Why does graphite conduct electricity whereas diamond does not, even though both are allotropes of pure carbon?
Answer and explanation
Correct answer: A. Each carbon forms 3 bonds in layers
- Step 1: In graphite, each carbon atom bonds covalently to three other carbons in 2D planar sheets. Step 2. The fourth valence electron is delocalised across the layer, acting as a mobile charge carrier that conducts electric current. Answer: Each carbon forms 3 bonds in layers. Why not the others: In diamond, all four valence electrons are locked into localised covalent bonds.
Practise more Chemical Reactivity, Bonding, Structure and Organic Chemistry.
Free questions with an explanation for every one, and full timed National Certificate of Educational Achievement mock tests.
More Chemical Reactivity, Bonding, Structure and Organic Chemistry questions
- What is the electron configuration of a neutral sodium atom (atomic number 11), and what ion does it form?
- Why do solid ionic compounds like sodium chloride (NaCl) not conduct electricity, whereas molten or aqueous NaCl conducts readily?
- What type of bonding exists within a diamond crystal, and why does diamond have an exceptionally high melting point?
- What gas is produced when dilute hydrochloric acid reacts with calcium carbonate (marble chips)?
- According to collision theory, why does increasing the temperature of a reaction mixture increase the rate of reaction?
- What is the pH of a neutral aqueous solution at 25 deg C, and what happens to pH as acidity increases?