Oxygen has four electrons in its p-orbital. This is not a sphere-shaped harmonic, and so four p-electrons cannot achieve a stable electron symmetry around a spherical core. Oxygen therefore cannot simply add its four p-electrons on top of the same (2s2) configuration that beryllium has, as shown here, because it would not be a stable configuration:
The asymmetry therefore causes oxygen to hybridize its 2s and 2p electrons in order to achieve tetrahedral symmetry. Its sp3 hybrid orbitals feature two di-electrons and 2 unpaired, degenerate electrons. This is why oxygen typically makes 2 bonds, and it is this tetrahedral geometry that gives the water (H2O) molecule its characteristic bent shape (see below). (The wireframe simply indicates the boundary of the n=2 shell, since there are no electrons defining the boundary of its sphere.)
The outer spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital and will decrease toward the nodal regions that divide the orbitals, where electron density will be lowest. In this case the orbitals containing the two lone pairs (dark blue) will each occupy slightly more volume than the two containing the unpaired electrons.
Oxygen is keen to obtain two extra electrons to fill its 2nd shell. It is the second most electronegative element after fluorine, and can therefore be quite reactive as it attracts electrons to itself. Oxygen can make one or more covalent bonds or it can gain two electrons in ionic interaction in order to reach the stability of a full 2nd shell. This is the same electron configuration as the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why oxygen forms a 2– ion. The negative ion is much larger than the neutral atom because electrons now outnumber protons by two. This results in a much lower effective nuclear charge — a lower average attraction by the nucleus on each electron.
When oxygen gains two electrons via covalent bonding with two hydrogen atoms, the water (H2O) molecule is formed. Its asymmetrical structure, and the fact that oxygen pulls electrons more strongly than hydrogen, gives it very important properties. The most significant is that it creates an electron imbalance which makes the side of the molecule where the hydrogen atoms attach slightly positive (𝛿+) and the opposite (oxygen) side slightly negative (𝛿–). This polarity causes water molecules to stick to each other and to certain substances rather effectively. It makes it harder to change water’s temperature, causing oceans and lakes to moderate local climates. It makes ice less dense than water, causing ice to float, ensuring that aquatic species survive the winter. It allows a great variety of substances (like NaCl) to dissolve in water, and it allows (like oils) others to resist water and form cellular structures within a water environment. It is truly a remarkable molecule.
The dioxygen (O2) molecule has a surprisingly strong paramagnetism. This means it is strongly attracted into a magnetic field.
Gaseous dioxygen (O2(g)) has a magnetic susceptibility of χm=3,449, which is stronger than the rare earth metal cerium (χm=2,450). Liquid dioxygen (O2(l)) has a χm=7,699, which is stronger than the rare earth metal neodymium (χm=5,628). Furthermore, dioxygen gas has more than six times the paramagnetism of manganese, which has five unpaired electrons.
The Molecular Orbital Theory attributes the paramagnetism of dioxygen to the presence of two unpaired electrons in anti-bonding molecular orbitals. An alternate explanation for this phenomenon may emerge from the theory of Sub-Quantum Chemistry.
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