Metals and Non-metals

Metals are elements that are generally hard, lustrous, malleable, ductile, and good conductors of heat and electricity. They have a tendency to lose electrons in chemical reactions and form cations. Examples of metals include copper, iron, gold, silver, aluminum, and zinc.

Non-metals, on the other hand, are elements that are generally not lustrous, not malleable, not ductile, and poor conductors of heat and electricity. They have a tendency to gain electrons in chemical reactions and form anions. Examples of non-metals include oxygen, nitrogen, carbon, sulfur, and chlorine.

The division of elements into metals and non-metals is based on their physical and chemical properties. However, there are some elements that have properties of both metals and non-metals, and they are called metalloids. Examples of metalloids include silicon, germanium, arsenic, and boron.

The properties of metals and non-metals are very different from each other and are used in a wide variety of applications. For example, metals are commonly used in construction, electrical wiring, and machinery, while non-metals are used in the production of fertilizers, plastics, and medicines.

Metals have several physical properties that distinguish them from non-metals. Some of the most notable physical properties of metals include:

1. Luster: Most metals have a shiny, metallic luster that reflects light and makes them appear polished and reflective.
2. Ductility: Metals can be drawn into thin wires without breaking, a property known as ductility.
3. Malleability: Metals can be hammered into thin sheets without breaking, a property known as malleability.
4. Conductivity: Metals are good conductors of heat and electricity, meaning they allow heat and electricity to flow through them easily.
5. Density: Metals are generally dense, meaning they have a high mass per unit volume.
6. Melting and boiling points: Metals typically have high melting and boiling points, meaning they require a lot of heat to melt or boil.
7. Hardness: Metals can be relatively hard, although this can vary depending on the metal.

These physical properties make metals useful in a wide range of applications, from construction and machinery to electrical wiring and electronics.

Non-metals have several physical properties that distinguish them from metals. Some of the most notable physical properties of non-metals include:

1. Lack of luster: Non-metals do not have the shiny, reflective appearance of metals. Instead, they tend to be dull in appearance.
2. Brittle: Non-metals are often brittle, meaning they are prone to breaking or shattering when subjected to stress or force.
3. Non-conductivity: Non-metals are typically poor conductors of heat and electricity. This means they do not allow heat or electricity to flow through them easily.
4. Low density: Non-metals are generally less dense than metals, meaning they have a lower mass per unit volume.
5. Low melting and boiling points: Non-metals typically have low melting and boiling points, meaning they require less heat to melt or boil.
6. Softness: Non-metals tend to be relatively soft, although this can vary depending on the specific non-metal.

These physical properties make non-metals useful in a range of applications, from producing fertilizers and plastics to creating medicines and cosmetics. However, their poor conductivity can make them unsuitable for use in electronics and other applications where electrical conductivity is important.

Ans : 1. (i) Mercury (Hg) is a metal that is liquid at room temperature. (ii) Sodium (Na) is a metal that can be easily cut with a knife. (iii) Silver (Ag) is considered the best conductor of heat among all metals. (iv) Lead (Pb) is a metal that is a poor conductor of heat.

2. Explain the meanings of malleable and ductile

Ans: The terms “malleable” and “ductile” are used to describe the ability of a metal to be shaped or molded without breaking. Malleability refers to the ability of a metal to be hammered, rolled, or pressed into thin sheets or foils, while ductility refers to the ability of a metal to be drawn into thin wires without breaking. These properties are particularly important in industries like construction and metallurgy, where metals are often shaped and molded into specific forms for use in various applications. Malleable and ductile metals can also be easily worked into intricate shapes, making them useful in jewelry and decorative arts. Examples of malleable metals include gold and silver, while examples of ductile metals include copper and aluminum.

