I have 3 questions:
1) how might the actual structure of a diamond help explain why it is sometimes found in the form of large, single crystals?
2) what feature of graphite's structure might
account for its usefulness as a lubricant?
3) why are fullerenes powfery as solids rather than composed of large-scale chunks?
On 1, I'm not quite sure what your prof is looking for. We know carbon bonds to itself in relatively strong bonds and we know diamond has the face centered cubic structure with four tetrahedra inside a unit cell. Thus carbon corners are held together as well as each face is bonded. That explains why it is hard but not necessarily why it occurs as chunks. We also know that it cleaves very nicely along the planes because of the cubic structure. We know it comes up in volcanic pipes and the force of the eruption plus abrasion along the way may cleave large pieces into smaller pieces. I may be answering the question a little backwards; my point is that they were formed in larger pieces but cleaved on their way to the surface or as volcanic ashes were eroded to leave the diamonds in smaller pieces.
2. Graphite occurs in sheets. Air is adsorbed (not absorbed) on the surface of graphite. The sheet formation makes it easy for one layer to slide across another.
3. Fullerenes are large molecules (as compared to many others); however, they are not in the same league as diamonds. Furthermore, fullerenes crystallize in pentagonal or heptagonal structures that fit together to form the "buckyball." There is nothing to hold one buckyball to another buckyball; therefore, it is more likely to occur in powder form. In fact, soot contains some fullerenes.
Check my thinking. Do some research of your own and use that plus the above to make your own decisions.
1) One of the reasons why diamonds are found in the form of large, single crystals is due to the unique crystal structure of diamonds. Diamonds are made up of a repeating lattice structure of carbon atoms in a tetrahedral arrangement. Each carbon atom is covalently bonded to four neighboring carbon atoms, forming a strong and rigid three-dimensional structure. This structure makes diamonds extremely hard and gives them their characteristic brilliance.
The strong covalent bonds between the carbon atoms in a diamond make it difficult for impurities or defects to occur within the crystal structure. As a result, when the conditions are right for diamond formation, such as high-pressure environments deep within the Earth, the carbon atoms can naturally align and bond together to form large, single crystals.
2) Graphite's usefulness as a lubricant can be attributed to its layered crystal structure. In graphite, carbon atoms are arranged in stacked layers held together by weak van der Waals forces. Within each layer, the carbon atoms are bonded together through strong covalent bonds.
The weak interlayer forces make it easy for the layers to slide over each other, allowing the graphite to act as a lubricant. When graphite is applied between two surfaces, such as in machinery, the layers of carbon atoms break away and create a smooth surface that reduces friction. This property enables graphite to be an effective lubricant.
3) Fullerenes, such as buckyballs (C60 molecules), are powdery as solids rather than composed of large-scale chunks due to their molecular structure. Fullerenes have a spherical cage-like structure made up of carbon atoms, similar to a soccer ball or a geodesic dome.
The carbon atoms in a fullerene are arranged in a pattern of pentagons and hexagons, forming a closed network. This arrangement creates a curved surface, and the resulting molecule is not easily stackable or compressible like other regular solids.
As a result, fullerenes tend to exist as individual molecules or small clusters instead of forming large-scale chunks. This is why they appear powdery as solids because their unique structure does not allow them to form extensive crystal lattices or solid structures like diamonds or graphite.