Each electron shell contains electrons up to a fixed maximum. Atoms with unfilled electron shells are less stable than those with full shells. Atoms of the noble gases all have filled electron shells- this makes them very stable (inert), and they rarely react to form molecules. All other elements have unfilled electron shells, so they undergo the process of bonding, in which outer electrons are shared or transferred between atoms, so that each atom has a stable, full, outer shell.
Bond length: the distance between two atomic nuclei when the atoms are linked by a chemical bond.
Covalent bond: a pair of shared electrons between two atoms- a balance between the attraction between positively charge atomic nuclei and negatively charged electrons, and the repulsion between nuclei/electrons and themselves. Each atom donates one electron to the shared pair, resulting in an area of high electron density that pulls the two nuclei together, thus forming a molecule.
e.g. a hydrogen atom has one electron in its outer (and only shell). It can become more stable by having a filled outer shell containing two electrons. If two hydrogen atoms share their electrons in a covalent bond, each has “access” to the two electrons required for stability. In this way, a hydrogen molecule is formed.
A covalent bond in a graphical chemical formula is represented by a straight line. Double and triple bonds can occur between atoms, where two or three pairs of electrons are shared between two atoms, e.g. H-O-H, O=O N: N are graphical formulae indicating the covalent bonding in some small molecules.
Covalent/atomic radius: half the bond length when two atoms of the same element are linked by a single covalent bond.
Dative covalent/coordinate bonding is where the two electrons forming a covalent bond are donated by only one atom. Dative covalent bonds are represented by arrows pointing from the atom that donates the electrons to the atom that accepts them, e.g. N® H
Dot-cross diagrams are a way of showing how electrons are shared in covalent molecules and ionic compounds. The dots and crosses represent the outer electrons only.
There are two golden rules in dot cross diagrams (or at least there were for the AEB Chemistry A-Level examiners). First only use dots and crosses- even if you think squares and other symbols will make things clearer, don’t use them. Second, don’t draw in any covalent bonds. There’s nothing to stop you doing a rough graphical formula showing all the covalent bonds (in fact, this could be very helpful) but don’t put any bonds in the actual dot-cross diagram.
Note: if you’re drawing a dot-cross diagram of a charged species, put square brackets around the diagram and write the charge outside the bracket on the top right hand side.
Giant covalent structures have strong covalent bonds throughout a huge molecule containing thousands of atoms linked together, e.g. diamond, graphite, silicon dioxide. The number and strength of bonds within the molecule leads to a non volatile structure with high melting and boiling point. Giant structures are also hard, and often brittle. They are insoluble in all solvents. All giant structures except graphite are non conductors because they have no free electrons available to carry a current. Graphite can conduct because it has one free bonding electron per atom.
Simple/small molecules: molecules consisting of a small number of atoms linked together by covalent bonds, e.g. S8 HCl, CH4. Small molecules have strong intramolecular bonds (within a molecule) but weak intermolecular forces (between molecules). This makes them volatile with low melting and boiling points as these weak intermolecular forces are easily overcome. Simple molecular structures are generally gases or liquids at room temperature. They are non-conductors- electrons cannot be transferred between molecules. Generally, small molecules are soluble in non-polar solvents (e.g. CCl4) but insoluble in polar solvents. An explanation of this is given in the Intermolecular forces section.
Ionic Bond: the electrostatic force of attraction between two oppositely charged ions. Electrons are transferred from metal atoms to non-metal atoms so that each becomes an ion with a full outer shell. The ions are now oppositely charged, and so attract each other.
Ionic Lattice: the 3D structure of ions in a crystal, consisting of a regular arrangement of alternating metal and non-metal ions, held together by the electrostatic forces of attraction between oppositely charged ions. Ionic lattices form a giant, repeating structure.
Ionic radius: half the radius of an ion in a crystal.
Properties of ionic compounds:
- Conduct electricity when molten or dissolved in water as under these conditions the ions are free to move and carry a current. Ionic compounds do not conduct electricity when solid as the ions are in fixed positions.
- Dissolve in polar solvents (see Intermolecular forces)
- Ionic compounds are brittle. This is because applying a force can disrupt the alternating cation-anion lattice arrangement. If the layers in the ionic crystal are disrupted, ions with same charge may end up next to each other- these ions will repel each other and a crack will form.