and metals like lead and mercury which are poor conductors. Despite these exceptions, these tests provide a solid and practical basis for classification.

Chapter Summary

Here are the key takeaways from our detailed exploration of 'Metals and Non-metals':

• Elements are classified as metals and non-metals based on their properties.
• Metals are typically lustrous, malleable, ductile, sonorous, and are good conductors of heat and electricity. They are solid at room temperature (except mercury).
• Non-metals are generally non-lustrous, brittle, poor conductors of heat and electricity, and can exist as solids, liquids, or gases.
• Metals react with oxygen to form basic or amphoteric oxides. Non-metals form acidic or neutral oxides.
• The reactivity of metals with water and acids varies, and this behaviour is summarized in the reactivity series.
• A more reactive metal displaces a less reactive metal from its salt solution.
• Metals and non-metals react by transferring electrons to form ionic compounds, which are hard, have high melting points, and conduct electricity in a molten or dissolved state.
• Metallurgy is the process of extracting pure metals from their ores, involving concentration, extraction (based on reactivity), and refining.
• Corrosion is the degradation of metals, such as the rusting of iron, which can be prevented by methods like painting, galvanization, and alloying.
• An alloy is a mixture of metals (or a metal and a non-metal) created to enhance properties like strength and corrosion resistance.