Metals have several chemical properties that distinguish them from non-metals. Some of the most notable chemical properties of metals include:

1. Reactivity: Metals have a tendency to react with other substances, particularly non-metals, to form compounds. The reactivity of metals varies, with some metals such as sodium and potassium being highly reactive, while others such as gold and platinum are relatively unreactive.
2. Corrosion: Many metals are prone to corrosion, which occurs when they react with oxygen and other chemicals in the environment. Corrosion can lead to the formation of rust, tarnish, or other forms of damage.
3. Oxidation: Metals can undergo oxidation, which is a chemical reaction in which they lose electrons. This can result in the formation of metal oxides or other compounds.
4. Electrochemical properties: Metals have unique electrochemical properties that make them useful in batteries and other electrochemical devices. For example, metals such as lithium and zinc are commonly used in batteries because they can undergo oxidation and reduction reactions easily.
5. Reactivity with acids: Most metals will react with acids to form salts and hydrogen gas. This property can be used to identify certain metals in a laboratory setting.
6. Alloy formation: Metals can be mixed with other metals or non-metals to form alloys, which often have unique properties that make them useful in various applications. For example, steel is an alloy of iron and carbon that is stronger and more durable than pure iron.

These chemical properties make metals useful in a wide range of applications, from construction and machinery to electronics and batteries.

When metals are burnt in air, they undergo a chemical reaction with oxygen, which is present in the air. This process is known as oxidation. The exact nature of the reaction depends on the specific metal, but in general, metals react with oxygen to form metal oxides.

For example, when magnesium (Mg) is burnt in air, it reacts with oxygen (O2) to form magnesium oxide (MgO):

2Mg + O2 -> 2MgO

Similarly, when iron (Fe) is burnt in air, it reacts with oxygen to form iron oxide (Fe2O3):

4Fe + 3O2 -> 2Fe2O3

The burning of metals in air can also produce light and heat, depending on the specific metal and the conditions of the reaction. This is why some metals, such as magnesium, are used in flares and fireworks to produce bright, colorful displays. However, the burning of certain metals, such as sodium or potassium, can be dangerous and should only be done under controlled conditions.

When metals react with water, they can undergo a variety of chemical reactions depending on the specific metal and the conditions of the reaction.

1. Some metals, such as sodium (Na) and potassium (K), are highly reactive and can react violently with water. The reaction produces hydrogen gas and a metal hydroxide:

2Na + 2H2O -> 2NaOH + H2

2K + 2H2O -> 2KOH + H2

These reactions are highly exothermic, meaning they release a significant amount of heat, and can even result in explosions.

2. Other metals, such as calcium (Ca) and magnesium (Mg), also react with water, but the reaction is less violent. The reaction produces hydrogen gas and a metal hydroxide, just like the more reactive metals:

Ca + 2H2O -> Ca(OH)2 + H2

Mg + 2H2O -> Mg(OH)2 + H2

3. Some metals, such as iron (Fe), can react with water in the presence of oxygen (in the air) to form a hydrated iron(III) oxide, commonly known as rust:

4Fe + 3O2 + 6H2O -> 4Fe(OH)3

This reaction is relatively slow and can take place over a long period of time, such as the rusting of iron in a damp environment.

In general, the reactivity of metals with water depends on factors such as the specific metal, the temperature of the water, and the concentration of any dissolved salts or other substances in the water.

When metals react with acids, they can undergo a variety of chemical reactions depending on the specific metal and the acid used.

1. One common reaction is the formation of hydrogen gas. When a metal reacts with an acid, the acid donates hydrogen ions (H+) to the metal, which results in the formation of hydrogen gas and a metal salt:

Zn + 2HCl -> ZnCl2 + H2

Mg + 2HNO3 -> Mg(NO3)2 + H2

2. Another type of reaction that can occur is the reduction of the acid. Some metals, such as copper (Cu) and silver (Ag), are less reactive than hydrogen, so when they react with an acid, they do not produce hydrogen gas. Instead, the metal reduces the acid, resulting in the formation of a metal salt and water:

Cu + 2HCl -> CuCl2 + H2O

Ag + HNO3 -> AgNO3 + H2O + NO2

3. Certain metals, such as iron (Fe) and zinc (Zn), can react with dilute acids to produce metal salts and hydrogen gas. However, if the acid is concentrated, a different reaction can occur. In concentrated acid, the metal can react with the acid to produce a metal salt and oxides of nitrogen:

Fe + 6HNO3 -> Fe(NO3)3 + 3H2O + 2NO

Zn + 2HNO3 -> Zn(NO3)2 + H2O + 2NO2

In general, the reactivity of metals with acids depends on factors such as the specific metal, the concentration and type of acid used, and the temperature and pressure of the reaction.

