Bonding, Structure and the Properties of Matter

Chemical bonding explains how atoms are held together in substances. The three main types of strong chemical bond are ionic, covalent and metallic bonding. Ionic bonding involves electrostatic attraction between oppositely charged ions, covalent bonding involves shared pairs of electrons between atoms, and metallic bonding involves positive metal ions attracted to delocalised electrons. The type of bonding and structure determines many physical properties, including melting point, boiling point, electrical conductivity, hardness and state at room temperature.

Ionic bonding normally occurs when metals react with non-metals. Metal atoms lose electrons from their outer shell and form positively charged ions, while non-metal atoms gain electrons and form negatively charged ions. Metals in Groups 1 and 2 typically form ions with charges of +1 and +2, while non-metals in Groups 6 and 7 typically form ions with charges of −2 and −1. The oppositely charged ions are held together by strong electrostatic forces of attraction. Students should be able to work out ion charges from group numbers and represent ionic bonding using dot-and-cross diagrams.

Ionic compounds form giant ionic lattices containing large numbers of positive and negative ions arranged in a regular repeating structure. Strong electrostatic attractions act between oppositely charged ions in all directions throughout the lattice. Students should recognise ionic structures from diagrams, interpret models of ionic lattices and determine empirical formulae from the ratio of ions present. Sodium chloride is an important example of a giant ionic lattice.

Covalent bonding occurs when atoms share pairs of electrons. A covalent bond is a strong electrostatic attraction involving the shared pair of electrons and the nuclei of the bonded atoms. Covalent bonding occurs in non-metal elements and compounds made from non-metals. Students should be able to draw dot-and-cross diagrams and structural formulae for hydrogen, chlorine, oxygen, nitrogen, hydrogen chloride, water, ammonia and methane, and interpret diagrams showing covalent bonding.

Many covalently bonded substances consist of small molecules. They usually have relatively low melting and boiling points because only the weak intermolecular forces between molecules need to be overcome when the substance melts or boils; the strong covalent bonds inside each molecule are not broken. Larger molecules generally have stronger intermolecular forces and therefore higher melting and boiling points. Small molecular substances usually do not conduct electricity because they do not contain free-moving charged particles.

Polymers are very large molecules made from many repeating units joined by strong covalent bonds. The forces between polymer molecules are generally stronger than those between small molecules, so many polymers are solids at room temperature. Students should be able to recognise polymers from structural diagrams, identify repeating units and distinguish a polymer structure from a small molecular structure.

Metallic bonding occurs in metals and alloys. Metals contain a giant structure of positive metal ions surrounded by delocalised electrons. These electrons are free to move throughout the structure. Strong electrostatic attraction between the positive metal ions and the negatively charged delocalised electrons holds the metallic lattice together. This model explains the high melting points of many metals and several of their characteristic physical properties.

Pure metals consist of layers of atoms arranged in a regular pattern. These layers can slide over one another, making many pure metals malleable and allowing them to be bent or shaped. Alloys are mixtures containing a metal and one or more other elements. Different-sized atoms disrupt the regular layers, making it more difficult for them to slide over each other, so alloys are generally harder than pure metals. Brass and bronze are examples of alloys.

Metals conduct electricity because their delocalised electrons are free to move through the metallic structure and carry electrical charge. This explains why metals are commonly used as electrical conductors. Metals are also good conductors of thermal energy because energy can be transferred rapidly through the structure by the delocalised electrons and by vibrations of particles within the metallic lattice.

The three common states of matter are solid, liquid and gas. In solids, particles are closely packed in fixed positions and vibrate; in liquids, particles remain close together but can move past each other; in gases, particles are widely separated and move freely. Melting, freezing, boiling and condensing involve changes in energy and particle arrangement. The temperature at which a substance changes state depends on the strength of the forces between its particles. State symbols used in equations are (s), (l), (g) and (aq).

Giant covalent structures contain very large networks of atoms connected by strong covalent bonds. Examples include diamond, graphite and silicon dioxide. These substances generally have very high melting points because many strong covalent bonds must be broken to melt them. Their other properties depend on how the atoms are arranged and whether any electrons are free to move through the structure.

Diamond and graphite are different forms of carbon with different structures and properties. In diamond, each carbon atom forms four covalent bonds in a rigid three-dimensional giant structure. This makes diamond extremely hard and gives it a very high melting point, but it does not conduct electricity because it has no delocalised electrons. In graphite, each carbon atom forms three covalent bonds, producing layers of hexagonal rings. Weak forces between the layers allow them to slide, making graphite soft. Each carbon also contributes one delocalised electron, explaining why graphite can conduct electricity.

Graphene is a single layer of graphite consisting of carbon atoms arranged in a hexagonal structure. It is extremely strong, very thin and a good conductor of electricity, making it useful in electronics and composite materials. Fullerenes are molecules made entirely of carbon with hollow structures, including spherical molecules such as Buckminsterfullerene, C60. Carbon nanotubes are cylindrical fullerenes with very high length-to-diameter ratios. Their properties make them useful in nanotechnology, electronics and materials.

Nanoparticles generally have dimensions between 1 nm and 100 nm. Their properties can differ from those of the same material in bulk because nanoparticles have a much larger surface area to volume ratio. As particle size decreases, the proportion of atoms at the surface increases. This means smaller quantities of a nanoparticle material may sometimes produce the same effect as larger quantities of the bulk material. Students should compare nanoscale dimensions with atoms, molecules and larger particles and understand how surface area to volume ratio changes with particle size.

Nanoparticles have applications in medicine, electronics, cosmetics, sun creams, deodorants and catalysts. Their high surface area to volume ratio can make them particularly useful where high reactivity or a large active surface is required. However, nanoparticles may behave differently from larger particles of the same substance, and their effects on human health and the environment are not always fully understood. Students should be able to evaluate the advantages and disadvantages of using nanoparticles for a specified purpose and explain possible risks.