Occurence of Metals Class 10

Occurrence of Metals

1. The Earth's crust is the primary source of metals. Sea water also contains salts of metals such as sodium chloride and magnesium chloride.
2. Most metals are highly reactive and do not occur as free elements in nature; instead, they are found in the form of compounds with other elements. These compounds include oxides, carbonates, sulphides, and chlorides.
3. Metals found in their native state, i.e., as free elements, are referred to as occurring in the "native state." Examples of such metals include copper, silver, gold, and platinum.
4. Some less reactive metals like copper and silver can be found in both the native state and the combined state (as compounds).

Reactivity Series and Occurrence

1. Metals high up in the reactivity series, such as potassium, sodium, calcium, magnesium, and aluminium, are extremely reactive and are never found in nature as free elements. They are always found in combined states.
2. Metals in the middle of the reactivity series, like zinc, iron, and lead, are moderately reactive and are also found in the combined state.
3. Metals placed above copper in the reactivity series are typically found in nature only as compounds.
4. The least reactive or unreactive metals, such as copper, silver, gold, and platinum, are found in their free state as metals. However, copper and silver are mainly found in the combined state as their sulphides or oxides.

Minerals and Ores

1. Minerals are natural materials in which metals or their compounds are found in the Earth's crust.
2. Some minerals contain a high percentage of metal, while others may contain only a small percentage.
3. The presence of impurities in some minerals can make the extraction of metals difficult.
4. Ores are specific types of minerals that contain a good percentage of metal and have no objectionable impurities. Ores are economically and conveniently used to extract metals.
5. Not all minerals are ores, but all ores are minerals.

Common Ores

Some examples of common ores and the metals they contain are:

1. Sodium: Rock salt (Sodium chloride - NaCl)
2. Aluminium: Bauxite (Aluminium oxide - Al₂O₃.2H₂O)
3. Manganese: Pyrolusite (Manganese dioxide - MnO₂)
4. Zinc: Calamine (Zinc carbonate - ZnCO₃) and Zinc blende (Zinc sulphide - ZnS)
5. Iron: Haematite (Iron (III) oxide - Fe₂O₃)
6. Copper: Cuprite (Copper (I) oxide - Cu₂O) and Copper glance (Copper (I) sulphide - Cu₂S)
7. Mercury: Cinnabar (Mercury (II) sulphide - HgS)

Extraction of Metals

Extraction of metals refers to the process of obtaining pure metals from their naturally occurring compounds, which are usually found in the form of ores. The extraction process can vary depending on the reactivity of the metal, the type of ore, and the desired purity of the final metal product.

General Steps Involved in the Extraction of Metals:

1. Mining and Crushing: The first step in the extraction process is to mine the ore from the Earth's crust. Ores are typically rocks or minerals that contain a high concentration of the desired metal along with impurities, known as gangue. Once mined, the ore is often crushed into smaller pieces to facilitate further processing.
2. Concentration of Ore: Ores are rarely pure, and they usually contain a mixture of the desired metal compound and other impurities. The concentration of ore involves removing as much of the impurities as possible to obtain a concentrated form of the metal compound. Several methods can be used for concentration, including gravity separation, froth flotation, magnetic separation, and leaching.
3. Conversion of Concentrated Ore into Metal: This step involves reducing the metal compound in the concentrated ore to obtain the pure metal. The method of reduction depends on the reactivity of the metal.

I. Extraction of Highly Reactive Metals

Metals like potassium, sodium, calcium, magnesium, and aluminium are highly reactive and cannot be reduced by carbon. They are typically extracted through the process of electrolysis, where electric current is passed through a molten metal compound (usually a chloride or oxide), causing the metal to be deposited at the cathode (negative electrode) while the anode (positive electrode) reaction generates oxygen or chlorine gas.

1. Extraction of Sodium (Na)

1. Source: Sodium is commonly found in nature as sodium chloride (NaCl), also known as "rock salt."
2. Concentration: The mined sodium chloride ore is first crushed and purified to remove impurities. This process yields a concentrated form of sodium chloride.
3. Electrolysis: The concentrated sodium chloride is then melted (becomes molten) and subjected to electrolysis. In this process, electric current is passed through the molten sodium chloride.
4. Cathode Reaction: At the cathode (negative electrode), sodium ions (Na⁺) from the molten salt are attracted. These sodium ions gain electrons and are reduced to form sodium atoms (sodium metal).
   2Na⁺ + 2e^– → 2Na
5. Anode Reaction: At the anode (positive electrode), chloride ions (Cl^–) from the molten salt are attracted. These chloride ions lose electrons and are oxidised to form chlorine gas (Cl₂).
   2Cl^– → Cl₂ + 2e^–
6. Result: Sodium metal is produced at the cathode, while chlorine gas is produced at the anode.
7. Final Product: The sodium metal obtained is highly reactive and must be handled carefully as it reacts vigorously with water and air.

