Question 1: Define a covalent bond and explain its significance in the structure of diamond and graphite.

Answer: A covalent bond is a strong attraction between two non-metal atoms sharing a pair of electrons. In diamond, each carbon atom forms four covalent bonds in a tetrahedral structure, making it very hard. In graphite, each carbon atom forms three covalent bonds in a planar hexagonal arrangement, with weak forces between layers.

Question 2: Describe the difference in atomic arrangement between diamond and graphite.

Answer: Diamond has a three-dimensional tetrahedral network where each carbon is bonded to four others, forming a rigid structure. Graphite has layered sheets of hexagonally arranged carbon atoms with weak forces between layers, allowing them to slide over each other.

Question 1: Calculate the number of covalent bonds in a single crystal of diamond consisting of 1,000,000 carbon atoms. (Assume each carbon forms four covalent bonds).

Answer: 4,000,000 bonds

Question 2: If the density of graphite is 2.26 g/cm³ and the molar mass of carbon is 12.01 g/mol, calculate the number of carbon atoms in 1 cm³ of graphite.

Answer: 1.42 x 10^23 atoms

Question 1: Describe a safe method to test the electrical conductivity of diamond and graphite samples.

Answer: Connect each sample to a circuit with a power supply and a bulb. If the bulb lights, the material conducts electricity. Graphite conducts due to delocalised electrons; diamond does not because it lacks free electrons. Ensure safety by handling electrical equipment carefully and avoiding water contact.

Question 2: Identify the independent, dependent, and control variables in an experiment testing the hardness of diamond versus graphite.

Answer: Independent variable: type of material (diamond or graphite). Dependent variable: hardness measured (e.g., via scratch test). Control variables: sample size, testing method, pressure applied.

Question 1: A sample of diamond has a melting point of approximately 3550°C, while graphite sublimates at about 3652°C. Explain why these values differ significantly.

Answer: Diamond's strong 3D covalent network requires a lot of energy to break, resulting in a high melting point. Graphite sublimates at a slightly higher temperature because its layers are held together by weak forces, allowing sublimation at high temperatures with less energy needed to break covalent bonds.

Question 2: The thermal conductivity of diamond is extremely high, whereas graphite's is moderate. Explain the reason based on their structures.

Answer: Diamond's rigid 3D covalent network allows efficient transfer of heat through vibrations of atoms. In graphite, free delocalised electrons and layered structure facilitate heat transfer within layers but less so between layers, resulting in moderate thermal conductivity.

Question 1: Explain the differences in bonding and structure between diamond and graphite and how these differences influence their physical properties such as hardness and electrical conductivity. (6 marks)

Answer: Diamond has a three-dimensional tetrahedral network where each carbon atom forms four covalent bonds, making it extremely hard and insulating. Graphite consists of layers of hexagonally arranged carbon atoms bonded covalently within layers, with weak forces between layers, making it softer and able to conduct electricity due to delocalised electrons. These structural differences directly affect their hardness, electrical conductivity, and other physical properties.

Question 1: Describe one industrial use of diamond and explain why its structure makes it suitable for this purpose.

Answer: Diamonds are used in cutting and grinding tools due to their exceptional hardness, which results from their strong covalent network structure. This makes them ideal for cutting through tough materials like metals and stone.