Class 10 Science Chapter 3 Metals and Non-metals

Key Concept: Critical Thinking, Complex Scenarios

b) Hardness
[Solution Description] Sonority refers to the ability of a material to produce sound when struck. However, not all metals exhibit this characteristic due to variations in hardness. Metals like mercury, being liquid at room temperature, do not produce sound upon striking, thus contradicting the generalized statement.

Sonority is the property of metals to produce sound when struck. However, not all metals exhibit this characteristic due to differences in hardness. For example, mercury, being liquid at room temperature, does not produce sound upon striking. This highlights that while many metals are sonorous, exceptions exist based on their physical state and hardness.

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Key Concept: Balanced Equation

b) $2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$
[Solution Description]
Magnesium reacts with oxygen in the air to form magnesium oxide. The unbalanced equation is $\text{Mg} + \text{O}_2 \rightarrow \text{MgO}$. To balance the equation, we need two magnesium atoms and one molecule of diatomic oxygen to produce two molecules of magnesium oxide:
$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$.
Therefore, the balanced equation has the coefficients 2 for Mg and MgO.

Magnesium reacts with oxygen in the air to form magnesium oxide. This reaction is an example of a combination reaction where two reactants combine to form a single product. The balanced chemical equation for this reaction is $ 2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO} $. This equation illustrates the stoichiometry of the reaction, indicating that two moles of magnesium react with one mole of oxygen to produce two moles of magnesium oxide.

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Key Concept: Basic Concept

b) Basic
[Solution Description] Magnesium reacts with oxygen in the air to form magnesium oxide (MgO), which is a basic oxide. The reaction is represented by the equation:
$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$
This shows that magnesium oxide is formed, and it is basic in nature.

Aluminium, along with metals like iron and zinc, does not react with cold or hot water but can react with steam to form metal oxides and hydrogen gas. For instance, aluminium reacts with steam to form aluminium oxide $(Al_2O_3)$ and hydrogen gas, as shown in the equation $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$. This behavior is typical for certain metals that require higher energy conditions, such as steam, to react.

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Key concept: Exceptions in Properties

c) Mercury
[Solution Description] Mercury is the only metal that is liquid at room temperature. Most metals are solid except for mercury which exists as a liquid.

Mercury is an exception among metals as it is the only metal that exists in liquid form at room temperature. Unlike most metals, which are solid, mercury’s unique property allows it to remain in a liquid state under standard conditions.

Notes:

Key Concept: Exothermic Reaction

c) Potassium with water
[Solution Description] To determine which reaction is the most exothermic, we need to consider the nature of the metals and their reactivity with water. Sodium and potassium are highly reactive with water, and their reactions are notably exothermic, often resulting in hydrogen catching fire. Given these options, the reaction of potassium with water is more exothermic compared to calcium and magnesium due to its high reactivity.

Potassium reacting with water:
$2K + 2H_2O \rightarrow 2KOH + H_2 + \text{heat}$
This reaction releases a significant amount of heat energy.

Potassium’s reaction with water is highly exothermic, producing potassium hydroxide and hydrogen gas. This reaction is so vigorous that the hydrogen gas evolved often ignites. The general reaction for metals with water is $ \text{Metal} + \text{Water} \rightarrow \text{Metal hydroxide} + \text{Hydrogen gas} $. For potassium, the equation is $ 2\text{K} + 2\text{H}_2\text{O} \rightarrow 2\text{KOH} + \text{H}_2 + \text{heat} $.

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Key Concept: Reaction Prediction

c) Sodium sulfide
[Solution Description] When sulfur reacts with a strong base such as sodium hydroxide (NaOH), it forms sodium sulfide $(Na_2S)$. This reaction typically involves breaking down of sulfur in presence of a hydroxide ion leading to formation of a sulfide compound.

When considering reactions between non-metals and bases, it is important to note that non-metals typically form acidic or neutral oxides. These oxides do not generally react with bases to form salts or other compounds, as non-metals do not readily donate electrons or protons in such reactions.

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Key Concept: Element Identification

d) Chlorine
[Solution Description] The electronic configuration given is 2,8,7, which means there are two electrons in the first shell (K), eight electrons in the second shell (L), and seven electrons in the third shell (M). This matches with chlorine, which has an atomic number of 17.

Chlorine is a non-metal with the electronic configuration 2,8,7, indicating it has 17 electrons. This configuration shows that chlorine has two electrons in the K shell, eight in the L shell, and seven in the M shell. Chlorine tends to gain one electron to achieve a stable octet configuration, forming a chloride anion $(Cl^-)$.

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Key Concept: Properties of Metal Oxides

c) Aluminium oxide $(Al_2O_3)$

[Solution Description]

An amphoteric oxide can react with both acids and bases. Given the options, we recall that aluminium oxide $(Al_2O_3)$ demonstrates this behavior. It reacts with hydrochloric acid to form aluminium chloride and water, as well as reacting with bases. Therefore, $Al_2O_3$ is an amphoteric oxide.

