Key concepts in chemistry Revision Notes

Deducing Empirical Formula from Molecular Formula

The empirical formula is the simplest whole-number ratio of atoms in a compound. To deduce it from a molecular formula, divide all the subscripts in the molecular formula by the largest common factor. For instance, the molecular formula for glucose is C₆H₁₂O₆. The largest common factor for 6, 12, and 6 is 6. Dividing by 6 gives C₁H₂O₁, so the empirical formula is CH₂O.

Deducing Molecular Formula from Empirical Formula and Relative Molecular Mass

If you have the empirical formula and the relative molecular mass (M₉) of a compound, you can find its molecular formula. First, calculate the relative empirical formula mass. Then, divide the relative molecular mass by the relative empirical formula mass. This gives you a whole number, which is the factor by which you multiply the subscripts in the empirical formula to get the molecular formula. For example, if the empirical formula is CH₂ and the M₉ is 42, the empirical formula mass is 12.0 + (2  1.0) = 14.0. The factor is 42.0 / 14.0 = 3. So the molecular formula is (CH₂)₃, which is C₃H₆.

Allotropes of Carbon

Diamond and graphite are allotropes of carbon, meaning they are different structural forms of the same element. Their distinct properties arise from the different ways carbon atoms are bonded together in their giant covalent structures.

Giant Covalent Structure Definition

Giant covalent structures consist of a vast network of atoms held together by strong covalent bonds. Unlike simple molecular structures, these do not form individual molecules but rather continuous, extended lattices. This type of bonding typically results in high melting points and hardness.

Diamond as a Giant Covalent Structure

In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. This forms a rigid, three-dimensional lattice, making diamond exceptionally hard and giving it a very high melting point. It does not conduct electricity due to the absence of delocalised electrons.

Graphite as a Giant Covalent Structure

In graphite, each carbon atom is covalently bonded to three other carbon atoms, forming hexagonal rings in layers. Within these layers, the bonds are strong covalent bonds. Between the layers, however, there are only weak intermolecular forces, allowing the layers to slide over each other. This structure makes graphite soft and slippery. It also has delocalised electrons between the layers, allowing it to conduct electricity.

Diamond Structure

Diamond is a giant covalent structure. Each carbon atom is bonded to four other carbon atoms. The arrangement forms a strong, rigid, three-dimensional lattice. This structure accounts for diamond's extreme hardness and very high melting point. All outer shell electrons are used in bonding, so diamond does not conduct electricity.

Graphite Structure

Graphite is also a giant covalent structure, but it has a layered arrangement. Each carbon atom bonds to three other carbon atoms within a layer, forming hexagonal rings. Weak intermolecular forces hold the layers together, allowing them to slide. Each carbon atom has one delocalised electron, which allows graphite to conduct electricity and heat.

Key Differences and Properties

Diamond's strong 3D network makes it very hard and a poor electrical conductor. Graphite's layered structure with weak forces makes it soft, slippery, and a good electrical conductor due to delocalised electrons. Both have high melting points due to strong covalent bonds requiring significant energy to break.

Diamond Structure and Properties

Diamond has a giant covalent structure where each carbon atom forms strong covalent bonds with four other carbon atoms in a rigid, tetrahedral arrangement. This makes diamond extremely hard and gives it a very high melting point. All valence electrons are localised in strong covalent bonds, meaning there are no free electrons to carry charge, so diamond does not conduct electricity.

Graphite Structure and Properties

Graphite also has a giant covalent structure but is arranged in layers of hexagonal rings. Within each layer, each carbon atom is bonded to three other carbon atoms, leaving one valence electron delocalised. These delocalised electrons are free to move, allowing graphite to conduct electricity, making it suitable for electrodes. The layers are held together by weak intermolecular forces, allowing them to slide easily over each other, which makes graphite a good lubricant.

Applications based on Structure

Diamond's extreme hardness (due to strong 3D covalent network) makes it ideal for cutting tools. Graphite's ability to conduct electricity (due to delocalised electrons) makes it useful for electrodes. Its slipperiness (due to weak forces between layers) allows it to be used as a lubricant.