Living organisms carry out a diverse set of tasks, such as building and maintaining physical structures, moving, synthesizing macromolecules, maintaining electrochemical gradients, and maintaining a constant body temperature. All of these processes require energy. One fundamental problem living organisms face is how to obtain energy to carry out these tasks.

When considered at their most basic level, all of the activities mentioned above involve chemical reactions. Any chemical reaction can, in principle, proceed in one of two directions, in which the “reactants” are transformed into “products”, or vice versa. Whether or not a chemical reaction requires energy to proceed in a particular direction depends on the change in Gibbs free energy (ΔG, which equals Gproducts – Greactants) that takes place when the reaction occurs. The ΔG for a given reaction depends in part on the inherent characteristics of the reaction and the molecules involved, and in part on the molecules’ concentrations and environmental factors.

If ΔG for a reaction equals zero, then the reaction is said to be at equilibrium. At equilibrium, the forward and reverse reactions occur to the same extent and no net reaction takes place. If ΔG is negative, (i.e., <0), the forward reaction is more energetically favourable, or spontaneous. If all molecules begin at equal concentration, there will be a net conversion of reactants to products without addition of energy. This will continue until equilibrium is reached, at which time the concentration of products will be higher than that of reactants.

However, if ΔG for a reaction is greater than zero, the forward reaction is energetically unfavourable, and at equilibrium reactants have a higher concentration than products. To take a specific example, the reaction of combining glucose and fructose to make sucrose has a standard free energy change (ΔG°) of +5.5 kcal/mole. (Note: the small circle in ΔG° indicates that the value is calculated using the G for each molecule under “standard” conditions, with defined temperature, pressure, and concentration. The ΔG° incorporates information about only the natures of the molecules involved, and not about the actual conditions under which the reaction occurs in a particular system. Even though the standard conditions are not normally met in biological systems, they are useful as a point of reference.) If sucrose is to be made in appreciable quantities, the natural tendency of the glucose and fructose molecules to remain apart must be overcome.

One strategy to make products even when the ΔG° is positive is to change the concentrations of the reactants and products, such that the ratio of reactants to products is high. This is frequently done in living systems, but is not always practical or desired.

An alternate strategy to drive an unfavourable reaction forward is to couple that reaction to an energetically favourable reaction (i.e., one with a negative ΔG), such that the total ΔG for both reactions is negative. If the total ΔG is negative, then both reactions will proceed spontaneously. Knowing this, the question becomes: is there a suitable energetically favourable reaction that organisms can couple to many different unfavourable reactions? If so, where do the reactants for the energetically favourable reaction come from?

The vast majority of energy used to sustain life comes ultimately from the sun. Light energy is captured by plants and stored in carbohydrates (this process, photosynthesis, is not covered in this course). The plants can then use cellular respiration to break down the carbohydrates, releasing energy that the plants use to carry out processes necessary for life. Animals likewise use respiration to derive energy from sugars, fats, and proteins obtained by eating plants and other animals.

However, the breakdown of (for example) sugar molecules like glucose is not used to directly power chemical reactions in the cell. A major reason for this is that a lot more energy is available in a glucose molecule than is needed to drive individual chemical reactions in the cell. Most of the energy would be wasted as heat. It is more desirable to convert the energy available in glucose into a form that can be used in small amounts, a little at a time, as needed. The universal energy carrier in biological systems is adenosine triphosphate.

Hydrolysis of ATP to form ADP and inorganic phosphate (Pi) is energetically favourable (ΔG° of -7.3 kcal/mol). Combining the hydrolysis of ATP with unfavourable reactions such as synthesis of sucrose from glucose and fructose can result in an overall reaction that is favourable. Another example of an energetically unfavourable reaction that is coupled to ATP hydrolysis is the synthesis of glutamine from glutamic acid and ammonia.

ATP is often said to have a “high-energy” phosphoanhydride bond. This is a shorthand way of saying that the ATP to ADP + Pi hydrolysis reaction is energetically favourable. There is nothing special about the particular oxygen-phosphorus bond that is broken in this reaction that makes it “high-energy”, except for its surroundings. In this course, I will use the shorthand phrase “high-energy bond” because it is convenient, but remember that it means that an energetically favourable reaction takes place at that bond.

ATP can also be hydrolyzed to AMP and pyrophosphate (PPi). The PPi ion can be further hydrolyzed to release more energy. Other compounds with “high-energy” bonds exist, some of which release more energy than ATP when hydrolyzed. One example is creatine phosphate, which is used to quickly transfer phosphate groups to ADP during muscle contraction.

Hydrolysis of ATP drives a wide variety of energetically unfavourable cellular reactions. To have enough readily available energy to do work, the cell must continuously regenerate the ATP consumed.

In addition to requiring energy, some biochemical processes require molecules to be reduced. Chemically speaking, reduction is the gain of electrons. Oxidation is the opposite of reduction, and happens when a molecule loses electrons. In any red-ox reaction, at least one molecule is oxidized and at least one is reduced; there cannot be oxidation of one molecule without a corresponding reduction of another. So cells require a source of electrons for use when a molecule needs to be reduced.

In organic chemistry a carbon atom is said to become more oxidized as more electronegative atoms (such as oxygen) bond to it.

Two key biological molecules involved in red-ox reactions are Nicotinamide Adenine Dinucleotide (NAD) and Nicotinamide Adenine Dinucleotide Phosphate (NADP). These molecules carry electrons on the nicotinamide moiety. In its oxidized form, nicotinamide carries a positive charge, and the molecules are abbreviated NAD+ or NADP+. Net addition of a hydride ion (H, one proton and two electrons) to nicotinamide neutralizes the charge, and the molecules are abbreviated NADH or NADPH. Similar to the way cells use ATP as an intermediate source of energy, they use NADH and NADPH as sources of reducing power (i.e., reducing agents or electron sources).

Although they are both reducing agents, NADH and NADPH have different roles in metabolism. The reducing power of NADH is primarily used to regenerate ATP from ADP, while the reducing power of NADPH is primarily used in the synthesis of biomolecules, notably fatty acids and cholesterol.

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