Q1: Explain the process of extraction of metals from their respective ores using carbon reduction. Take the example of the extraction of iron (Fe) from its ore, hematite (Fe₂O₃). Write the balanced chemical equation for the reduction reaction and describe the role of carbon monoxide (CO) in the process.
Ans: Extraction of metals from their ores using carbon reduction is a common method known as the "blast furnace process." This process is used for extracting metals that are less reactive than carbon.
Example of Iron Extraction from Hematite:
Hematite ore (Fe₂O₃) is commonly used for the extraction of iron. In the blast furnace, iron oxide reacts with carbon monoxide (CO) to produce iron metal and carbon dioxide (CO₂).
Balanced Chemical Equation:
3Fe₂O₃ + 3CO → 2Fe₃O₄ + 3CO₂
Fe₃O₄ + 4CO → 3Fe + 4CO₂
Role of Carbon Monoxide (CO):
Carbon monoxide is a reducing agent in this process. It reacts with the iron oxide in hematite to form iron metal and carbon dioxide. The carbon monoxide reduces the iron oxide by removing oxygen from it. The iron produced collects at the bottom of the blast furnace and is tapped off as molten iron, while the slag (impurities) floats on top and is also removed.
Q2: Discuss the reactivity series of metals and its significance in predicting the behavior of metals with various substances. Provide examples of metals that can displace hydrogen from dilute acids and metals that cannot. Write the balanced chemical equation for the reaction of zinc (Zn) with hydrochloric acid (HCl).
Ans: The reactivity series of metals arranges metals in order of their reactivity with water, acids, and other metal ions. It helps predict the behavior of metals when they react with various substances.
Metals that Can Displace Hydrogen from Dilute Acids:
Metals higher in the reactivity series can displace hydrogen from dilute acids. For example, zinc (Zn), iron (Fe), and magnesium (Mg) can displace hydrogen from acids like hydrochloric acid (HCl).
Metals that Cannot Displace Hydrogen from Dilute Acids:
Metals lower in the reactivity series cannot displace hydrogen from dilute acids. For instance, copper (Cu), silver (Ag), and gold (Au) cannot displace hydrogen from hydrochloric acid.
Balanced Chemical Equation for Zinc Reaction with Hydrochloric Acid:
Zn + 2HCl → ZnCl₂ + H₂
In this reaction, zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. One molecule of zinc reacts with two molecules of hydrochloric acid to yield one molecule of zinc chloride and one molecule of hydrogen gas.
Q3: Describe the process of corrosion of metals with the help of an example. Explain how the presence of moisture and oxygen accelerates the corrosion process. Write the balanced chemical equation for the formation of rust (iron oxide) on iron (Fe).
Ans: Corrosion is the process by which metals deteriorate and become less useful due to the reaction with their environment. An example of corrosion is the rusting of iron.
Corrosion Process: Iron reacts with moisture (water) and oxygen from the air to form hydrated iron(III) oxide, commonly known as rust. The presence of moisture is crucial because it provides the necessary ions for the electrochemical reactions involved in corrosion. Oxygen is also necessary for the oxidation of iron.
Balanced Chemical Equation for Formation of Rust on Iron:
4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃
In this reaction, iron reacts with oxygen and water to produce iron(III) hydroxide, which is rust. Four atoms of iron react with three molecules of oxygen and six molecules of water to yield four molecules of iron(III) hydroxide.
Q4: Explain the term "amphoteric oxides" and provide examples. Describe the reaction of an amphoteric oxide with an acid and a base. Write the balanced chemical equation for the reaction of aluminum oxide (Al₂O₃) with sodium hydroxide (NaOH).
Ans: Amphoteric oxides are oxides that can exhibit both acidic and basic properties. They can react with both acids and bases to form salts and water.
Examples of Amphoteric Oxides: Aluminum oxide (Al₂O₃) and zinc oxide (ZnO) are examples of amphoteric oxides.
Reaction with Acid: Amphoteric oxides react with acids to form salts and water. For example, aluminum oxide reacts with hydrochloric acid to produce aluminum chloride and water.
Balanced Chemical Equation for Reaction with Acid:
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
Reaction with Base:
Amphoteric oxides also react with bases to form salts and water. For instance, aluminum oxide reacts with sodium hydroxide to produce sodium aluminate and water.
Balanced Chemical Equation for Reaction with Base:
Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄
Q5: Explain the process of extraction of highly reactive metals like sodium (Na) and potassium (K) from their ores. Describe the role of electrolysis in the extraction process and write the balanced chemical equation for the electrolytic extraction of sodium from sodium chloride (NaCl) melt.
Ans: Highly reactive metals like sodium and potassium are extracted by the process of electrolysis. This is because these metals cannot be extracted using traditional carbon reduction methods due to their high reactivity.
Extraction Process: The extraction of sodium and potassium involves the electrolysis of their molten chlorides. A mixture of sodium chloride (NaCl) and calcium chloride (CaCl₂) is heated to a high temperature to form a molten mixture. When an electric current is passed through the molten mixture, sodium and potassium ions migrate towards the cathode (negative electrode), where they are reduced to form metal atoms.