When metals are placed in solutions of other metal salts, they can undergo a variety of reactions depending on the specific metals and the conditions of the reaction. The reaction that occurs depends on the relative reactivity of the metals involved.

1. Displacement reaction: A more reactive metal can displace a less reactive metal from its salt solution. For example, if iron (Fe) is placed in a solution of copper(II) sulfate (CuSO4), a displacement reaction will occur:

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

The iron (Fe) is more reactive than the copper (Cu), so it displaces the copper from its sulfate solution and forms iron(II) sulfate. Copper metal is deposited on the surface of iron.

2. No reaction: If a less reactive metal is placed in a solution of a more reactive metal salt, no reaction occurs. For example, if copper (Cu) is placed in a solution of iron(II) sulfate (FeSO4), no reaction will occur because copper is less reactive than iron.

Cu(s) + FeSO4(aq) → No Reaction

3. Precipitation reaction: When two metal ions in solution react to form a solid precipitate, a precipitation reaction occurs. For example, when solutions of silver nitrate (AgNO3) and sodium chloride (NaCl) are mixed, a white precipitate of silver chloride (AgCl) forms:

AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

In this case, the silver ions (Ag+) in the silver nitrate react with the chloride ions (Cl-) in the sodium chloride to form a solid silver chloride precipitate.

Overall, the reaction of metals with solutions of other metal salts depends on the relative reactivity of the metals involved and the specific conditions of the reaction.

The reactivity series is a list of metals arranged in order of their reactivity towards chemical reactions, particularly towards oxygen and water. The reactivity series is important in predicting the products of displacement reactions, the feasibility of redox reactions, and the ease with which metals can be extracted from their ores.

Here is the reactivity series in order of decreasing reactivity:

1. Potassium (K)
2. Sodium (Na)
3. Calcium (Ca)
4. Magnesium (Mg)
5. Aluminum (Al)
6. Zinc (Zn)
7. Iron (Fe)
8. Nickel (Ni)
9. Tin (Sn)
10. Lead (Pb)
11. Hydrogen (H)
12. Copper (Cu)
13. Silver (Ag)
14. Gold (Au)
15. Platinum (Pt)

Metals at the top of the reactivity series (such as potassium and sodium) are highly reactive and readily react with oxygen and water to form oxides and hydroxides. Metals at the bottom of the reactivity series (such as gold and platinum) are less reactive and do not readily react with oxygen and water.

The reactivity series can also be used to predict the products of displacement reactions. A more reactive metal can displace a less reactive metal from its salt solution, as seen in the example given in the previous answer.

Sodium metal is highly reactive and readily reacts with oxygen and moisture in the air, which can cause it to ignite or explode. To prevent this from happening, sodium is often stored in containers filled with kerosene oil or other non-reactive liquids such as mineral oil.

Kerosene oil is used because it is non-reactive with sodium and does not contain any water, which can react with the sodium metal. The oil also provides a barrier between the sodium and the air, preventing oxygen and moisture from coming into contact with the metal. This helps to keep the sodium from reacting and potentially causing a fire or explosion.

When sodium is needed for use in a chemical reaction, it is typically removed from the kerosene oil using tongs or a spatula and then washed with anhydrous solvents like ether or hexane to remove any residual oil before it is used.

(i) Iron with steam:

Iron reacts with steam (water vapor) to form iron(II) oxide and hydrogen gas.

Fe(s) + H2O(g) → FeO(s) + H2(g)

(ii) Calcium and potassium with water:

Calcium and potassium react with water to form metal hydroxide and hydrogen gas.

