Elements, compounds and mixtures Study Guide

The term 'pure' carries different meanings depending on whether it's used in a scientific context or in everyday language. In everyday use, 'pure' often implies something is natural, unadulterated, or free from harmful substances. For example, 'pure orange juice' suggests it contains only juice with no added sugar or artificial ingredients. Similarly, 'pure gold' might refer to gold with a very high proportion of gold metal, even if it contains trace amounts of other elements.

Scientifically, however, a 'pure substance' has a much more precise definition. A pure substance consists of only one type of element or one type of compound. This means all the particles (atoms or molecules) in a pure substance are chemically identical. For instance, pure water (H₂O) consists solely of water molecules and nothing else. Pure iron (Fe) contains only iron atoms. The key characteristic of a scientifically pure substance is that it has a fixed and sharp melting point and boiling point. If a substance starts to melt or boil over a range of temperatures, it indicates that it is an impure substance, or a mixture, because the presence of impurities disrupts the orderly arrangement of particles and requires different amounts of energy to overcome intermolecular forces. Impurities lower the melting point and raise the boiling point of a substance, and broaden the temperature range over which these phase changes occur. Therefore, scientific purity is rigorously tested by observing these physical properties. For example, a pure substance will boil at a specific temperature (e.g., 100°C for pure water at standard pressure), whereas an impure sample will boil over a range of temperatures above 100°C.

Carbon exists in several allotropic forms, each with distinct properties determined by its atomic structure and bonding. Diamond, graphite, fullerenes, and graphene are all made of carbon atoms, but their arrangements lead to vastly different characteristics. Diamond has a giant covalent structure where each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. This strong 3D network requires a huge amount of energy to break, making diamond incredibly hard with a very high melting point, and it does not conduct electricity because all outer electrons are used in bonding, leaving none free to move. Graphite also has a giant covalent structure, but its carbon atoms are arranged in hexagonal layers. Within each layer, each carbon atom is bonded to three others, creating strong covalent bonds. However, the layers are held together by weak intermolecular forces. This allows the layers to slide over each other, making graphite soft and slippery. Furthermore, because each carbon atom only forms three bonds, one outer electron per carbon is delocalised and free to move within the layers, making graphite an excellent electrical conductor. Fullerenes are molecules of carbon atoms forming hollow spheres or tubes, such as Buckminsterfullerene (C60). Their spherical or tubular structures allow them to trap other atoms or be used as lubricants or drug delivery systems. They do not conduct electricity well in their pure form, but some can be modified to do so. Graphene is a single layer of graphite, just one atom thick. It possesses exceptional strength due to its strong covalent bonds arranging carbon atoms in a hexagonal lattice. Its 2D structure and delocalised electrons mean it is an excellent conductor of electricity and heat, and it is also transparent and incredibly light. Understanding these structural differences is key to explaining the varied applications of carbon allotropes.