2. Extraction of Aluminum (Al)

1. Source: Aluminium is primarily extracted from bauxite, an ore that contains aluminium oxide (Al₂O₃).
2. Concentration: Bauxite ore is first crushed and purified to remove impurities. This process yields a concentrated form of aluminium oxide.
3. Conversion to Aluminum Chloride: Aluminium oxide is then mixed with molten cryolite (Na₃AlF₆) to lower its melting point. This mixture forms a substance known as aluminium chloride (AlCl₃).
4. Electrolysis: The aluminium chloride is subjected to electrolysis in a specialised cell called the Hall-Héroult cell.
5. Cathode Reaction: At the cathode (negative electrode), aluminium ions (Al³⁺) from the molten aluminium chloride are attracted. These aluminium ions gain electrons and are reduced to form aluminium atoms (aluminium metal).
   2Al³⁺ + 6e^– → 2Al
6. Anode Reaction: At the anode (positive electrode), oxygen ions (O^2-) from the molten aluminium chloride are attracted. These oxygen ions lose electrons and are oxidised to form oxygen gas (O₂).
   6O^2- → 3O₂ + 12e^–
7. Result: Aluminium metal is produced at the cathode, while oxygen gas is produced at the anode.
8. Final Product: The aluminium metal obtained is used for various industrial applications, especially in the aerospace and automotive industries.

II. Extraction of Moderately Reactive Metals

Extraction of moderately reactive metals involves converting their ores into metal oxides and then reducing these oxides to obtain the free metal. Moderately reactive metals include zinc, iron, tin, lead, and others that are in the middle of the reactivity series.

Steps involved in the extraction process for moderately reactive metals:

1. Concentration of Ore

1. After mining, the ore contains the metal in the form of compounds, often mixed with impurities like rocky materials, earthy particles, and other non-metallic substances known as gangue.
2. The first step is to concentrate the ore by removing the impurities. This is typically done through various physical and chemical methods, depending on the nature of the ore and the impurities present.

2. Conversion of Ores into Metal Oxides

Before the metal can be extracted, the ore is often converted into a metal oxide. This step is crucial and depends on the type of ore being processed.

Calcination:

1. If the ore is a carbonate ore, it is subjected to calcination.
2. Calcination involves heating the carbonate ore strongly in the absence of air (or in a limited supply of air).
3. The carbonate ore decomposes to form the metal oxide and carbon dioxide is released as a byproduct.
4. Example: Zinc carbonate (ZnCO₃) is calcined to produce zinc oxide (ZnO).
   ZnCO₃ (s) + Heat → ZnO (s) + CO₂ (g)

Roasting:

1. If the ore is a sulphide ore, it is subjected to roasting.
2. Roasting involves strongly heating the sulphide ore in the presence of air.
3. The sulphide ore decomposes to form the metal oxide and sulphur dioxide as a byproduct.
4. Example: Zinc sulphide (ZnS) is roasted to produce zinc oxide (ZnO).
   2ZnS (s) + 3O₂ (g) → 2ZnO (s) + 2SO₂ (g)

3. Reduction of Metal Oxide

Once the metal oxide is obtained, it is reduced to obtain the free metal. The choice of reducing agent depends on the reactivity of the metal.

Reduction with Carbon:

1. The oxides of moderately reactive metals like zinc, iron, tin, lead, and copper are often reduced using carbon as the reducing agent.
2. This process involves mixing the metal oxide with carbon (usually in the form of coke) and heating the mixture in a furnace.
3. Carbon reduces the metal oxide to produce the free metal and carbon monoxide as a byproduct.
4. Example: Reduction of zinc oxide with carbon.
   ZnO (s) + C (s) → Zn (s) + CO (g)

Reduction with Aluminum:

1. For more reactive metals like manganese and chromium, aluminium can be used as a reducing agent.
2. Aluminium is used when the metal oxide is of higher reactivity than zinc and cannot be efficiently reduced by carbon.
3. Aluminium powder reduces the metal oxide to metal, while aluminium itself is oxidised to aluminium oxide.
4. Example: Extraction of manganese metal using aluminium as the reducing agent.
   3MnO₂ (s) + 4Al (s) → 3Mn (l) + 2Al₂O₃ (s) + Heat

Thermite Reaction:

1. The reduction of metal oxides by aluminium, especially when it results in the production of molten metal, is known as a thermite reaction.
2. Thermite reactions are highly exothermic and release a significant amount of heat.
3. This property is used in thermite welding, where broken iron pieces are joined together using the intense heat generated during the thermite reaction.