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

When a dilute hydrochloric acid is added to a reactive metal such as magnesium or zinc, hydrogen gas is produced. The chemical reaction can be represented as follows:

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

When iron reacts with dilute sulfuric acid, hydrogen gas is also produced along with iron(II) sulfate. The chemical reaction can be represented as follows:

Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)

When zinc is added to a solution of iron(II) sulfate, a displacement reaction takes place. Zinc is more reactive than iron, so it displaces iron from the iron(II) sulfate solution, forming zinc sulfate and iron metal. The iron metal may appear as a solid precipitate, which can be observed as a color change in the solution from green to colorless. The zinc sulfate remains in solution and will not be visible.

The chemical equation for the reaction is:

Zn(s) + FeSO4(aq) → Fe(s) + ZnSO4(aq)

Metals and non-metals have different properties and, therefore, react differently with each other. Generally, metals tend to lose electrons to form cations, while non-metals tend to gain electrons to form anions. When metals and non-metals react, they form ionic compounds.

For example, when sodium (a metal) reacts with chlorine (a non-metal), they form an ionic compound called sodium chloride. During the reaction, sodium donates one electron to chlorine, which becomes a chloride ion, and sodium becomes a sodium ion. The chemical equation for this reaction is:

2Na(s) + Cl2(g) → 2NaCl(s)

When two non-metals react, they can form covalent compounds by sharing electrons. For example, hydrogen and oxygen are both non-metals that react to form water, which is a covalent compound. The chemical equation for this reaction is:

2H2(g) + O2(g) → 2H2O(l)

The reaction between non-metals can also involve the exchange of electrons, forming ionic compounds. For example, when chlorine gas reacts with hydrogen gas, they can form hydrogen chloride, which is an ionic compound. The chemical equation for this reaction is:

H2(g) + Cl2(g) → 2HCl(g)

The nature of the reaction between metals and non-metals and non-metals and non-metals depends on their respective properties and the conditions under which the reaction occurs.

Ionic compounds are formed by the reaction between metals and non-metals, resulting in the transfer of electrons from the metal to the non-metal to form ions. These ions then combine to form a crystal lattice structure, which is characteristic of ionic compounds. Some of the properties of ionic compounds are:

1. High melting and boiling points: Ionic compounds have high melting and boiling points due to the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice. It requires a large amount of energy to overcome these forces and separate the ions.
2. Hard and brittle: Ionic compounds are usually hard and brittle due to the rigid and ordered arrangement of ions in the crystal lattice. The strong electrostatic forces holding the ions in place cause the crystal to shatter easily when subjected to external force.
3. Solubility: Ionic compounds are often soluble in polar solvents such as water due to the strong attraction between the polar solvent molecules and the charged ions. However, they are insoluble in non-polar solvents.
4. Conductivity: Solid ionic compounds do not conduct electricity as the ions are fixed in position in the crystal lattice. However, when dissolved in water or molten state, they can conduct electricity due to the movement of the ions.
5. Crystal shape: Ionic compounds form characteristic crystal shapes due to the ordered arrangement of ions in the crystal lattice. For example, sodium chloride forms a cubic crystal lattice shape.

The properties of ionic compounds are largely determined by the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice.

Some of the properties of ionic compounds and relevant chemical reaction examples are:

1. High melting and boiling points: Ionic compounds have high melting and boiling points due to the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice. For example, when solid NaCl (sodium chloride) is heated, it melts at 801°C and boils at 1413°C due to the strong ionic bonds between Na+ and Cl- ions.
2. Hard and brittle: Ionic compounds are usually hard and brittle due to the rigid and ordered arrangement of ions in the crystal lattice. When an external force is applied, the ions of the crystal lattice shift from their positions and cause the crystal to shatter easily. For example, when an external force is applied to NaCl crystal, it shatters into small pieces.
3. Solubility: Ionic compounds are often soluble in polar solvents such as water due to the strong attraction between the polar solvent molecules and the charged ions. For example, NaCl dissolves in water as the water molecules surround the ions and separate them from the crystal lattice.

