YOU CAN'T WIN. YOU CAN'T BREAK-EVEN. YOU can't get out of the GAME! (Part 4 of 5)
By Herve Bensabat

Free Energy

So much for a brief history of thermodynamics. Is there a simpler way of defining the laws of thermodynamics?

Generally speaking, in matters of nutrition only the first two laws are substantially meaningful. The zeroth law that establishes the concept of temperature and the third law that describes absolute zero do not pertain directly to our discussion.

Let us consider energy for a moment. There are two forms of energy that comprise the total energy of a system: Potential energy and Kinetic energy.

Potential energy is the energy an object has because of its position or condition rather than motion. An object held in a person's hand has potential energy, which turns to kinetic energy when the person lets it go, and it drops to the ground. So kinetic energy is the energy an object has because of its motion (equal to one half the mass of the body times the square of its speed). Releasing potential energy transforms it into kinetic energy of motion.

A macronutrient possesses potential energy before its release in metabolism. Its release often encompasses chemical reactions.

The term exergonic describes a chemical reaction that results in the release (freeing) of energy to its surroundings, i.e. one that gives up energy resulting in a decline of free energy (‘useful' energy available or ‘free' to do biologic work). If heat is the form of energy released, than the term exothermic is used (spontaneous reaction is another synonymous term).

Conversely, the term endergonic describes a chemical reaction that stores or absorbs energy from its surroundings resulting in an increase in free energy for biologic work. Similarly, if heat is the form of energy absorbed, the reaction is endothermic.

To honor Willard Gibbs (1839-1903), whose research provided the foundation for biochemical thermodynamics, the symbol G is used to denote free energy. Thus, to describe free energy quantitatively, the following formula is used:

G = H - TS

where, within a cell when pressure and volume are relatively stable, free energy (G) equals the potential energy within a molecule's chemical bonds (called enthalpy or H), minus the energy unavailable because of randomness ( entropy or S), times the absolute temperature T (degree C + 273).

Changes in free energy (ΔG) occur when the bonds in reactant molecules form new molecules with different bonding. Under conditions of constant pressure, volume and temperature, the equation to incorporate these changes can be expressed as follows:

ΔG = ΔH - TΔS

Consider Energy (E) for a moment. In natural events it is difficult to assess any change in energy (ΔE) because some work is done in the atmosphere.

Enthalpy (H) takes into account this work by adjusting ΔE for changes in pressure and volume.

ΔH = ΔE + P Δ V

where ΔH is the change in enthalpy, ΔE is the change in energy, P is pressure and ΔV is the change in volume.

Fortunately, volume changes at constant pressures are extremely small and for most biological reactions, P ΔV = 0. Hence,

ΔH = ΔE or ΔE = ΔH

Thus, we can now see more clearly how we arrived at equation (7) since, in effect, enthalpy reflects the change in energy, or, as it was defined earlier, it is the potential energy within a molecule's chemical bonds.

In the biological world, many important reactions are endergonic in that they require energy inputs. Often they are linked or coupled to, and driven by, exergonic reactions.

Usually when energy yielding (exergonic) and energy-consuming (endergonic) reactions are coupled, part of the energy available from exergonic processes is captured in the form of ATP. Endergonic reactions are catalyzed by enzymes that use the energy of ATP.

The first law of thermodynamics tells us that energy forms can be exchanged, but it doesn't say in what direction these exchanges occur. The second law tells us that these changes always go in the direction of randomness or disorder, i.e. when energy is exchanged, the efficiency will be imperfect and some energy will escape — generally as heat, thus increasing entropy (ΔS).

In essence, because of the second law, free energy (ΔG) carries a negative sign, that is energy is given up by a reaction and becomes available to do work. For example, in human muscle, 25% of the enthalpy change (ΔH) is available to do work with the remaining 75% used to maintain homeostasis and is wasted as heat, i.e. muscle efficiency during locomotion approximates 25%.

While humans can be said to be 25% efficient, machines don't always fare better. A gasoline engine for example, is 20-30% efficient. A diesel engine is 30-40% efficient and an ordinary light bulb is about 20% efficient.

From equation (7), if ΔG = ΔH - T ΔS, then rewriting the equation for ΔH yields:

ΔH = ΔG + T ΔS

Said another way, the transfer of potential energy to degrade to kinetic energy of motion always proceeds in a direction that decreases the capacity to perform work (i.e. increased entropy).

Ultimately, all of the potential energy in a system degrades to the unusable form of kinetic or heat energy.

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