linked hydrogen nuclei can be considered to have its shell filled with two electrons. Thus the two hydrogen nuclei are fixed together by an electron arrangement that provides the cohesive force between them. This arrangement makes the molecular orbital equivalent to the elec- tron configuration of the noble gas helium.

Oxygen, in the second row of the periodic table, also forms a di- atomic (O₂) molecule, but follows a different rule, the so-called octet rule. This means that the molecule must be surrounded by eight electrons in its outermost s and p orbitals to mimic the energetically favourable electron configuration of the noble gas that lies next to oxygen, neon (Ne). To form a diatomic stable molecule, the two oxy- gen nuclei need to share at least two pairs of electrons so that they can acquire the stability of an outer electron octet. In this case, each oxygen atom in the molecule also has two pairs of electrons not involved in the bonding. These electron pairs are called lone pairs (Fig. 1.9a). Another possibility is that the oxygen molecule share three pairs of electrons for covalent bonding (Fig. 1.9b). In this case, six electrons are not involved in the covalent bonding, and exist as two lone pairs and two single electrons. In this electronic structure, oxygen is a biradical with a high oxidative reactivity.

An ionic bond results from the electrical attraction between posi- tively charged cations and negatively charged anions. Cations are formed when elements capable of transferring some of their nega- tively charged electrons donate electrons to elements capable of accepting electrons, that form anions. This mechanism can be best demonstrated by the reaction between sodium (Na) and a counter element, chlorine (Cl). Sodium, with the atomic number 11, has 11 positively charged protons counterbalanced by 11 negatively charged electrons. Its valence shell holds only one electron. If sodium donates this electron to chlorine, which is able to accept this 'gift', it keeps 10 electrons in three shells, two in 1s, and eight in 2s and 2p. Sodium has then acquired the electron configuration of neon. However, as a re- sult the 11 protons are counterbalanced by only 10 electrons so that sodium is left with a net electrical charge of 1+. A positively charged cation Na⁺ is created, that alone is not stable but needs to combine with a partner counterbalancing its net charge. Chlorine, with atomic number 17, has 17 protons counterbalanced by 17 electrons. Its valence shell holds seven electrons and requires only one more to complete the stable octet. It can capture one electron from some other atom that is willing to release it, for example from a sodium atom (Fig. 1.10a). This electron transfer results in the formation of a chloride anion that carries a net charge of 1-. As a result of the electron transfer both ions are attracted to each other by a strong electrostatic force and the