4.4.1: Bonding of Oxygen and Fluorine

Let us now consider oxygen (O) which has eight electrons, two in the core and six valence (1s², 2s², 2px², 2py¹, 2pz¹). As with nitrogen, oxygen does not use all its electrons to form six bonds because it is too small and the orbitals that would need to be used to make six bonds are too high in energy to be energetically accessible; that is, not enough energy would be released upon bond formation to “pay for” that energy.

The simplest oxygen-containing molecule is molecular oxygen, O₂. On our simple covalent bond model the two oxygen atoms are connected by a σ and a π bond, forming a double bond⁷⁷. The next simplest, stable, most common, and by far the most important compound of oxygen at least from the perspective of living organisms, is water (H₂O). In water there are two O–H bonds and two lone pair non-bonding orbitals. As in the case of nitrogen, the orbitals are sp³ hybrids and the oxygen atom is surrounded by four centers of electron density (see a pattern here?), two bonds, and two lone pairs. Again, the lone pair orbitals are larger than the O–H bonding orbitals, which distorts the tetrahedral symmetry of the molecule. Instead of equal angles of 109º between the orbitals, the angle between the O–H bonds is 104.5º. When we use a Lewis structure to represent the structure of H₂O, it is critical to include all valence shell electrons.

Compound	Molar mass (g/mole)	Boiling point	Bond type	Bond length (pm)	Atomic radius (pm)
CH4	16	–161 °C	C–H (in CH4)	109	C - 70
NH3	17	– 33 °C	N–H in (NH3)	101	N - 65
H2O	18	100 °C	O–H (in H2O)	96	O - 60
HF	20	19.5 °C	F–H in (HF)	92	F - 50
Ne	20	–246.08°C	not applicable	not applicable	Ne - 38

Continuing on across the periodic table we see that fluorine is the next element after oxygen. It has nine electrons: two core and seven valence. Rather than forming seven bonds fluorine only forms a single bond for basically the same reasons that oxygen only forms two bonds. Hydrogen fluoride, HF, has one bond, but four centers of electron density around the fluorine. Because HF has only two atoms, they must by definition lie on a line and therefore we do not need to discuss its shape.

As we will see, a valid Lewis structure makes it possible to extrapolate a significant amount of information about a molecule’s chemical and physical properties. A confusing point is that the Lewis structure can be written in a number of apparently different ways, which are actually equivalent. The key to remember is that the Lewis structure does not attempt to depict a molecule’s actual three-dimensional structure. It is a shorthand (a “cartoon” if you like) that assumes you already know the arrangement of orbitals. No matter how it is drawn, the actual structure of a H₂O molecule is the same with a 104.5º bond angle between the O–H bonds.

The tendency to form four centers (bonds or non-bonding pairs) has led to the rather misleading “octet rule”, which states that some elements tend to form molecules that have eight electrons around any atom (except for hydrogen). Unfortunately, the octet rule is far from being a rule because there are many exceptions, as we will see later. For example many of the elements past the second row of the periodic table are capable of bonding to more than four other atoms and some elements form stable compounds with less than eight electrons. It is important to remember that the octet rule is not the reason why atoms bond with each other, but it is a useful heuristic when constructing Lewis structures for the second row elements (C, N, O, F).

CH4	NH3	H2O	HF	Ne
-258.7°F (-161.5°C)	-28.01°F (-33.34°C)	212°F (100°C)	67.1°F (19.5°C)	-410.9°F (-246.1°C)

References

⁷⁷ Interestingly O₂ cannot be well described by a simple valence bond model, because it can be shown that molecular oxygen has two unpaired electrons (it is a di-radical). The bonding is best explained by using molecular orbital theory.