In the real world, simple chemical reaction systems are rare. While chemistry lab experiments typically involve mixing pure chemicals together in well defined amounts under tightly controlled conditions, in the wild, things are messier. There are usually a number of chemical species present, and that can lead to competing reactions. While laboratory systems are effectively “closed”, and the results are analyzed only after the reaction has reached equilibrium, real systems are usually open and rarely reach equilibrium. This is particularly true for living systems - which tend to be dead if they reach equilibrium (or become enclosed). In fact most real systems are subject to frequent short and long term perturbations. For example, as we saw in the last chapter, perturbations to a system (adding or taking away a product or a reactant) lead to compensatory changes - the system responds, as described by Le Chatelier’s principle. This essentially simple behavior can, in the context of more complex system, produce quite dramatic behaviors. Life is an example of such a system, a system that has (in its various forms) survived uninterrupted for over 3.5 x 10⁹ years.

In this chapter we will examine a range of complex systems. We will consider how living systems manage to keep the concentration of important chemical species at a reasonable level (for example by buffering the pH), how they manage to use differences in concentrations of chemical species to drive cellular processes (like thought), and how reactions that release energy (by forming more stable compounds) can be coupled to reactions that require energy to occur.

Systems composed of one reaction

We begin by considering a few of the reactions that are important to us and that can either be moved backwards or forwards depending on conditions. Molecular oxygen (O₂) is a vital component in a number of reactions in our bodies, particularly energy capture from food, through an evolutionarily ancient process known as aerobic respiration. O₂ must be transported to every cell so that it can participate in cellular reactions. O₂ diffuses into the bloodstream in the lungs, but O₂ is not very soluble in water (the main component of blood). If we relied on O₂‘s solubility in water to transport it around the body we would be in trouble. Instead O₂ reacts with (we usually say binds to, but this is definitely a chemical reaction) a protein called hemoglobin. The structure of hemoglobin is complex; it is composed of four polypeptide subunits, each polypeptide is associated with a heme group. The heme group contains an iron ion (Fe2+) complexed to four nitrogenous bases? linked into a ring (called a porphyrin) to form a more or less planar arrangement. A similar molecule, known as chlorophyll, differs most dramatically from heme in that the iron ion is replaced by magnesium (see below). Its function is not to bind O₂ (or CO₂), but rather to absorb visible light and release an energetic electron as part of the photosynthetic process.

Iron is a transition metal. Recall that these elements have d orbitals, some of which are empty and available for bonding. Iron II (Fe2+) has plenty of energetically available orbitals, and therefore can form Lewis acid-base complexes with compounds that have available electrons (such as nitrogenous bases). Within the porphyrin ring, four N’s interact with the Fe2+ ion. Typically transition metals form complexes that are octahedral in geometry. In the case of the heme group, four of these interactions involve nitrogens from the four rings, a fifth involves an N of histidine residue of one of the protein’s polypeptides that approaches from below the ring plane. This leaves one site open for the binding of an O2 molecule, which has available lone electron pairs.190 When an O2 binds to one of these heme groups, Hemoglobin + O2 ↔ Hemoglobin-O2.

This way of depicting the reaction is an oversimplification. As we said initially, each hemoglobin molecule contains four polypeptides each of which is associated with a heme group (green in the figure), so there are four heme groups in a single hemoglobin molecule. Each heme group can bind one O₂ molecule. As an O₂molecule binds to the heme iron, there are structural and electronic changes that take place within the protein as a whole - this leads to a process known as cooperativity - the four heme groups do not act independently.

Binding of O2 to one of the four heme groups in hemoglobin causes structural changes to the protein, which increases the affinity of remaining three heme groups for O2. Similarly, when a second O2 binds, it again increases the affinity for O2 of the remaining two heme groups.

As you might suspect, this process is reversible. Imagine a hemoglobin protein with four bound O₂s. When an O₂ is released from the hemoglobin molecule, the affinity between the remaining O2s and the heme groups is reduced, making it more likely that more of the bound O₂s will be released. This is an equilibrium reaction, and we can apply Le Chatelier’s principle to it. Where O₂ is present in high abundance (in the lungs) the reaction shifts to the right (binding and increasing affinity for O₂), and where O₂ is present at low levels, the reaction shifts to the left, releasing O₂ and reduced affinity for O₂.) The end result is a molecule (hemoglobin) that has high capacity for binding O₂ where O₂ is present at high concentrations and readily releases O₂ where O₂ is present at low concentrations. In the blood [hemoglobin] ranges between 135-170 g/L, approximately 2 mM, and since there are 4 O₂ binding sites per hemoglobin, this leads to about ~8 mM concentration of O₂ - this compares to O₂2s’s solubility in water which is ~ 8mg/L at 37ºC, or 250 μM). The reaction can be written like this

O₂ in the air ↔ O₂ in the blood (liquid) + hemoglobin ↔ hemoglobin-O₂ + O₂ in the blood ↔ hemoglobin-2O₂ + O₂ in the blood ↔ hemoglobin-3O₂ + O₂ in the blood ↔ hemoglobin-4O₂

When the hemoglobin reaches areas of the body where [O₂] is low, the oxygen dissociates from the hemoglobin into the blood. The dissolved O₂ is then removed from the blood by aerobic (that is, oxygen-utilizing) respiration.   C₆H₁₂O₆ + 6O₂ ↔ 6CO₂ + 6H₂O

Again, the combination of Le Chatelier’s principle and the cooperativity of the O₂ + hemoglobin reaction now leads to the complete dissociation of the hemoglobin-4O₂ complex, releasing O₂. The products of aerobic respiration - which is essentially a combustion reaction - are carbon dioxide and water.

Question to answer:

• What complicates reaction systems in the real world (outside the lab)?
• Why is O₂ not very soluble in water?
• By what factor does binding with hemoglobin increase solubility of O₂ in water?
• Draw Lewis structures for O₂ and CO, why do you think they bind in similar ways to hemoglobin?
• Why would you expect CO₂ to react with hemoglobin differently from the way O₂ interacts with hemoglobin?
  

Questions to ponder:

• Why does it make physiological sense that O₂ would bind to oxygen-free hemoglobin (deoxyhemoglobin) relatively weakly and cooperatively?