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Key Concept: Property Identification

b) Acidic
[Solution Description] Non-metal oxides typically form acidic solutions when dissolved in water because they react with water to produce acids.

Non-metal oxides generally exhibit acidic properties when dissolved in water. They react with water to form acids, which is why solutions of non-metal oxides are typically acidic. This is a key characteristic that distinguishes non-metal oxides from metal oxides, which are usually basic.

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Key Concept: Process Comparison

c) Calcination
[Solution Description] Calcination involves heating the ore in a limited supply of air or in the absence of air. This process is typically used for carbonate ores where the desired outcome is to remove volatile components such as carbon dioxide $(\text{CO}_2)$, leaving behind the metal oxide. In contrast, roasting involves heating sulfide ores in the presence of excess air. Hence, calcination is the process for this transformation.

Calcination is a process used to convert carbonate ores into oxides by heating them in a limited supply of air. This process removes volatile components such as carbon dioxide, leaving behind the metal oxide. Calcination is distinct from roasting, which involves heating sulphide ores in the presence of excess air to convert them into oxides.

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Key concept: Basic Identification

b) Carbon
[Solution Description] Non-metals are elements that typically lack the properties associated with metals, such as luster, malleability, and good electrical conductivity. Common examples of non-metals include carbon, oxygen, sulfur, and iodine.

Non-metals are elements that typically lack the properties associated with metals, such as luster, malleability, and good electrical conductivity. Common examples of non-metals include carbon, oxygen, sulfur, and iodine. These elements are generally poor conductors of heat and electricity and are not malleable or ductile.

Notes:

Key Concept: Reaction Equation

c) $Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2$
[Solution Description] When magnesium reacts with hot water, it forms magnesium hydroxide and hydrogen gas. The balanced chemical equation for this reaction is:
newline $Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$.

When magnesium reacts with hot water, it forms magnesium hydroxide and hydrogen gas. The reaction can be represented by the equation: $ Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g) $. This reaction is less vigorous compared to the reactions of metals like sodium and potassium with water.

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Key Concept: Metal Reaction Order

d) Potassium, Calcium, Iron

[Solution Description]

According to the activity series, potassium $(K)$ is highly reactive and reacts vigorously with water, followed by calcium $(Ca)$, which reacts less violently than potassium but more vigorously than iron $(Fe)$. Iron reacts very slowly with water. Hence, the order of reactivity with water from highest to lowest is: Potassium, Calcium, Iron.

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Key Concept: Process Application

c) Calcination
[Solution Description] Zinc carbonate is converted into zinc oxide through the process of calcination. This involves heating the carbonate ore in limited air, leading to the decomposition of $ZnCO_3$ to ZnO and $CO_2$. The chemical equation for this reaction is:
$\text{ZnCO}_3(s) \xrightarrow{\text{Heat}} \text{ZnO}(s) + \text{CO}_2(g)$
Calcination is specifically used for carbonates, while roasting is used for sulphides.

Calcination is a process used to convert carbonate ores into oxides by heating them in limited air. This process is specifically used for carbonate ores like zinc carbonate $(ZnCO_3)$, which decomposes into zinc oxide (ZnO) and carbon dioxide $(CO_2)$ when heated. The chemical reaction for this process is: $ \text{ZnCO}_3(s) \xrightarrow{\text{Heat}} \text{ZnO}(s) + \text{CO}_2(g) $.

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Key Concept: Electron Configuration

c) $1s^2 2s^2 2p^6 3s^2 3p^5$
[Solution Description] Chlorine has an atomic number of 17, so its electron configuration is $1s^2 2s^2 2p^6 3s^2 3p^5$ before reacting with magnesium. It needs one more electron to complete its octet.

Chlorine, before reacting with magnesium, has an electron configuration of $ 1s^2 2s^2 2p^6 3s^2 3p^5 $. Chlorine has an atomic number of 17, which means it has 17 electrons. It needs one more electron to complete its octet, making it highly reactive and ready to form compounds like magnesium chloride.

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Key concept: Lustre Identification

c) Metallic lustre
[Solution Description] Metals have a characteristic shiny appearance because they have a property known as metallic lustre. This means that when light hits their surface, it reflects in such a way that gives these metals a brilliant shine.

Metals, in their pure state, exhibit a shiny appearance known as metallic lustre. This property is due to the ability of metals to reflect light from their surface, which is why metals like gold and silver are often used in decorative items and jewelry.

Notes :

• Lustre: The property of shining appearance due to the reflection of light from the surface of a material.

Metals:

1. Metallic Lustre: Metals have a shiny appearance when freshly cut, polished, or struck.
2. Examples:
   • Gold, Silver, and Copper: Known for their bright and reflective surface.
   • Aluminum and Zinc: Exhibit a less pronounced shine but still have a metallic appearance.
3. Reason: Free electrons in metals reflect light, producing lustre.