Balanced Chemical Equation for Electrolytic Extraction of Sodium:
2Na⁺ + 2e⁻ → 2Na
In this equation, sodium ions (Na⁺) are reduced by gaining two electrons (2e⁻) to form sodium atoms (Na). This occurs at the cathode during the electrolysis of sodium chloride melt. The chlorine ions (Cl⁻) migrate towards the anode (positive electrode), where they are oxidized to form chlorine gas (Cl₂).
The overall process of electrolysis allows the extraction of highly reactive metals like sodium and potassium by using electrical energy to drive the reduction of metal ions at the cathode.
Q6: Explain the term "alloy" and provide an example. Describe the advantages of using alloys over pure metals. Write the balanced chemical equation for the formation of brass, an alloy of copper (Cu) and zinc (Zn).
Ans: An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal, to create a material with improved properties compared to those of the individual components.
Example of Alloy:
Bronze is an example of an alloy, which is composed of copper and tin.
Advantages of Using Alloys over Pure Metals:
Balanced Chemical Equation for Formation of Brass:
Cu + Zn → CuZn
In this equation, copper (Cu) reacts with zinc (Zn) to form brass, an alloy composed of copper and zinc.
Q7: Explain the process of electrolytic refining of impure copper (Cu). Describe how this process helps in obtaining pure copper and write the balanced chemical equation for the electrolytic refining of impure copper.
Ans: Electrolytic refining is a process used to purify metals obtained from ore or other sources. In the case of impure copper, this process involves the use of electrolysis to remove impurities and obtain pure copper.
Process of Electrolytic Refining: Impure copper is used as the anode, while a thin sheet of pure copper serves as the cathode. Both are immersed in an electrolyte solution containing copper sulfate (CuSO₄) and sulfuric acid (H₂SO₄). When electric current passes through the electrolyte, copper from the anode dissolves into the solution as copper ions (Cu²⁺) and moves towards the cathode. At the cathode, copper ions are reduced and deposited as pure copper.
Balanced Chemical Equation for Electrolytic Refining of Copper:
The overall process ensures that impurities from the anode settle at the bottom as anode mud or dissolve into the electrolyte. The pure copper deposited at the cathode is collected, resulting in the purification of copper.
Q8: Discuss the importance of the "activity series" of metals in predicting displacement reactions. Provide an example of a displacement reaction involving metals. Write the balanced chemical equation for the reaction between zinc (Zn) and copper sulfate (CuSO₄) solution.
Ans: The activity series of metals is a list that arranges metals in order of their reactivity. It is a useful tool for predicting displacement reactions, where a more reactive metal displaces a less reactive metal from its compound.
Importance of Activity Series: The activity series helps determine whether a given metal can displace another metal from its compound during a reaction. A metal higher in the activity series will displace a metal lower in the series from its compound.
Example of Displacement Reaction: Zinc (Zn) is more reactive than copper (Cu). When a piece of zinc is placed in a copper sulfate (CuSO₄) solution, a displacement reaction occurs where zinc displaces copper from copper sulfate.
Balanced Chemical Equation for Zinc and Copper Sulfate Reaction:
Zn + CuSO₄ → ZnSO₄ + Cu
In this reaction, zinc reacts with copper sulfate to form zinc sulfate and copper metal. The zinc displaces copper from its compound, leading to the change of color in the solution from blue to colorless, and the formation of reddish-brown copper metal.
Q9: Explain the process of roasting as a method for the extraction of metals from their sulfide ores. Provide an example of a metal extracted by this method. Write the balanced chemical equation for the roasting of zinc sulfide (ZnS).
Ans: Roasting is a method used to convert metal sulfide ores into their corresponding oxides. It involves heating the ore in the presence of excess oxygen to remove sulfur and obtain the metal oxide.
Process of Roasting: In roasting, metal sulfide ore is heated in the presence of oxygen to produce metal oxide and sulfur dioxide gas. The metal oxide is then further processed to extract the metal.
Example of Roasting: Zinc is often extracted from its sulfide ore, zinc sulfide (ZnS), through the roasting process. Roasting converts zinc sulfide into zinc oxide (ZnO).
Balanced Chemical Equation for Roasting of Zinc Sulfide:
2ZnS + 3O₂ → 2ZnO + 2SO₂
In this reaction, zinc sulfide reacts with oxygen to produce zinc oxide and sulfur dioxide gas. The sulfur dioxide gas is a byproduct of the roasting process and is often captured for further industrial use.
Q10: Discuss the properties of non-metals that make them suitable for use as insulators. Provide examples of materials used in electrical cables and wires that utilize the insulating properties of non-metals. Write the chemical formula for sulfur dioxide (SO₂) and nitrogen dioxide (NO₂), which are common non-metal compounds.
Ans: Non-metals possess properties that make them excellent insulators. They have high resistance to the flow of electricity and poor thermal conductivity, which make them suitable for use in insulating materials.
Properties of Non-Metals as Insulators:
Examples of Insulating Materials in Electrical Cables: Materials like rubber and plastic, which are composed of non-metals, are commonly used as insulating coatings for electrical cables and wires. These materials prevent the flow of electric current and minimize the risk of electric shock.
Chemical Formulas for Non-Metal Compounds: