Problem 34
Question
Assertion : Among the carbon allotropes, diamond is an insulator, whereas, graphite is a good conductor of electricity. Reason : Hybridization of carbon in diamond and graphite are \(s p^{3}\) and \(s p^{2}\), respectively. [Main Online April 10, 2016] (a) Both assertion and reason are correct, but the reason is not the correct explanation for the assertion (b) Both assertion and reason are correct, and the reason is the correct explanation for the assertion (c) Both assertion and reason are incorrect (d) Assertion is incorrect statement, but the reason is correct. This question contains STATEMENT-1 (Assertion/ Statement ) and STATEMENT-2 (Reason/Explanation) and has 4 choices (a), (b), (c) and (d) out of which ONLY ONE is correct. (a) Statement-1 is True, Statement-2 isTrue; Statement-2 is a correct explanation for Statement-1 (b) Statement-1 is True, Statement-2 is True; Statement-2 is NOT a correct explanation for Statement-1 (c) Statement-1 is True, Statement- 2 is False (d) Statement-1 is False, Statement- 2 is True.
Step-by-Step Solution
Verified Answer
(b) Both assertion and reason are correct, and the reason is the correct explanation for the assertion.
1Step 1: Understand the Assertion
The assertion states that among carbon allotropes, diamond is an insulator and graphite is a good conductor of electricity. This is a true statement because in diamond, carbon atoms are arranged in a three-dimensional tetrahedral structure, resulting in no free electrons to conduct electricity. On the other hand, graphite has a layered structure where electrons can move freely between layers, making it a good conductor.
2Step 2: Analyze the Reason
The reason provided is that the hybridization of carbon in diamond is \( sp^{3} \) and in graphite is \( sp^{2} \). In diamond, each carbon forms four bonds using \( sp^{3} \) hybrid orbitals, creating a strong covalent bond network. In graphite, each carbon atom forms three \( sp^{2} \) hybridized bonds and one p-orbital remains unhybridized, allowing delocalized electrons that contribute to conductivity.
3Step 3: Check the Correlation between Assertion and Reason
Since the reason given explains the presence of conductance in graphite due to \( sp^{2} \) hybridization but does not directly link the hybridization types to the electrical properties naming diamond specifically as an insulator, we look at whether the hybridization itself explains the electrical characteristics. In this case, the hybridization type directly affects bonding and electronic properties explaining why diamond lacks free electrons and graphite allows electron movement.
4Step 4: Determine Correct Option
Given that both the assertion and reason are true, the reason correctly explains why diamond is an insulator (due to \( sp^{3} \) hybridization) and graphite is a conductor (due to \( sp^{2} \) hybridization), the correct option is (b) as both statements are true and the reason correctly explains the assertion.
Key Concepts
DiamondGraphiteHybridizationElectrical Conductivity
Diamond
Diamond is a fascinating carbon allotrope, known for its extraordinary hardness and brilliance. In diamond, each carbon atom forms four strong covalent bonds with other carbon atoms, creating a robust three-dimensional structure. This configuration is called a tetrahedral lattice. The exceptional strength of diamonds comes from these strong chemical bonds, which result in a highly stable and rigid network of carbon atoms.
One distinctive characteristic of diamond is its status as an insulator. Insulators are materials that do not allow the flow of electrical current. In diamonds, all the electrons are tightly bound in these covalent bonds, leaving no free electrons to move around. Therefore, this structure lacks the mobile charge carriers needed for conducting electricity, thus categorizing diamond as an electrical insulator. Despite their electrical non-conductivity, diamonds are widely used in various industrial applications due to their hardness.
One distinctive characteristic of diamond is its status as an insulator. Insulators are materials that do not allow the flow of electrical current. In diamonds, all the electrons are tightly bound in these covalent bonds, leaving no free electrons to move around. Therefore, this structure lacks the mobile charge carriers needed for conducting electricity, thus categorizing diamond as an electrical insulator. Despite their electrical non-conductivity, diamonds are widely used in various industrial applications due to their hardness.
Graphite
Graphite stands as a stark contrast to diamond, despite both being allotropes of carbon. Graphite has a layered structure where each layer comprises carbon atoms arranged in a hexagonal lattice. These layers are held together by weaker van der Waals forces, allowing them to slide easily over one another. This gives graphite a slippery feel and is why it's used as a lubricant and in pencils.
One of the remarkable properties of graphite is its ability to conduct electricity, distinguishing it from diamond. This is due to its structure, where each carbon atom in graphite is bonded to only three other carbon atoms in a planar arrangement. The fourth electron is free to move, leading to electron delocalization across the layers. This presence of free-moving electrons allows graphite to conduct electricity efficiently. Therefore, graphite is often utilized in applications requiring electrical conduction, such as electrodes in batteries and solar panels.
One of the remarkable properties of graphite is its ability to conduct electricity, distinguishing it from diamond. This is due to its structure, where each carbon atom in graphite is bonded to only three other carbon atoms in a planar arrangement. The fourth electron is free to move, leading to electron delocalization across the layers. This presence of free-moving electrons allows graphite to conduct electricity efficiently. Therefore, graphite is often utilized in applications requiring electrical conduction, such as electrodes in batteries and solar panels.
Hybridization
Hybridization is a concept in chemistry that helps explain the shape and bonding of molecules. In carbon allotropes like diamond and graphite, different types of hybridization occur, affecting their distinct properties.
In diamond, carbon atoms undergo \( sp^3 \) hybridization. This means that one s-orbital mixes with three p-orbitals to form four equivalent hybrid orbitals. These \( sp^3 \) hybrid orbitals form tetrahedral configurations, leading to the strong and rigid covalent bonds seen in diamonds.
On the other hand, graphite involves \( sp^2 \) hybridization. In this scenario, one s-orbital combines with two p-orbitals to create three \( sp^2 \) hybrid orbitals. Each carbon atom forms three strong bonds within the layer. The remaining p-orbital is left unhybridized, contributing to \( \pi \) bonding and resulting in electron delocalization. This is why different hybridization types in diamond and graphite lead to distinct properties in terms of hardness, structure, and electrical conductivity.
In diamond, carbon atoms undergo \( sp^3 \) hybridization. This means that one s-orbital mixes with three p-orbitals to form four equivalent hybrid orbitals. These \( sp^3 \) hybrid orbitals form tetrahedral configurations, leading to the strong and rigid covalent bonds seen in diamonds.
On the other hand, graphite involves \( sp^2 \) hybridization. In this scenario, one s-orbital combines with two p-orbitals to create three \( sp^2 \) hybrid orbitals. Each carbon atom forms three strong bonds within the layer. The remaining p-orbital is left unhybridized, contributing to \( \pi \) bonding and resulting in electron delocalization. This is why different hybridization types in diamond and graphite lead to distinct properties in terms of hardness, structure, and electrical conductivity.
Electrical Conductivity
Electrical conductivity is a property that defines how well a material can allow the flow of electric current. Materials that exhibit high electrical conductivity have many free electrons that can move freely, enabling the transfer of electricity.
In the case of carbon allotropes, graphite is an excellent conductor of electricity due to its layered structure and electron delocalization. In graphite, some of the electrons aren't bound tightly in covalent bonds, which facilitates the flow of electric charge across the layers. This makes graphite suitable for use in applications like electrodes and conductive materials.
By contrast, diamond's nature as an insulator stems from its strong, fixed covalent bonds in a three-dimensional network. Since all electrons are involved in these bonds, there are no free electrons to support electrical conduction. Thus, the comparison between diamond and graphite offers a compelling example of how atomic structure and bonding can influence a material's electrical properties.
In the case of carbon allotropes, graphite is an excellent conductor of electricity due to its layered structure and electron delocalization. In graphite, some of the electrons aren't bound tightly in covalent bonds, which facilitates the flow of electric charge across the layers. This makes graphite suitable for use in applications like electrodes and conductive materials.
By contrast, diamond's nature as an insulator stems from its strong, fixed covalent bonds in a three-dimensional network. Since all electrons are involved in these bonds, there are no free electrons to support electrical conduction. Thus, the comparison between diamond and graphite offers a compelling example of how atomic structure and bonding can influence a material's electrical properties.
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