NaCl(s) → Na+(aq) + Cl-(aq)

4. Conductivity: Solid ionic compounds do not conduct electricity as the ions are fixed in position in the crystal lattice. However, when dissolved in water or molten state, they can conduct electricity due to the movement of the ions. For example, when NaCl is dissolved in water, it dissociates into Na+ and Cl- ions, which can move freely and conduct electricity.

NaCl(aq) → Na+(aq) + Cl-(aq)

5. Crystal shape: Ionic compounds form characteristic crystal shapes due to the ordered arrangement of ions in the crystal lattice. For example, NaCl forms a cubic crystal lattice shape with Na+ and Cl- ions arranged alternately.

The properties of ionic compounds are largely determined by the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice, as well as their interactions with solvents and external forces.

Metals occur naturally in the earth’s crust and can be found in various forms such as ores, minerals, and rocks. The abundance of different metals varies in different regions of the world, and some metals are more commonly found than others.

The most abundant metal in the earth’s crust is aluminum, followed by iron, calcium, sodium, and potassium. Other metals like copper, zinc, lead, and gold are also found in significant quantities.

Metals can be found in their elemental form or as compounds in ores. Ores are rocks that contain high concentrations of metals or their compounds. The process of obtaining metals from ores involves several steps, including mining, beneficiation, smelting, and refining.

Mining involves the extraction of ore from the earth’s crust. Beneficiation involves separating the ore from other rock and impurities. Smelting involves heating the ore to high temperatures to extract the metal from the ore. Refining involves purifying the metal to remove any impurities.

Metals can also be obtained from seawater, although the concentration of metals in seawater is very low. Technologies like electrolysis and solvent extraction are used to extract metals from seawater.

Overall, the occurrence of metals in nature is determined by various geological processes, and the methods for extracting metals from ores and other sources depend on the specific metal and its concentration.

Extraction of metals refers to the process of obtaining metals from their ores or other sources. Different methods are used depending on the nature of the ore and the desired metal. Here are some common methods of metal extraction:

1. Pyrometallurgy: In this process, high temperatures are used to extract metals from their ores. For example, iron is extracted from iron ore using a blast furnace, which heats the ore to high temperatures and separates the metal from the rock.
2. Hydrometallurgy: This method involves using aqueous solutions to extract metals from their ores. For example, gold can be extracted from its ores by dissolving it in a cyanide solution.
3. Electrometallurgy: In this method, metals are extracted by applying an electrical current to a solution containing the metal ions. For example, aluminum is extracted by electrolysis of molten aluminum oxide.
4. Biometallurgy: This method uses microorganisms to extract metals from their ores. For example, bacteria can be used to extract copper from low-grade ores.

After the metal has been extracted from its source, it often needs to be further purified to remove any impurities. This is typically done using chemical or physical methods, such as electrolysis or fractional distillation.

The extraction of metals is an important industrial process that allows us to obtain the materials we need for various applications, from construction and manufacturing to electronics and renewable energy.

Enrichment of ores refers to the process of increasing the concentration of a metal or mineral in an ore to make it economically viable for extraction. There are several methods used for ore enrichment, including:

1. Gravity Separation: This method is based on the difference in densities of the ore and gangue particles. The ore is crushed and subjected to gravitational forces to separate it from the gangue.
2. Magnetic Separation: This method uses magnets to separate magnetic minerals from non-magnetic ones. The ore is crushed and passed through a magnetic separator to separate the magnetic minerals.
3. Froth Flotation: In this method, the ore is mixed with water and chemicals, and air is blown into the mixture to create a froth. The froth contains the mineral particles, which can then be skimmed off.
4. Leaching: This method involves dissolving the metal or mineral from the ore using a chemical solution. The metal or mineral is then recovered from the solution.
5. Smelting: This method involves heating the ore to high temperatures to extract the metal or mineral. The metal or mineral is then further purified by other methods.

The choice of ore enrichment method depends on the nature of the ore and the desired metal or mineral. Enrichment of ores is an important step in the process of obtaining metals and minerals, as it makes it economically viable to extract them from low-grade ores.

Metals that are low in the activity series, such as copper, silver, and gold, are less reactive and can be extracted using relatively simple methods. Here are the general steps involved in extracting metals low in the activity series:

1. Mining: The ore containing the metal is extracted from the ground through mining. The ore is typically crushed and ground into fine particles.
2. Concentration: The ore is concentrated to increase the metal content using techniques such as gravity separation, magnetic separation, or froth flotation.
3. Smelting: The concentrated ore is then heated in a furnace to a high temperature to extract the metal from the ore. In the case of copper, the concentrated ore is mixed with a flux, such as limestone, and heated with carbon to produce crude copper metal.
4. Refining: The crude metal is then refined to remove any impurities, such as sulfur, arsenic, or lead, that may have been present in the ore. Refining can be done through processes such as electrolysis, which uses electricity to purify the metal.

The process of extracting metals low in the activity series is relatively simple and involves a combination of mining, concentration, smelting, and refining. The process can vary depending on the nature of the ore and the desired metal, but the general steps remain the same.

One example of a metal low in the activity series that is extracted using simple methods is copper. Here is the general process of extracting copper from its ore, along with the chemical reactions involved:

1. Mining: Copper ore is extracted from the ground using either underground or open-pit mining.
2. Concentration: The copper ore is crushed and then concentrated using froth flotation. In this process, the ore is mixed with water and chemicals, including frothing agents, collectors, and modifiers. Air is then blown into the mixture, creating froth that contains the copper mineral particles.
3. Smelting: The concentrated copper ore is heated in a furnace to produce copper matte, which is a mixture of copper sulfide and iron sulfide.

CuFeS2 + O2 → 2CuS + FeS + SO2

4. Converting: The copper matte is then converted to blister copper by blowing air through it to remove the iron sulfide and convert the copper sulfide to copper metal.

2CuS + 3O2 → 2Cu + 2SO2

5. Refining: The blister copper is refined using an electrolytic process called electrorefining. In this process, the blister copper is used as the anode, and a sheet of pure copper is used as the cathode. An electric current is passed through the solution, causing the copper ions from the anode to deposit onto the cathode.

Cu2+ (aq) + 2e- → Cu (s)

The process of extracting copper from its ore involves several chemical reactions and refining steps. However, it is still considered a relatively simple process compared to the extraction of more reactive metals.

Metals in the middle of the reactivity series, such as iron and zinc, can be extracted from their ores through reduction using carbon or carbon monoxide. Here is an example of the extraction of iron from its ore, along with the chemical reactions involved:

1. Mining: Iron ore is extracted from the ground using either underground or open-pit mining.
2. Processing: The iron ore is crushed and then concentrated using magnetic separation. In this process, the ore is mixed with water and a magnetic field is applied to separate the iron mineral particles from the non-magnetic particles.
3. Reduction: The concentrated iron ore is then reduced in a blast furnace using carbon monoxide as the reducing agent. The carbon monoxide reacts with the iron oxide in the ore, producing molten iron and carbon dioxide gas.

Fe2O3 + 3CO → 2Fe + 3CO2

4. Refining: The molten iron obtained from the blast furnace is then refined to remove impurities, such as carbon and silicon. This is typically done through the basic oxygen steelmaking (BOS) process, which involves blowing oxygen into the molten iron to burn off the impurities.

C + O2 → CO2 Si + O2 → SiO2

The process of extracting iron from its ore involves several chemical reactions and refining steps. Similarly, zinc can also be extracted using carbon or carbon monoxide reduction. The chemical reactions involved are different but follow the same basic principles.

