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

Introduction to Metals and Non-metals

Welcome, students! In the world of science, particularly chemistry, understanding the building blocks of matter is fundamental. You've already learned about elements, the purest substances from which everything else is made. But did you know these elements can be broadly classified into distinct groups based on their properties? This chapter, 'Metals and Non-metals', delves into this crucial classification. We are surrounded by elements in our daily lives, from the iron in our window grills and the copper in electrical wires to the oxygen we breathe and the carbon in our pencils. Understanding their classification helps us predict their behaviour, their reactions, and their uses.

This chapter will take you on a journey through the fascinating characteristics that distinguish metals from non-metals. We will start by exploring their physical properties—like shine, hardness, and conductivity—and then move on to their chemical properties, which dictate how they react with other substances like air, water, and acids. We will uncover the secrets behind why some metals are highly reactive while others are not, introducing the pivotal concept of the 'reactivity series'. Furthermore, we will learn about the very nature of chemical bonding, specifically how metals and non-metals interact to form ionic compounds. Finally, we will explore the practical science of metallurgy—the process of extracting pure metals from their natural ores—and discuss the common problem of corrosion and how we can prevent it. This chapter forms the bedrock of understanding inorganic chemistry and is essential for building a strong foundation in the subject.

Physical Properties of Metals and Non-metals

The most straightforward way to begin classifying elements is by observing their physical characteristics. These are properties that can be seen or measured without changing the chemical identity of the substance. Let's explore them in detail.

Physical Properties of Metals

Metals form a large group of elements with a set of common physical traits that make them easily identifiable and incredibly useful in various applications.

• Lustre: Metals, in their pure state, have a characteristic shining surface. This property is called metallic lustre. Think of the shine of gold, silver, or a freshly cut piece of iron. This lustre makes them ideal for making jewellery and decorative items.
• Hardness: Most metals are generally hard. The hardness varies from metal to metal. Iron, for instance, is very hard and strong, which is why it's used in construction. However, there are exceptions. Alkali metals like sodium (Na) and potassium (K) are so soft that they can be easily cut with a knife.
• Malleability: This is one of the most remarkable properties of metals. Malleability is the ability of a substance to be beaten into thin sheets. Gold and silver are among the most malleable metals. You might have seen extremely thin silver foil used to decorate sweets. Aluminium foil is another common example used in packaging food.
• Ductility: This is the ability of a substance to be drawn into thin wires. Most metals are ductile. Gold is the most ductile metal; it's astonishing to know that a single gram of gold can be drawn into a wire about 2 kilometres long! Copper and aluminium are also highly ductile, which is why they are used to make 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 is the best conductor, followed closely by copper. This is why cooking utensils are made of metals like copper and aluminium, and electrical wires are made of copper. The poorest conductor of heat among metals is lead.
• Sonority: Metals are sonorous, which means they produce a ringing sound when struck. This is why school bells and musical instruments like cymbals are made of metals.
• State: At room temperature, all metals are solid, with one famous exception: Mercury (Hg). Mercury is a liquid metal, which makes it useful in thermometers and barometers.
• High 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. However, gallium (Ga) and caesium (Cs) have very low melting points; they can melt on the palm of your hand!

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. However, there is an exception: Iodine (I) is a non-metal, but it has a lustrous, shiny appearance.
• Hardness: Non-metals are generally soft. A good example is sulphur or phosphorus. The exception here is diamond, which is an allotrope (a different structural form) of carbon. Diamond is the hardest natural substance known.
• Malleability and Ductility: Non-metals are brittle. This means they cannot be beaten into sheets or drawn into wires. If you hit a solid non-metal like a piece of coal (carbon) or sulphur with a hammer, it will break into pieces.
• Conductivity of Heat and Electricity: Non-metals are poor conductors of heat and electricity. They act as insulators. This is why the handles of cooking utensils and the coatings of electrical wires are made from non-metallic materials like plastic or rubber. An important exception is graphite, another allotrope of carbon, which is a good conductor of electricity and is used to make electrodes in batteries.
• Sonority: Non-metals are non-sonorous. They do not produce a ringing sound when struck.
• State: Non-metals exist in all three states at room temperature. For example, carbon and sulphur are solids, bromine (Br) is a liquid, and oxygen, nitrogen, and hydrogen are gases.
• Low Melting and Boiling Points: Non-metals generally have low melting and boiling points compared to metals (with exceptions like diamond and graphite).

