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Unit Test (Solutions): Metals & Non-metals | Science Class 10 PDF Download

Time: 1 hour

M.M. 30

Attempt all questions.

  • Question numbers 1 to 5 carry 1 mark each.
  • Question numbers 6 to 8 carry 2 marks each.
  • Question numbers  9 to 11 carry 3 marks each.
  • Question number 12 & 13 carry 5 marks each

Q1: Which of the following elements is a metal?  (1 Mark)
(a) Carbon (C)
(b) Oxygen (O)
(c) Chlorine (Cl)
(d) Sodium (Na)
Ans: (d)
Sodium (Na) is a metal. Metals are typically found on the left-hand side of the periodic table and have properties like good conductors of heat and electricity, malleability, and ductility. Carbon (C), oxygen (O), and chlorine (Cl) are non-metals.

Q2: The chemical formula of magnesium chloride is:  (1 Mark)
(a) MgCl2
(b) Mg2Cl
(c) MgCl
(d) MgCl
3
Ans: (a)
The chemical formula of magnesium chloride is MgCl2. The subscript 2 indicates that there are two chlorine atoms for every one magnesium atom in the compound.

Q3: Which of the following oxides is amphoteric in nature?  (1 Mark)
(a) Carbon monoxide (CO)
(b) Sulphur dioxide (SO2)
(c) Aluminium oxide (Al2O3)
(d) Nitrogen dioxide (NO2)

Ans: (c)
Amphoteric oxides are those oxides that can react with both acids and bases to form salts and water. Aluminium oxide (Al2O3) is an example of an amphoteric oxide. Carbon monoxide (CO), sulfur dioxide (SO2), and nitrogen dioxide (NO2) are not amphoteric.

Q4: Name one metal and one non-metal that exist as a liquid at room temperature.  (1 Mark)
Ans: Metal: Mercury (Hg)
Non-metal: Bromine (Br2)
Mercury (Hg) is a metal that exists as a liquid at room temperature. Bromine (Br2) is a non-metal that also exists as a liquid at room temperature.

Q5: Define the term "displacement reaction" in the context of metals and give an example.  (1 Mark)
Ans: Displacement Reaction: 
A displacement reaction is a type of chemical reaction where a more reactive metal displaces or replaces a less reactive metal from its salt solution.
Example: Zinc (Zn) displaces copper (Cu) from copper sulfate solution:
Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s)
In this reaction, zinc (a more reactive metal) displaces copper from copper sulfate solution, resulting in the formation of zinc sulfate and copper metal.

Q6: Explain why metals are good conductors of electricity.  (2 marks)
Ans: 
Metals are good conductors of electricity due to the presence of delocalized electrons in their atomic structure. In metallic bonding, metal atoms lose valence electrons to form positive ions (cations). These delocalized electrons move freely throughout the entire metal lattice, creating a "sea of electrons." When an electric potential is applied, these electrons can move easily, carrying the electric charge throughout the metal. This delocalized electron cloud is responsible for the high electrical conductivity of metals.

Q7: Write the balanced chemical equation for the reaction between sulphuric acid (H2SO4) and calcium hydroxide (Ca(OH)2).  (2 marks)
Ans: 
H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2H2O(l)
In the reaction between sulphuric acid (H2SO4) and calcium hydroxide (Ca(OH)2), a double displacement reaction occurs. Sulphuric acid donates two hydrogen ions (H+) to calcium hydroxide, resulting in the formation of calcium sulfate (CaSO4) and water (H2O).

Q8: What happens when metals react with acids? Explain with an example.  (2 marks)
Ans: 
When metals react with acids, they undergo a chemical reaction where the metal displaces hydrogen from the acid, resulting in the formation of metal salts and hydrogen gas.
Example: Zinc (Zn) reacts with hydrochloric acid (HCl):
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
In this reaction, zinc reacts with hydrochloric acid, displacing hydrogen from the acid. Zinc chloride (ZnCl2) is formed as a salt, and hydrogen gas (H2) is released.

Q9: Explain the reactivity series of metals.  (3 marks)
Ans: 
The reactivity series of metals is a list of metals arranged in decreasing order of their reactivity. It helps us understand how metals react with other substances, including acids and water. The general reactivity trend is that more reactive metals displace less reactive metals from their compounds.
The reactivity series from most reactive to least reactive is as follows:

  • Potassium (K)
  • Sodium (Na)
  • Calcium (Ca)
  • Magnesium (Mg)
  • Aluminium (Al)
  • Zinc (Zn)
  • Iron (Fe)
  • Lead (Pb)
  • Copper (Cu)
  • Silver (Ag)
  • Gold (Au)


Q10: Compare the physical properties of metals and non-metals.  (3 marks)
Ans: Physical Properties of Metals:

  • Metals are generally solid at room temperature (except mercury).
  • They have high density and are heavy.
  • Metals are good conductors of heat and electricity.
  • They have a lustrous appearance when freshly cut (except for a few like lead).
  • Metals are malleable (can be hammered into thin sheets) and ductile (can be drawn into wires).
  • Metals have a high melting and boiling point compared to non-metals.