Here is an example of the extraction of sodium from its ore, along with the chemical reactions involved:

1. Mining: Sodium chloride (rock salt) is extracted from underground mines or from salt pans near the sea.
2. Purification: The impurities, such as calcium and magnesium ions, are removed from the rock salt by fractional crystallization.
3. Electrolysis: The purified sodium chloride is then melted and electrolyzed using a Downs cell. In this process, the cathode is made of liquid mercury, which forms an amalgam with the sodium metal produced. The anode is made of graphite, which reacts with the chloride ions to form carbon dioxide gas.

2NaCl(l) → 2Na(l) + Cl2(g) 2C(s) + O2(g) → 2CO(g)

The sodium metal formed at the cathode rises to the surface and is collected, while the chlorine gas produced at the anode is removed.

The overall process of extracting sodium from its ore involves several chemical reactions and purification steps. Similarly, potassium and calcium are also extracted using electrolysis, but with different electrolytic cells and chemical reactions.

The refining of metals is the process of removing impurities from crude metals obtained from the extraction process. The refining process aims to improve the quality and purity of the metal, making it suitable for various applications. There are several methods of refining metals, some of which are described below:

1. Distillation: This method is used to refine metals with low boiling points, such as zinc and mercury. The metal is heated to its boiling point and then condensed to separate it from impurities.
2. Electrolytic refining: This method is used to refine metals such as copper, nickel, and silver. The impure metal is made the anode in an electrolytic cell and a thin sheet of pure metal is made the cathode. The electrolyte is a solution of the metal salt. As electric current passes through the cell, the impure metal dissolves from the anode and gets deposited on the cathode in a pure form.
3. Vapour phase refining: This method is used to refine metals such as titanium and zirconium. The metal is heated in a vacuum, and its vapour is collected on a cooled surface. The impurities remain behind in the residue.
4. Zone refining: This method is used to refine metals with a high melting point and low solubility in their own molten state, such as germanium, silicon, and gallium. A narrow zone of the metal is melted and moved along the rod or wire. As the impurities are less soluble than the metal, they get segregated in the molten zone, and the impurities are removed from the end of the rod or wire.
5. Cupellation: This method is used to refine precious metals such as gold and silver. The impure metal is melted with lead, which oxidizes to form a lead oxide slag that absorbs the impurities. The precious metal is left behind in a pure form.

The refining process plays a crucial role in the production of high-quality metals that meet the required specifications for various applications.

1. (i) Mineral: A naturally occurring inorganic substance that has a definite chemical composition and crystal structure. (ii) Ore: A mineral deposit from which a metal or valuable mineral can be extracted profitably. (iii) Gangue: The unwanted material present in an ore that has no commercial value.
2. Two metals that are found in nature in the free state are gold (Au) and platinum (Pt).
3. The chemical process used for obtaining a metal from its oxide is reduction. In this process, the metal oxide is heated with a reducing agent such as carbon or hydrogen, which removes the oxygen from the metal oxide and leaves behind the pure metal. The reaction can be represented as:

Metal oxide + reducing agent → Metal + by-products

For example, the reduction of iron oxide (FeO) with carbon (C) to obtain iron (Fe) can be represented as:

FeO + C → Fe + CO

Corrosion is a natural process that involves the gradual destruction or deterioration of a material, usually a metal, due to its reaction with the environment. It is an electrochemical process that involves the transfer of electrons from the metal to the environment. Corrosion can occur in various forms, such as rusting of iron or steel, tarnishing of silver, and pitting of aluminum.

Corrosion occurs when a metal is exposed to a corrosive environment such as air, water, or chemicals. The corrosion process involves the formation of metal ions and the release of electrons from the metal. The metal ions combine with the environment to form corrosion products such as rust, which is a common example of corrosion.

Corrosion can be prevented or minimized by using protective coatings, such as paints or galvanizing, or by using corrosion-resistant materials such as stainless steel or aluminum. Proper maintenance and regular inspections can also help prevent corrosion. In addition, the use of sacrificial anodes or cathodic protection can also help prevent corrosion by providing an alternate path for electron transfer, thereby protecting the metal from corrosion.

Moment of Inertia

Meaning of Delta Symbol

Acids Bases and Salts