The reaction that turns molecular oxygen (O2) into water releases lots of energy, and all animals need that energy to drive their bodies functioning. The half reaction and associated free energy change are:

O2 + 4H3O+ + 4e- --> 6H2O     delta G = -305 kJ/mol

There has to be a biological mechanism for capturing oxygen as O2 (in its high-energy, zero oxidation state) and bringing it to a place where it can be turned into H2O (this is a reduction reaction) in such a way that the energy released by this process can be profitably used by the organism. This is accomplished in most organisms via hemoglobin and/or myoglobin, which are often referred to as oxygen carrier proteins. An iron atom in the center of myoglobin binds to O2 and takes to where the energy is needed. Hemoglobin works almost exactly the same way except that where myoglobin has only one iron-containing protein subunit, hemoglobin has four. When one of hemoglobin's four irons binds an O2 molecule, the other three protein subunits' iron atoms can bind O2 more easily. This is called the "cooperativity effect" and is not completely understood. It is hypothesized that the binding of the first O2 acts as a trigger that stresses the protein chains into a configuration more amenable to accepting additional O2's.

Following are a few important characteristics of hemoglobin and myoglobin that make them so effective for life as we know it:

  • O2 binding is pH dependent. It is observed experimentally that O2 is released more readily when the pH is acidic; i.e. when there is lots of CO2 around. (Recall the reaction CO2 + H2O --> H2CO3.) This is important because it provides a means by which O2 is released only where it is needed. If there is lots of CO2 around, then the CO2 to O2 ratio probably needs some adjusting.

  • O2 is not a very good ligand for this system. Other ligands such as CO and CN- can bind more strongly to the iron because they are better pi acceptors than oxygen, and the fact that they can block the O2 binding site so effectively is what makes them so very toxic. However, it is critical for the body that O2 be only weakly bound because it needs to be able to pop on and off of the iron where it is needed. In other words, the reversibility of the reaction is important. While O2 has two electrons in its pi* orbital, CO and CN- have lots of empty, electron-accepting pi* space that strengthens their bonds to iron. Note that although CO and CN- are isoelectronic (meaning that their MO configuration is the same) the negative charge on CN makes it a slightly worse electron acceptor than CO and therefore CO binds more strongly to the iron than CN-. Paradoxically, CN- is more toxic per mole than CO. This is probably because it catalyzes some other configuration change in the protein chain that permanently deactivates the heme unit.

Here is a picture of the grand hemoglobin molecule itself. The iron atom is where all the O2 binding action happens, but notice that it is actually a very small part of the molecule as a whole.

Now consider the molecular orbitals in hemoglobin. The MO's are different when oxygen is bound and when it is not bound, and this accounts for the color change: hemoglobin is red when oxygen is bound and blue when oxygen is not bound. Not coincidentally, this is why your blood is red in your arteries and blue in your veins. To understand this, it is important to think about the geometry of the molecule. When oxygen is not bound, the iron atom is in the +2 oxidation state. It's slightly too big to fit into the hole in the center of the plane of the immediately-surrounding "heme," pictured below, so it rests just on top of the heme plane.

I looked through a bunch of textbooks for a picture of iron in a heme plane but at first it was difficult to find one; all I could find was a picture of a synthetic model of a system that is extremely similar. (Except that the iron atom fits all the way into the plane instead of resting on top of it in this picture.) The synthetic system is called a "picket fence" porphyrin and was made by none other than Professor Collman.

In the biological hemoglobin molecule, there's a nitrogen atom underneath the heme plane, but without oxygen bound the iron is prevented from being able to be approximated as octahedral by the fact that there's nothing on top. See the right side of the figure below for iron's atomic orbital structure in this unbound configuration. You can rationalize it using simple crystal field theory: xz and yz are approximately degenerate and are the lowest in energy because they are the furthest removed from the ligands. The xy orbital is next because although it is in the plane of the heme, the lobes of that orbital do not point toward the nitrogen lone pairs. Then comes z2; one of its lobes points toward the nitrogen on the bottom and its doughnut is in the plane of the heme nitrogens, but its top lobe points toward nothing. Last and highest in energy is the x2-y2 orbital, whose lobes point directly toward the lone pair electrons of the heme nitrogens.

Of course, when O2 binds, the Fe and O2 orbitals bond and their energies change. In particular, the O2 pi* orbitals interact with the Fe xz and z2 orbitals. Note that the same would be true for the interaction between the heme Fe and CN- or CO, only the interaction in either of those two cases would be stronger; CN- and CO are even more enthusiastic pi* acceptors because their pi* orbitals start out empty. A side effect of the pi* accepting properties of all three of these ligands is that iron formally changes from being in the +2 oxidation state to the +3 oxidation state. You can think of it as having "given up" one of its electrons to the ligand via the ligand's pi* orbitals. As a side effect of losing one of its outer electrons, the iron atom is now small enough to fit snugly into the hole in the center of the heme plane.