III. Extraction of Less Reactive Metals

The extraction of less reactive metals, such as copper and mercury, is relatively simple compared to highly reactive metals. These metals are typically found in nature as compounds or ores, and their extraction involves specific processes.

Steps involved in the extraction of less reactive metals are:

1. Extraction of Mercury:

Mercury is a relatively unreactive metal that can be found in its ore known as cinnabar, primarily mercury(II) sulphide (HgS).

The extraction of mercury from cinnabar ore involves two main steps:

1. Roasting: The concentrated cinnabar ore is heated in the presence of air, a process known as roasting. During roasting, mercury(II) sulphide (HgS) is converted into mercury(II) oxide (HgO) and sulphur dioxide (SO₂).
   2HgS (s) + 3O₂ (g) → 2HgO (s) + 2SO₂ (g)
2. Reduction: The mercury(II) oxide obtained from roasting is further heated, typically to around 300°C, in the absence of air. This causes the mercury oxide to decompose, resulting in the formation of liquid mercury.
   2HgO (s) → 2Hg (l) + O₂ (g)

Mercury metal, which is in a liquid state at room temperature, can be collected.

2. Extraction of Copper:

Copper is another less reactive metal that can be obtained from its ore, often referred to as copper glance or copper(I) sulphide (Cu₂S).

The extraction of copper from copper glance ore also involves two primary steps:

1. Roasting: The concentrated copper(I) sulphide ore (copper glance) is subjected to roasting in the presence of air. This process converts some of the copper(I) sulphide into copper(I) oxide (Cu₂O) and sulphur dioxide (SO₂).
   2Cu₂S (s) + 3O₂ (g) → 2Cu₂O (s) + 2SO₂ (g)
2. Reduction: After a sufficient amount of copper(I) sulphide has been converted into copper(I) oxide, the supply of air is cut off. In the absence of air, copper(I) oxide reacts with the remaining copper(I) sulphide to produce copper metal and sulphur dioxide.
   2Cu₂O (s) + Cu₂S (s) → 6Cu (s) + SO₂ (g)

Copper metal is obtained in solid form.

IV. Refining (Purification) of Impure Metal

Metals obtained through various reduction processes often contain impurities, making them impure. To obtain pure metals, a process called refining is employed. Refining of metals involves the removal of these impurities and improving the quality of the metal. The choice of refining method depends on the type of metal and the nature of impurities present. One of the most widely used methods for refining impure metals is electrolytic refining.

1. Electrolytic Refining: Electrolytic refining is a process that involves purifying metals through electrolysis. Several metals, including copper, zinc, tin, lead, chromium, nickel, silver, and gold, are refined using this method.
2. Anode and Cathode: A thick block of the impure metal is made the anode, which is connected to the positive terminal of the battery. Simultaneously, a thin strip of pure metal is made the cathode, connected to the negative terminal of the battery.
3. Electrolyte: An electrolyte solution containing a water-soluble salt of the metal being refined is used in the process. Often, dilute sulfuric acid is added to the solution.
4. Process: When an electric current is passed through the setup, the impure metal from the anode dissolves into the electrolyte solution, while pure metal from the solution deposits onto the cathode. Impurities are removed from the impure metal, either dissolving into the solution or settling at the bottom of the anode as "anode mud."

Electrolytic Refining of Copper

1. Apparatus: The setup consists of an electrolytic tank containing acidified copper sulphate (CuSO₄) solution as the electrolyte. The impure copper metal is used as the anode, connected to the positive terminal of the battery, while a pure copper strip serves as the cathode, connected to the negative terminal.



2. Process: When an electric current is passed through the system, copper ions (Cu²⁺) from the copper sulphate solution move toward the cathode, where they gain electrons and become copper atoms. These copper atoms are then deposited onto the cathode, resulting in the production of pure copper metal.
   Soluble impurities in the impure copper move into the solution, while insoluble impurities settle as anode mud.
3. Explanation: In the copper sulphate solution, copper ions (Cu²⁺) and sulphate ions (SO₄^2-) are present. During electrolysis, copper ions are reduced at the cathode to form copper atoms, which get deposited on the cathode. This process effectively removes impurities from the impure copper and yields pure copper metal. The soluble impurities go into the solution, while the insoluble impurities collect below the anode as anode mud.

