In this section we will explore what happens when we combine two O atoms together to form O₂, the second most common molecule in Earth's atmosphere after N₂. The figure below shows a couple of the ways one might put the two atoms together.

25.1

Form 2 is perhaps the most obvious thing to try: orient the two atoms so that a sigma bond forms between two of the singly occupied 2p orbital and then form a pi bond between the other two singly occupied 2p orbitals. As the potential energy curves show, this gives a stable structure, but not the most stable one. In Form 1, the most stable structure, there is just a sigma bond. The two remaining singly occupied 2p orbitals both point off-axis with a 90° between them (making them π_x and π_y orbitals). Why would form 1 be more stable than form 2, when there appears to be more bonding in form 2?

We can see what's going on if we look at the orbitals at the minimum separation:

25.2

First, the top row shows the two O 2p orbitals that form the sigma bond. As we saw in N₂, each orbital is attracted to the other nucleus and delocalizes partly onto it. The second row shows the three electrons in π_x orbitals. There is a pair of electrons that looks very much like the pi bond pair in N₂: again, each 2p orbital is attracted toward the other nucleus. The difference between N₂ and O₂ is that O₂ has a third electron in another π_x orbital. It's in the same general space as the bond pair, so it's destabilizing what would otherwise be strong pi bonding. Note how the lobes of the singly occupied orbital are directed somewhat away from the bonding region between the nuclei to reduced their overlap with the bond pair. Bear in mind that there are also three electrons in π_y orbitals. They exhibit identical behavior.

We've seen the behavior that occurs when the a doubly occupied orbital interacts with a singly occupied orbital. Here it affects the π orbitals. Previously, (in Section 14) we observed what happens when the 1s orbital of H interacts with the lobe pair on Be. We saw a sigma recoupled pair bond form. In O₂ we are observing pi recoupled pair bonding. The following figure shows what happens:

25.3

In this diagram the off-axis 2p orbitals are split into large and small parts. The small parts each have an electron (four of the total six electrons in π_x and π_y orbitals), and they can be coupled together to form pi bonds. That leaves an electron in one of the large parts in both the x and y π orbitals. The diagram is more confusing than we'd like, however, so we won't use this form of diagram again.

Adding H to O₂. Next, we'll add H to the second form of O₂ and see what happens. Each O atom has a singly occupied orbital, so we can H to either O atom.

Adding H to O₂ is analogous to adding the second H to OH to form H₂O. The singly occupied O 2p orbitals in O₂ lie at a 90° angle from the O-O bond axis, so the nominal bond angle of HO₂ is also 90°. As shown in the figure below,

25.4

.... the actual bond angle is 104°, just as we found for H₂O. Let's look at some of the orbitals. When H approaches the O₂, we see that the orbitals that are affected by forming the new bond rearrange quite a bit:

25.5

The top row shows the orbitals for H and O₂ before they've changed very much, at R = R_e + 1 Å. From left to right, we see the H 1s¹ orbital, the singly occupied π orbital on O₂ and the doubly occupied π orbital on O₂. In the middle row (R = R_e + 0.4 Å), the OH bond is starting to form. The H orbital is being attracted toward O. The singly occupied O₂ orbital is starting to localize onto the O atom to which H is bonding, while the doubly occupied O₂ orbital is localizing onto the other O atom. The final row show the orbitals in HO₂. The two singly occupied orbitals on H and O are now coupled into an OH sigma bond pair, and the double occupied orbital is almost fully localized on the other O atom as a 2p² pair. The pi bond between the two O atoms is gone, because one of the orbitals has formed a bond pair with H.

Let's also examine the orbitals forming the O-O sigma bond:

25.6

The orbitals participating in the existing bond between the two O atoms change very little when the new OH bond forms.

Adding H to HO₂. There's still a singly occupied orbital on HO₂ that is perpendicular to the plane containing the three nuclei, so we expect we can add a second H to HO₂ to form H₂O₂.

We would expect the second H to form a bond with the O atom on the end, answer (c), not with either of the other two atoms. That would give us a structure where the atoms are connected as H–O–O–H. The actual structure of H₂O₂ (HOOH) is shown in the next figure, along with its 2D and 3D bonding diagrams.

25.7

Like HO₂, the nominal bond angle of the second OH bond is also 90° In the optimized structure of H₂O₂, both OH angles are 100°. There's another angle of interest. We see that H₂O₂ is not planar. The two H atoms do not lie in the same plane as the two O atoms. As shown below, there are two planes that be drawn:

25.8

The green line is the O-O bond axis, and both O atoms lie in both planes. The plane outlined in orange contains the two O atoms and the upper H atom, while the plane outlined in light blue contains the two O atoms and the other H atom. The angle between the two planes is 112°. We call this the torsion angle. Based on the bonding diagrams shown for H₂O₂, we expect its torsion angle to be close to 90° because that is the angle between the two off-axis singly occupied 2p orbitals in O₂. The real angle is 22° larger. This compound, H₂O₂, is known as hydrogen peroxide. HO₂ is known as the peroxyl radical.

Adding H to HO₂, part 2. In the last question we asked you which atom the second H atom forms a bond with. We found that H forms a bond with the O atom on the end (the terminal O atom). But if we try forming a bond with the central O, we find that there is another stable compound that can be formed. Here's is what it looks like:

25.9

We will label this compound H₂OO to distinguish it from the other form. It is sometimes known as water oxide or oxywater because it consists of water with an extra O bonded to it. Oxywater is much less stable that hydrogen peroxide: note that the bond energy for the second OH bond is only 48 kcal/mol for oxywater while it is 93 kcal/mol for H₂O₂, so the latter is nearly twice as large as the former.

The orbitals depicted in Figure 25.9 account both for why H₂OO can form at all and why it is so much less stable that H₂O₂. When the first OH bond forms, the out-of-plane π orbitals shift. The lobes are no longer evenly distributed on the two O atoms as they are in O₂. The pair of electrons shifts toward the O that is between the H and other O, while the the singly occupied π orbitals has more density on the terminal O atom. So one does expect the second H to preferential form a bond with the terminal O atom, but the singly occupied orbitals has enough density left on the orbital that the H can add there instead. But more rearrange ment is needed, and thus it is the less stable compound. Incidently, when two compounds can be formed with the same atoms (two H and two O atoms, in this case), they are called isomers.

There's another way to view the bonding in H₂OO, as a dative bond between water and the second O atom. When we first encountered dative bonding back in Section 11, we observed a bond form between He and H⁺ (see Figure 11.3 and following). The shared pair of electrons came from H, since H⁺ has no electrons, only empty orbitals. To form a dative bond, we need a pair of electrons and an unused orbital. Water has a 2s² lone pair, but how does O end up with an unoccupied orbital? We've usually seen it with two singly occupied 2p orbitals and one doubly occupied orbital. But way back in Section 7 we saw that there is an excited state of O atom that has two doubly occupied 2p orbitals and one unoccupied 2p orbital. That's the form of O atom that can form a bond with water. Here are the two forms of O atom again, ground and excited states:

25.10

We will see dative bonding with O again in the next section and will look at orbitals then.

Adding H to H₂O₂ or H₂OO. If we try to add a third H to either of these compounds, we find that we cannot form another OH bond.

In the next section, we will start with N and see what compounds we can make by adding O and H to it.