![]() Usually, we define the z axis as lying along the line between the two atoms we are looking at. ![]() The p z on one atom could interact with the p z on the other atom, however, because they are parallel to each other. If they do not lie parallel to each other - that is, if they are perpendicular to each other, such as a p x and a p y - then they cannot interact with each other at all. Remember, there are a couple of very different ways in which p orbitals can combine with each other, depending upon which axis they lie. The 2s orbitals aren't the only ones in the second shell. So we can ignore them and we aren't really missing anything. The effect of both those combinations being occupied is to cancel out the bonding those two pairs of electrons remain non-bonding. When those four electrons are filled into the MO diagram from the bottom up, they will occupy both the bonding σ 1s and the antibonding σ 1s*. That must mean that each atom has two 1s electrons of its own, for a total of four. In the context of MO, suppose we do have 2s electrons. When we drew Lewis structures, we gave oxygen six electrons, rather than eight we were ignoring the core. Ignoring the core electrons is pretty common if you recall, in atomic electron configurations we might write 2s 22p 4 instead of 1s 22s 22p 4 for oxygen we were ignoring the core. They are buried a little deeper in the atom, and they don't play a very important role in bonding. That's because if there are any 2s electrons, then those 1s electrons are really core electrons, not valence. Most of the time, we aren't going to see both the σ 1s and the σ 2s displayed in the diagram. These spherical orbitals would combine very much like 1s orbitals, and we would get a similar diagram, only at a slightly higher energy level. The next lowest set of atomic orbitals is the 2s level. Note that we have not added any electrons to that molecular orbital energy diagram yet, but when we do, we will just fill them in from the bottom up, just like we would an atomic orbital energy diagram. Altogether, the picture says that the 1s orbital on one atom and the 1s orbital on the other atom can combine in two different ways, producing the lower-energy, bonding σ 1sand the higher-energy, antibonding σ 1s*. The sides of the diagram just refer back to where those molecular orbitals came from, with dotted lines to guide you from one place to another. The order of energy so far is σ 1s, σ 1s*. It is analogous to the atomic orbital energy diagram (which goes 1s, 2s, 2p, 3s.). The middle of the diagram is just the molecular orbital energy diagram. What we see here is a molecular orbital interaction diagram. ![]() So, in a molecule, the lowest-energy molecular orbitals would be the ones formed from the lowest-energy atomic orbitals, the 1s orbitals. There are some departures from that rule, sometimes, but that's the simplest place to start. To a great extent, the order of molecular orbitals in energy can be considered to follow from the order of the atomic orbitals from which they are constructed. Just as we think of there being a progression of atomic orbitals from lowest energy to highest (1s, 2s, 2p, 3s.), we can organize these molecular orbitals by order of their energy. So far, we have looked at the ways in which pairs of atomic orbitals could combine to form molecular orbitals - to form bonds. ) ofĪs with H 2 +, the He 2 + ion should be stable, but the He–He bond should be weaker and longer than in H 2.\)
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