V. Final Processing and Alloy Formation

In some cases, the pure metal may undergo additional processing, such as casting, rolling, or machining, to shape it into useful products. Metals are often combined with other elements to form alloys, which can have improved properties like increased strength or corrosion resistance. Common alloys include steel (iron and carbon), bronze (copper and tin), and brass (copper and zinc).

Alloy

Alloys are mixtures of two or more metals, or sometimes metal and small amounts of non-metals, that combine to create materials with unique properties. Mixing metals in specific proportions while in a molten state and then cooling the mixture to room temperature creates alloys. Alloys exhibit properties different from their constituent metals, which makes them highly valuable in various applications.

Key characteristics of alloys include:

1. Increased Strength: Alloys are generally stronger than the individual metals they are composed of.
2. Enhanced Hardness: Alloys are typically harder than pure metals.
3. Improved Corrosion Resistance: Alloys often exhibit greater resistance to corrosion compared to pure metals.
4. Lower Melting Points: Alloys tend to have lower melting points than the constituent metals.
5. Reduced Electrical Conductivity: Alloys generally have lower electrical conductivity than pure metals.
6. Some common examples of alloys include:
   → Brass: An alloy of copper and zinc, known for its malleability, strength, and golden appearance. It is used in various applications, including utensils, screws, and ornaments.
   → Bronze: An alloy of copper and tin, valued for its toughness and resistance to corrosion. It is commonly used in statues, coins, and cooking utensils.
   → Solder: An alloy of lead and tin, with a low melting point, making it ideal for soldering or welding electrical wires together.
   → Amalgam: An alloy involving mercury and one or more other metals. Dentists use amalgam for tooth fillings.
   → Gold Alloys: Gold is alloyed with small amounts of other metals, like silver or copper, to increase its hardness for making jewellery. For instance, 22-carat gold is an alloy of gold with either silver or copper, making it suitable for crafting ornaments.

Corrosion

Corrosion is a natural chemical process that occurs when metals react with their environment, particularly in the presence of moisture or gases. It leads to the gradual deterioration of metals over time. Corrosion is an electrochemical process involving the transfer of electrons and ions between the metal and its environment.

Rusting of Iron

Rusting of iron is a specific type of corrosion that occurs when iron reacts with oxygen and moisture (water vapour or liquid water) in the air. The process involves the following reactions:

1. Oxidation: Iron (Fe) at the surface loses electrons and forms iron ions (Fe²⁺).
   Fe(s) → Fe²⁺(aq) + 2e^-
2. Reduction: Oxygen (O₂) in the air gains electrons and forms hydroxide ions (OH^-).
   O₂(g) + 4e^- + 2H₂O(l) → 4OH^-(aq)
3. Formation of Rust: Iron ions (Fe²⁺) combine with hydroxide ions (OH^-) to form iron hydroxide (Fe(OH)₂), which further reacts with oxygen to form hydrated iron(III) oxide, commonly known as rust (Fe₂O₃·xH₂O).

Conditions Necessary for the Rusting of Iron: Rusting of iron requires two essential conditions.

1. Presence of Air (Oxygen): Rusting needs oxygen from the air. Iron must come into contact with oxygen molecules for the corrosion process to occur.
2. Presence of Water (Moisture): Moisture, in the form of liquid water or water vapour, is necessary for rusting to happen. Water serves as an electrolyte, facilitating the flow of ions during the electrochemical reactions involved in corrosion.

Prevention of Rusting: Various methods can be employed to prevent or minimise the rusting of iron:

1. Painting: Applying a coat of paint or other protective coatings creates a physical barrier between the iron surface and the surrounding environment, preventing contact with moisture and oxygen.
2. Greasing or Oiling: Applying grease or oil to the iron surface forms a protective layer that inhibits contact with air and moisture.
3. Galvanisation: Galvanising iron involves coating it with a layer of zinc. Zinc is more reactive than iron and corrodes preferentially, protecting the iron underneath. This is commonly used for structures like fences, pipes, and car bodies.
4. Alloying: Creating alloys like stainless steel by adding elements like chromium and nickel can make iron more resistant to rusting. Stainless steel doesn't corrode readily because of the protective chromium oxide layer that forms on its surface.
5. Anodising (for Aluminum): Aluminium can be protected from corrosion by creating a thicker layer of aluminium oxide on its surface through a process called anodising. This thicker oxide layer enhances resistance to further corrosion.