Comparative Table: Physical Properties

Chemical Properties of Metals

While physical properties are useful for initial classification, chemical properties tell us how substances behave in chemical reactions. This is crucial for understanding their true nature and potential uses.

Reaction of Metals with Air (Oxygen)

What happens when you leave a metal out in the open? Almost all metals react with oxygen from the air to form metal oxides. The general reaction is:

Metal + Oxygen → Metal Oxide

However, the speed and conditions of this reaction vary greatly depending on the metal's reactivity.

• Highly Reactive Metals: Metals like potassium (K) and sodium (Na) react so vigorously with oxygen that they catch fire if kept in the open. To prevent this accidental combustion, they are stored immersed in kerosene oil.
• Less Reactive Metals: Metals like magnesium (Mg), aluminium (Al), zinc (Zn), and lead (Pb) react more slowly. They form a thin, protective layer of oxide on their surface. This layer prevents further oxidation (reaction with oxygen) and protects the metal underneath. This is why aluminium articles, despite being reactive, are resistant to corrosion.
• Even Less Reactive Metals: Iron (Fe) does not burn on heating, but iron filings can burn vigorously when sprinkled in a flame. Copper (Cu) does not burn but forms a black coating of copper(II) oxide when heated.
• Unreactive Metals: Metals like silver (Ag) and gold (Au) do not react with oxygen even at high temperatures.

Amphoteric Oxides: Most metal oxides are basic in nature, meaning they react with acids to form salt and water. However, some metal oxides, like aluminium oxide (Al₂O₃) and zinc oxide (ZnO), show both acidic and basic behaviour. Such oxides are called amphoteric oxides. They react with both acids and bases to produce salt and water.

Reaction with acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Reaction with base: Al₂O₃ + 2NaOH → 2NaAlO₂ (Sodium aluminate) + H₂O

Reaction of Metals with Water

Metals react with water to produce a metal oxide or a metal hydroxide, and hydrogen gas is released. The intensity of the reaction depends on the metal's reactivity.

Metal + Water → Metal hydroxide + Hydrogen (for reactive metals)

Metal + Steam → Metal oxide + Hydrogen (for less reactive metals)

• Reaction with Cold Water: Potassium (K) and sodium (Na) react violently with cold water. The reaction is so exothermic (releases so much heat) that the evolved hydrogen gas immediately catches fire. Calcium (Ca) reacts less violently, and the heat produced is not enough for the hydrogen to catch fire. Calcium starts floating because the bubbles of hydrogen gas stick to its surface.
• 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 for the same reason as calcium.
• Reaction with Steam: Metals like aluminium (Al), zinc (Zn), and iron (Fe) do not react with cold or hot water but react with steam to form the metal oxide and hydrogen.
• No Reaction with Water: Metals like lead (Pb), copper (Cu), silver (Ag), and gold (Au) do not react with water or steam at all.

Reaction of Metals with Acids (Dilute)

You've likely seen the fizzing when a metal reacts with an acid. Metals generally react with dilute acids to form a salt and hydrogen gas.

Metal + Dilute Acid → Salt + Hydrogen gas

The rate of this reaction also depends on the reactivity of the metal. For example, magnesium and zinc react readily with dilute hydrochloric acid (HCl), while iron reacts more slowly, and copper does not react at all. This is because copper is less reactive than hydrogen and cannot displace it from the acid.

A special case is nitric acid (HNO₃). When a metal reacts with nitric acid, hydrogen gas is usually not evolved. This is because HNO₃ is a strong oxidizing agent, and it oxidizes the H₂ produced to water (H₂O) and gets reduced itself to nitrogen oxides (like 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 classic example of a displacement reaction.

