8. Oxygen

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:

Each p-orbital lobe holds 1 electron. An electron pair occupies two opposite lobes.

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 indicates the boundary of the n=2 shell.)

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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. (This includes the inner and outer faces, as well as the sides of each orbital.) In the case of oxygen, 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.

Neutral oxygen (O) atom (left) compared to the larger oxide (O2–) ion (right)

Water (H2O)

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 others (like oils) to resist water and form cellular structures within a water environment. It is truly a remarkable molecule.

Formation of the water (H2O) molecule
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Dioxygen (O2)

The dioxygen (O2) molecule (usually just called “oxygen”) has a surprisingly strong paramagnetism, which means it is strongly attracted into a magnetic field (as shown in the video below). This effect implies that there must be unpaired electrons in the molecule. Unpaired electrons can align themselves with an external magnetic field, and this causes them to be attracted into it. However, the oxygen molecule seems, at first glance, to have no unpaired electrons, so its strong observed paramagnetism needs to be explained.

Liquid dioxygen (O2) poured between the poles of an electromagnet are attracted to it. (Source: Harvard Natural Sciences Lecture Demonstrations, Youtube)

The strength of this magnetic effect is called magnetic susceptibility (χm). By way of perspective, oxygen gas (O2(g)) has a magnetic susceptibility of χm=3,449, which is stronger than the rare earth metal cerium (χm=2,450). Liquid oxygen (O2(l)) has χm=7,699, which is stronger than the rare earth metal neodymium (χm=5,628) — which has four unpaired f-electrons. Furthermore, oxygen gas has more than six times the paramagnetism of manganese (χm=529), which has five unpaired d-electrons.

The Molecular Orbital Theory attributes the paramagnetism of dioxygen (O2) to the presence of two unpaired electrons in anti-bonding molecular orbitals (see Triplet Oxygen). An alternate (and at this point speculative) explanation for this phenomenon may emerge from the new theory of Sub-Quantum Chemistry, in which a set of quantum electron interactions unique to the oxygen molecule gives rise to a stronger than expected magnetic susceptibility (χm) value.

Combustion Reactions

Oxygen has a high electronegativity, which means that oxygen atoms pull electrons strongly toward themselves. It is difficult for an oxygen atom to do that, though, when it is bonding to another oxygen atom that is pulling electrons just as strongly. But when other atoms that hold their electrons less tightly are available, oxygen atoms will prefer to bond with them because their electron density will be more easily attracted toward the oxygen atoms. This is true of most atoms on the periodic table.

When combining with oxygen, we call the process combustion because it usually gives off so much heat that things catch fire. An example of this reaction occurs when a carbon and/or hydrogen compound reacts (or burns) in oxygen. When natural gas (methane CH4) burns, oxygen atoms from the oxygen (O2) molecules in the air let go of one another and instead bond to all available carbon and hydrogen atoms in the methane. Water vapor (H2O) and carbon dioxide (CO2) are formed as a result:

CH4 + 2O2 —> 2H2O + CO2

Other common combustion reactions occur in the formation of water (H2O) from hydrogen (H2) and oxygen (O2) and the combustion of octane (C8H18) in (the soon-to-be-obsolete) gasoline car engines.

Ozone (O3)

A different naturally occurring form (or allotrope) of oxygen is ozone (O3). It is formed in the stratosphere when incoming ultraviolet light hits oxygen molecules, and ozone is also broken back down into oxygen when hit by ultraviolet light. The ozone layer therefore protects us from harmful radiation by acting as a UV shield in two ways.

Ozone is a molecule that demonstrates an interesting electron state known as resonance. Resonance occurs when the electrons have more than one way to make the same structure, though the structure is not symmetrical overall.

The two equivalent Lewis Dot Structures for ozone (O3)

Both versions of this structure feature a single bond, a double bond, and 6 di-electrons. They are exact mirror images of one another, and the traditional understanding of this in chemistry is that the actual structure is some form of average between these two structures. Since that is difficult to depict, we show both alternatives and we imagine the resulting average.

In Sub-Quantum Chemistry, Lewis Dot Structure is extended in order to incorporate all of the quantum interactions between electrons. Quantum Lewis Structure allows us to depict a single symmetrical resonance state for ozone, and according to this new theory, it looks like this:

The Quantum Lewis Structure (left) and space-filling structure (right) for ozone (O3), according to Sub-Quantum Chemistry theory.

See Quantum Lewis Structure for more detail about ozone’s resonance structure.

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