Corrosion of Aluminum

Aluminium, like iron, undergoes corrosion, but it forms a different type of protective oxide layer. When aluminium is exposed to air, it reacts with oxygen to form a thin layer of aluminium oxide (Al₂O₃) on its surface. This oxide layer acts as a barrier, preventing further corrosion of the underlying aluminium. Aluminium is highly resistant to corrosion due to the protective nature of this oxide layer. In certain cases, anodizing aluminium can create a thicker and more durable oxide layer to enhance its corrosion resistance even further.

Allotropy

Allotropy is a fascinating phenomenon in chemistry where a single chemical element exists in multiple structural forms, known as allotropes, in the same physical state (solid, liquid, or gas). These allotropes have different arrangements of atoms or molecules, resulting in distinct physical and chemical properties. Allotropy is often observed in elements that can form different types of chemical bonds or possess diverse crystal structures.

Carbon (C)

Carbon is one of the most well-known elements that exhibit allotropy. It has several allotropes, including:

1. Diamond: In a diamond, each carbon atom forms strong covalent bonds with four other carbon atoms, creating a three-dimensional crystal lattice. This structure results in extreme hardness, transparency, and excellent thermal conductivity.
2. Graphite: In graphite, carbon atoms are arranged in hexagonal layers, with each carbon atom bonded to three others within the same layer. Weak van der Waals forces hold the layers together, making graphite slippery and a good conductor of electricity. It's commonly used as a lubricant and in pencils.
3. Graphene: Graphene is a single layer of carbon atoms arranged in a hexagonal lattice. It has exceptional electronic, thermal, and mechanical properties, making it a subject of intense scientific research.

Oxygen (O)

Oxygen exhibits allotropy with two primary allotropes:

1. Dioxygen (O₂): This is the most common form of oxygen, consisting of diatomic molecules (O₂) held together by a double covalent bond. Dioxygen is essential for respiration and combustion.
2. Ozone (O₃): Ozone is a triatomic molecule composed of three oxygen atoms. It forms the ozone layer in the Earth's atmosphere and plays a crucial role in protecting life from harmful ultraviolet (UV) radiation.
3. Phosphorus (P): Phosphorus has several allotropes, including white phosphorus, red phosphorus, and black phosphorus. Each allotrope has a different crystal structure and properties. White phosphorus is highly reactive and pyrophoric (ignites spontaneously in air), while red and black phosphorus are less reactive and have different applications.
4. Sulphur (S): Sulphur also has multiple allotropes, with the most common ones being rhombic sulphur (α-sulphur) and monoclinic sulphur (β-sulphur). These allotropes differ in their crystalline structures and properties. For example, rhombic sulphur is a yellow solid, while monoclinic sulphur is a red or orange solid.
5. Tin (Sn): Tin exhibits allotropy with two primary allotropes: white tin (alpha tin) and grey tin (beta tin). White tin is stable at lower temperatures and has a metallic lustre, while grey tin is less stable and becomes brittle and powdery, a phenomenon known as "tin pest."

Frequently Asked Questions

1. In what forms do metals occur in nature, and why are some metals found in pure form while others are found as compounds?

Metals occur either in their native state (pure form) or as compounds (ores) in nature. Unreactive metals like gold and platinum occur in their native state because they do not easily react with other elements. Reactive metals like iron, aluminium, and sodium are found as compounds (e.g., oxides, sulfides) because they readily react with oxygen, water, or other elements.

2. Why is electrolysis used to extract highly reactive metals like aluminium?

Highly reactive metals, such as aluminium, are extracted using electrolysis because they cannot be reduced by conventional chemical reactions, such as heating with carbon. These metals form strong bonds with oxygen or other elements in their ores, requiring a large amount of energy to break the bonds, which is supplied by the process of electrolysis.

3. How does the atomic structure of diamond differ from that of graphite, even though both are allotropes of carbon?

In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral structure, creating a rigid, three-dimensional lattice that gives diamond its hardness. In graphite, each carbon atom is bonded to three others in flat, hexagonal layers, with weak forces (Van der Waals forces) between the layers, allowing them to slide over one another, making graphite soft and slippery.

4. How does the allotropy of phosphorus affect its chemical reactivity and industrial applications?

White phosphorus is highly reactive, toxic, and flammable, making it dangerous to handle and store. Red phosphorus, in contrast, is more stable and less reactive, which is why it is used in safety matches. The different allotropic forms of phosphorus have distinct reactivities, affecting how they are used in industry.

5. Explain how carbon reduction works in the extraction of metals from their ores. Provide an example.

In carbon reduction, the ore is heated with carbon (usually in the form of coke), which acts as a reducing agent. Carbon removes the oxygen from metal oxides, forming carbon dioxide and leaving behind the pure metal. For example, in the extraction of iron from its ore (haematite).