Topic B: Microscopic World I

Microscopic World I explains the behaviour of matter in terms of atoms, their electron arrangements, and the way atoms join together. It builds the periodic table from electron structure and introduces ionic, covalent and metallic bonding, then links each type of structure to the physical properties we can observe.

Atomic structure and isotopes

An atom has a tiny central nucleus containing positively charged protons and neutral neutrons, surrounded by negatively charged electrons in shells. The atomic number is the number of protons and identifies the element; the mass number is the number of protons plus neutrons. A neutral atom has equal numbers of protons and electrons. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they have the same chemical properties but different masses. Relative atomic mass is the weighted average mass of an element's atoms compared with one twelfth of a carbon-12 atom, taking the abundance of each isotope into account, which is why chlorine has a relative atomic mass of about 35.5.

Electron arrangement and the periodic table

Electrons occupy shells around the nucleus, with the first shell holding up to 2 electrons, the second up to 8, and the third up to 8 for the first twenty elements. The electron arrangement is written as numbers separated by commas, for example sodium as 2,8,1. The modern periodic table arranges elements in order of increasing atomic number. The group number equals the number of outer-shell (valence) electrons, which controls chemical behaviour, while the period number equals the number of occupied shells. Elements in the same group react similarly; for instance, Group I metals all have one outer electron and form ions with a single positive charge.

Metals, non-metals and noble gases

Metals are found on the left and centre of the periodic table; they tend to lose electrons to form positive ions and are good conductors. Non-metals lie on the right and tend to gain or share electrons. The noble gases of Group 0 have full outer shells (helium has 2, the others 8), which makes them very stable and chemically unreactive. The drive of other atoms to reach this stable noble-gas electron arrangement is the key idea behind bonding: atoms transfer or share electrons so that each achieves a full outer shell. This octet idea lets you predict the charges of common ions and the formulae of simple compounds.

Ionic bonding and ionic structures

Ionic bonding occurs between a metal and a non-metal: the metal atom loses one or more electrons to form a positive cation, and the non-metal gains them to form a negative anion, each reaching a noble-gas arrangement. The oppositely charged ions are held together by strong electrostatic attraction acting in all directions, so they build a regular three-dimensional giant ionic lattice, as in sodium chloride. Ionic compounds have high melting and boiling points because much energy is needed to break the many strong attractions. They are brittle, do not conduct when solid, but conduct when molten or dissolved because the ions become free to move. Many dissolve in water.

Covalent bonding and covalent structures

Covalent bonding occurs between non-metal atoms, which share pairs of electrons so each gains a stable outer shell. A single bond shares one pair, a double bond two pairs. Simple molecular substances, such as water, carbon dioxide and methane, consist of small molecules held to each other by weak intermolecular forces, so they have low melting and boiling points and do not conduct electricity. Giant covalent (network) structures, such as diamond, graphite and silicon dioxide, have atoms joined by strong covalent bonds throughout, giving very high melting points. Diamond is hard and non-conducting, while graphite has delocalised electrons that let it conduct and layers that slide, making it a good lubricant.

Metallic bonding and properties of metals

In a metal, the atoms lose their outer electrons to form a regular lattice of positive ions surrounded by a sea of delocalised electrons free to move throughout the structure. The strong attraction between the positive ions and the mobile electrons is the metallic bond. This model explains the typical properties of metals: they conduct electricity and heat because the delocalised electrons can move and carry charge or energy; they are malleable and ductile because layers of ions can slide over one another without breaking the bonding; and most have high melting points and a shiny appearance. Stronger metallic bonding, from more delocalised electrons or smaller ions, gives higher melting points.