Physical Properties of Non-metals:

  • Non-metals can be solid, liquid, or gas at room temperature.
  • They have lower density compared to metals.
  • Non-metals are poor conductors of heat and electricity (except graphite).
  • Most non-metals are not lustrous; they are dull and do not have a shiny appearance.
  • They are brittle and cannot be easily hammered into sheets or drawn into wires.
  • Non-metals generally have lower melting and boiling points compared to metals.


Q11: Describe the extraction of iron from its ore in the blast furnace.  (3 marks)
Ans: The extraction of iron from its ore (haematite - Fe2O3) is done through the following steps in the blast furnace:

  • Charging the furnace: The blast furnace is loaded with alternate layers of haematite and coke (carbon) along with limestone (CaCO3) as a flux. The coke serves as the source of heat and reduces the iron ore to iron.
  • Reduction: As the furnace is heated to a high temperature (around 1500°C), carbon in the coke reacts with oxygen in the haematite to form carbon monoxide (CO). The carbon monoxide acts as a reducing agent, which reduces iron(III) oxide (Fe2O3) to iron(II) oxide (Fe3O4):
    Fe2O3 + 3CO → 2Fe3O4 + 3CO2
  • Further reduction: The iron(II) oxide then reacts with more carbon monoxide to form iron metal and carbon dioxide:
    Fe3O4 + 4CO -> 3Fe + 4CO2
  • Formation of slag: The impurities in the ore, along with some limestone, form a molten slag that floats on top of the molten iron. The slag is less dense than the molten iron, so it is easily separated.
  • Tapping: Molten iron is tapped off from the bottom of the furnace, and the slag is removed from the top.


Q12: Explain the corrosion of iron and its prevention methods.  (5 marks)
Ans:Corrosion of Iron: Corrosion of iron is the process in which iron reacts with oxygen and moisture from the air to form iron oxide, commonly known as rust. The chemical reaction involved is:
4Fe + 3O2 + 6H2O → 4Fe(OH)3

Prevention Methods:

  • Galvanization: One of the most effective ways to prevent the corrosion of iron is by galvanization. In this method, a thin layer of zinc is coated on the iron surface. Zinc acts as a sacrificial anode, preventing the iron from coming into direct contact with oxygen and moisture. The zinc corrodes instead of iron, protecting it from rusting.
  • Painting: Applying a layer of paint or varnish on the iron surface forms a barrier between the iron and the surrounding environment. It prevents oxygen and moisture from coming in contact with iron, thereby preventing corrosion.
  • Oil and Grease: Applying oil or grease on the iron surface creates a protective layer that prevents oxygen and moisture from reaching the iron, thus inhibiting rust formation.
  • Alloying: Mixing iron with other metals to form alloys like stainless steel can make the iron more corrosion-resistant. Stainless steel contains chromium, which forms a thin layer of chromium oxide on the surface, protecting the iron from corrosion.
  • Drying and Storage: Keeping iron objects dry and stored in a dry environment helps in preventing rust formation. Moisture is a crucial factor in the corrosion process, so eliminating it is essential.

It is essential to employ these preventive methods to increase the longevity of iron objects and structures and avoid damage due to corrosion.

Q13: Discuss the occurrence of metals in nature, giving examples of native metals and ores.    (5 marks)
Ans: 
Occurrence of Metals in Nature:

Native Metals: Some metals occur in nature in their pure elemental form and are known as native metals. They are relatively rare, but they do not require any extraction process. Examples of native metals include:

  • Gold (Au)
  • Silver (Ag)
  • Copper (Cu)
  • Platinum (Pt)

Ores: Most metals occur in nature as compounds, commonly known as ores. Ores are rocks or minerals from which metals can be extracted economically. Various methods are used to extract metals from their ores, depending on the reactivity of the metal. Examples of ores include:

  • Haematite (Fe2O3) and Magnetite (Fe3O4) are ores of iron (Fe).
  • Bauxite (Al2O3.2H2O) is an ore of aluminium (Al).
  • Copper pyrites (CuFeS2) and Malachite (CuCO3.Cu(OH)2) are ores of copper (Cu).
  • Zinc blend (ZnS) is an ore of zinc (Zn).
  • Cinnabar (HgS) is an ore of mercury (Hg).

These ores are processed through various metallurgical processes to obtain pure metals for industrial and commercial use. The extraction of metals from ores involves processes like roasting, calcination, reduction, and refining.

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