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

For example, if you place an iron nail (Fe) in a blue solution of copper sulphate (CuSO₄), the iron, being more reactive than copper, will displace the copper. The blue colour of the solution fades and turns light green (due to the formation of iron sulphate), and a reddish-brown coating of copper metal is deposited on the nail.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

If you were to place a copper strip in an iron sulphate solution, no reaction would occur because copper is less reactive than iron.

The Reactivity Series

Based on these displacement reactions, scientists have arranged metals in a vertical column in order of decreasing reactivity. This list is called the reactivity series or activity series. The most reactive metal is at the top, and the least reactive is at the bottom.

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

This series is a powerful tool for predicting the outcome of displacement reactions involving metals.

How do Metals and Non-metals React?

We've seen how metals and non-metals behave, but why do they react the way they do? The answer lies in their atomic structure, specifically their electrons.

Electronic Configuration and Valence Electrons

The reactivity of an element is determined by its tendency to attain a completely filled outermost shell (a stable octet), just like the noble gases. This is achieved by losing, gaining, or sharing electrons.

• Metals: Metals typically have 1, 2, or 3 electrons in their outermost shell (valence shell). It is easier for them to lose these electrons to achieve a stable configuration. When a metal atom loses electrons, it forms a positively charged ion called a cation.
• Non-metals: Non-metals typically have 4, 5, 6, or 7 electrons in their valence shell. It is easier for them to gain electrons to complete their octet. When a non-metal atom gains electrons, it forms a negatively charged ion called an anion.

Formation of Ionic Compounds

When a metal reacts with a non-metal, the metal atom transfers its valence electrons to the non-metal atom. The metal becomes a cation (positive ion), and the non-metal becomes an anion (negative ion). These oppositely charged ions are then held together by a strong electrostatic force of attraction, forming an ionic bond. The resulting compound is called an ionic compound or electrovalent compound.

Example 1: Formation of Sodium Chloride (NaCl)

Sodium (Na) has an electronic configuration of 2, 8, 1. It has one valence electron, which it readily loses to form a sodium ion (Na⁺) with a stable configuration of 2, 8.

Chlorine (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.

Na → Na⁺ + e⁻

Cl + e⁻ → Cl⁻

The Na⁺ and Cl⁻ ions are held together by a strong ionic bond to form NaCl.

Example 2: Formation of Magnesium Chloride (MgCl₂)

Magnesium (Mg) has a configuration of 2, 8, 2. It loses its two valence electrons to form Mg²⁺.

Chlorine (Cl) needs one electron. Therefore, one magnesium atom needs to react with two chlorine atoms. Each chlorine atom takes one electron from the magnesium atom to form two Cl⁻ ions. The resulting compound is MgCl₂.

Properties of Ionic Compounds

The strong forces of attraction between ions in ionic compounds give them some characteristic properties:

• Physical Nature: They are solids and are generally hard due to the strong inter-ionic attraction. 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 because a significant amount of energy is required to break the strong electrostatic forces of attraction between the ions.
• Solubility: They are generally soluble in water but insoluble in organic solvents like kerosene and petrol.
• Conduction of Electricity: In the solid state, ions are held in fixed positions and cannot move, so ionic compounds do not conduct electricity. However, in the molten state or when dissolved in water, the ions become free to move and can conduct electricity.

Occurrence of Metals (Metallurgy)

Metals are obtained from the Earth's crust. The process of extracting metals from their ores and then refining them for use is called metallurgy.

Minerals and Ores

The elements or compounds which occur naturally in the Earth’s crust are known as minerals. Some minerals contain a very high percentage of a particular metal, and the metal can be profitably extracted from it. These minerals are called ores. Therefore, all ores are minerals, but not all minerals are ores.

Extraction of Metals

The extraction of metal from its ore involves several steps, which can be summarized as follows:

1. Concentration of Ore (Enrichment): Ores mined from the earth are usually contaminated with large amounts of impurities like soil and sand, called gangue. The first step is to remove this gangue from the ore. This process is called concentration or enrichment.

2. Extraction of Metal from Concentrated Ore: This step depends on the reactivity of the metal.

3. Refining of Metal: The metal obtained in the previous step is often impure. The final step is to remove impurities to obtain the pure metal. This is called refining.

Extracting Metals of Low Reactivity

Metals at the bottom of the activity series (like Hg, Ag, Au) are very unreactive. Their ores are often oxides or sulphides which can be reduced to metals by heating alone. For example, cinnabar (HgS) is an ore of mercury. When heated in air, it is first converted into mercuric oxide (HgO), which then reduces to mercury on further heating.

2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g)

2HgO(s) → 2Hg(l) + O₂(g)

Extracting Metals of Medium Reactivity

Metals in the middle of the activity series (like Zn, Fe, Pb) are moderately reactive. They are usually present as sulphide or carbonate ores. It is easier to obtain a metal from its oxide than from its sulphide or carbonate. So, these ores are first converted into metal oxides.

• Roasting: Sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is called roasting. (e.g., for Zinc Blende, ZnS)
• Calcination: Carbonate ores are converted into oxides by heating strongly in a limited supply of air. This process is called calcination. (e.g., for Calamine, ZnCO₃)

The metal oxides are then reduced to the corresponding metal using a suitable reducing agent, such as carbon (coke). This is done in a furnace. This process is also known as smelting.

ZnO(s) + C(s) → Zn(s) + CO(g)

Sometimes, more reactive metals like sodium or aluminium are used as reducing agents. The reaction of iron(III) oxide (Fe₂O₃) with aluminium is used to join railway tracks or cracked machine parts. This is known as the Thermit reaction, and it is highly exothermic.

Extracting Metals of High Reactivity

Metals at the top of the activity series (like Na, K, Ca, Mg, Al) are very reactive. They cannot be obtained from their compounds by heating with carbon because these metals have a higher affinity for oxygen than carbon does. These metals are extracted by electrolytic reduction. This involves passing an electric current through the molten ore (or its chloride). The metal, being positively charged, is deposited at the negatively charged electrode (cathode), while the non-metal is liberated at the positively charged electrode (anode).

Refining of Metals

The most widely used method for refining impure metals is electrolytic refining. In this process:

• The impure metal is made the anode (positive electrode).
• A thin strip of pure metal is made the cathode (negative electrode).
• A solution of a salt of the same metal is used as the electrolyte.

When current is passed, pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is then deposited on the cathode. The soluble impurities go into the solution, whereas the insoluble impurities settle down at the bottom of the anode and are known as anode mud. This method is used for refining copper, zinc, tin, nickel, silver, and gold.

Corrosion and its Prevention

You have likely noticed the reddish-brown coating on old iron objects or the green layer on copper vessels. This degradation of metals is a natural process called corrosion.

What is Corrosion?

Corrosion is the process by which metals are slowly eaten away by the action of air, moisture, or a chemical (such as an acid) on their surface. For the corrosion of iron, known as rusting, both oxygen (from air) and water (or water vapour) are necessary conditions. Rust is hydrated iron(III) oxide (Fe₂O₃.nH₂O).

• Silver articles become black after some time as they react with sulphur in the air to form a coating of silver sulphide.
• Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface, gaining a green coat of basic copper carbonate.

Prevention of Corrosion

Corrosion causes enormous damage to buildings, bridges, ships, and all objects made of metals, especially iron. Preventing corrosion is therefore of great economic importance. Some common methods include:

• Painting, Oiling, Greasing: Applying a coat of paint, oil, or grease on the surface of a metal object cuts off the contact with air and moisture, thus preventing rusting.
• Galvanization: This is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc. Zinc is more reactive than iron, so it corrodes first, protecting the iron object. This is called sacrificial protection.
• Chromium Plating and Anodising: Plating a layer of a less reactive metal like chromium or tin on iron objects also prevents corrosion. Anodising is a process of forming a thick oxide layer of aluminium, which makes it resistant to further corrosion.
• Alloying: This is one of the best methods. An alloy is a homogeneous mixture of a metal with other metals or non-metals. Alloying changes the properties of the metal, often making it more resistant to corrosion. For example, pure iron is soft and rusts easily, but when mixed with nickel and chromium, we get stainless steel, which is hard and does not rust.

Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the primary metal and then dissolving the other elements in it in definite proportions. Some common alloys include:

• Brass: An alloy of copper and zinc (Cu + Zn).
• Bronze: An alloy of copper and tin (Cu + Sn).
• Solder: An alloy of lead and tin (Pb + Sn) with a low melting point, used for welding electrical wires.
• Amalgam: An alloy where one of the metals is mercury.

Important Questions and Answers

Q1: Why is sodium kept immersed in kerosene oil?

Answer: Sodium is a highly reactive metal. It reacts vigorously and exothermically with both oxygen and moisture present in the air. This reaction can be so intense that it catches fire. To prevent this accidental fire and to cut off its contact with the atmosphere, sodium is stored immersed in kerosene oil, as it does not react with kerosene.

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

Answer:

(i) Iron with steam: Iron is a moderately reactive metal that does not react with cold or hot water but reacts with steam to form iron(II,III) oxide (a mixed oxide) and hydrogen gas. The reaction is reversible.

3Fe(s) + 4H₂O(g) ⇌ Fe₃O₄(s) + 4H₂(g)

(ii) Calcium and potassium with water:

• Calcium: Calcium reacts readily but less violently with cold water to form calcium hydroxide and hydrogen gas. The heat evolved is not sufficient for the hydrogen gas to catch fire.

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

• Potassium: Potassium is a highly reactive metal that reacts violently with cold water, producing potassium hydroxide and hydrogen gas. The reaction is highly exothermic, causing the hydrogen gas to ignite immediately.

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

Q3: Differentiate between metal and non-metal on the basis of their chemical properties.

Answer: Metals and non-metals can be differentiated based on the following chemical properties:

Q4: What are amphoteric oxides? Give two examples of amphoteric oxides.

Answer: 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 a characteristic of certain metal oxides.

Two examples are:

1. Aluminium Oxide (Al₂O₃):
   • Reaction with an acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
   • Reaction with a base: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
2. Zinc Oxide (ZnO):
   • Reaction with an acid: ZnO + 2HCl → ZnCl₂ + H₂O
   • Reaction with a base: ZnO + 2NaOH → Na₂ZnO₂ + H₂O

Chapter Summary

Here are the key takeaways from this comprehensive chapter on Metals and Non-metals:

• Elements are classified as metals and non-metals based on their physical and chemical properties.
• Metals are lustrous, malleable, ductile, sonorous, and good conductors of heat and electricity. They are solid at room temperature (except mercury).
• Non-metals are generally non-lustrous, brittle, and poor conductors. They can exist as solids, liquids, or gases.
• Metals form basic or amphoteric oxides by reacting with oxygen. Non-metals form acidic or neutral oxides.
• The reactivity series lists metals in order of their decreasing chemical reactivity, which helps predict displacement reactions.
• Metals react with non-metals by transferring electrons to form ionic compounds, which are held together by strong electrostatic forces.
• Ionic compounds are hard solids with high melting points, are soluble in water, and conduct electricity in molten or aqueous states.
• The process of extracting metals from their naturally occurring ores is called metallurgy. It involves concentration, extraction, and refining.
• The extraction method depends on the metal's position in the reactivity series: electrolytic reduction (high reactivity), reduction with carbon (medium reactivity), or heating alone (low reactivity).
• Corrosion is the gradual degradation of metals due to reactions with their environment. Rusting of iron requires both air and moisture.
• Corrosion can be prevented by painting, galvanizing, alloying, and other methods.
• An alloy is a homogeneous mixture of a metal with other metals or non-metals, often created to enhance properties like strength and corrosion resistance.