Non-Metals:

1. Non-Lustrous: Most non-metals do not exhibit a shiny appearance.
2. Exceptions:
   • Iodine: A non-metal with a shiny, metallic-like surface.
   • Graphite: Exhibits a dull metallic lustre.
3. Reason: Lack of free electrons results in poor light reflection.

• Practical Test: Scratching or polishing surfaces helps in identifying lustre in metals.

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Key Concept: Brittle Nature

c) Sulphur
[Solution Description] Sulphur is a non-metal that is known for its brittleness, breaking easily when pressure is applied.

Sulphur is a non-metal known for its brittleness, meaning it can easily break or crumble when pressure is applied. This property is typical of non-metals, which do not possess the malleability or ductility found in metals.

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Key Concept: Extraction Method Selection

b) Electrolysis
[Solution Description] Metals high up in the reactivity series such as sodium and magnesium are extracted using electrolysis because they cannot be reduced by carbon due to their high affinity for oxygen. Therefore, electrolytic reduction is employed to extract these metals from their compounds:
$\text{Na}^+ + e^- \rightarrow \text{Na}$
At cathode and $\text{2Cl}^- \rightarrow \text{Cl}_2 + 2e^-$
At anode.

Electrolysis is the method used for extracting metals at the top of the reactivity series, such as sodium and magnesium, because these metals cannot be reduced by carbon. During electrolysis, the metal ions are reduced at the cathode, while the anions are oxidized at the anode. For example, in the electrolysis of molten sodium chloride:
At cathode: $\text{Na}^+ + e^- \rightarrow \text{Na}$
At anode: $2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$

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Key Concept: Electrolysis Setup

a) Graphite
[Solution Description] In the electrolysis of aluminium oxide, often using molten cryolite to dissolve alumina, the cathode is made from carbon. Electric current causes aluminum ions to move towards this negatively charged electrode where they gain electrons and deposit as aluminum metal.

According to the reactivity series, metals like iron are less reactive compared to zinc. Iron can be extracted from its ores using carbon reduction methods. This is because iron is positioned below carbon in the reactivity series, allowing carbon to reduce iron oxides to iron metal.

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Environmental Impact, Complex Reactions

b) Both Assertion and Reason are true, but Reason is NOT the correct explanation of Assertion.
[Solution Description] Non-metal oxides indeed include carbon dioxide $(CO_2)$, which is a significant greenhouse gas contributing to global warming. The reason states the fact that carbon dioxide is released through the combustion of fossil fuels, which is true. However, not all non-metal oxides are greenhouse gases; some contribute primarily to issues like acid rain or smog rather than directly to global warming. Therefore, while both statements are correct, the reason provided does not exclusively explain why non-metal oxides as a group are major contributors to greenhouse gases.

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Key Concept: Ionic Conductivity

b) Molten

[Solution Description] Ionic compounds conduct electricity when they are molten or dissolved in water because their ions are free to move and carry charge. In the solid state, ions are fixed in place within the crystal lattice and cannot move freely, hence ionic compounds do not conduct electricity in this state.

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Key Concept: Free State Metals

b) Copper
[Solution Description] Metals like gold and silver are often found in their free state due to their low reactivity compared to other metals.

Copper is commonly found in nature in the free state due to its low reactivity compared to other metals. While copper can also be found in combined states as sulphide or oxide ores, its presence in the free state is attributed to its resistance to oxidation and corrosion.

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Key Concept: Process Comparison

b) Calcination
[Solution Description] The conversion of carbonate ores into metal oxides involves calcination, which is heating in limited air.

Calcination is a process used to convert carbonate ores into metal oxides. This involves heating the carbonate ores strongly in limited air, which results in the decomposition of the carbonate to form metal oxide and carbon dioxide gas. This process is distinct from roasting, which is used for sulphide ores and involves heating in the presence of excess air.

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Key concept: Ionic Compound Properties, Conduction of Electricity

c) Lattice breaks down releasing free ions in molten state
[Solution Description] Magnesium chloride in solid state has ions fixed in place within its crystalline lattice structure, restricting their movement. Thus, it cannot conduct electricity. However, when it melts, the electrostatic forces between ions are overcome due to heat, allowing ions to move freely. These free-moving ions facilitate the conduction of electricity in the molten state.

Magnesium chloride in solid state has ions fixed in place within its crystalline lattice structure, restricting their movement. Thus, it cannot conduct electricity. However, when it melts, the electrostatic forces between ions are overcome due to heat, allowing ions to move freely. These free-moving ions facilitate the conduction of electricity in the molten state.

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Key Concept: Oxide Solubility

d) Sodium oxide
[Solution Description] Some metal oxides dissolve in water to produce alkalis. Sodium oxide dissolves in water to form sodium hydroxide, which is an alkaline solution. Therefore, the correct answer is sodium oxide.

Some metal oxides dissolve in water to produce alkaline solutions, known as alkalis. Sodium oxide is one such metal oxide that dissolves in water to form sodium hydroxide, an alkaline solution. The reaction can be represented as: $Na_2O(s) + H_2O(l) → 2NaOH(aq)$.

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