ENERGY BALANCE

Energy Input to the Body is the Sum of Energy Output and Storage

The first law of thermodynamics states that energy can neither be created nor destroyed; in a closed system, total energy is constant. This concept is illustrated in Figure 57-1. Humans acquire almost all their energy from ingested food, store it in different forms, and expend it in different ways. In the steady state, energy uptake must equal energy output.

The gastrointestinal (GI) tract breaks down ingested carbohydrates, fats, and proteins into smaller components and then absorbs them into the bloodstream for transport to sites of metabolism. For example, the GI tract reduces ingested carbohydrates to simple sugars (e.g., glucose), which are then transported to muscle cells and either oxidized to release energy or converted to glycogen for storage. Oxidation of fuels not only generates free energy but also waste products and heat (thermal energy).

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Figure 57-1 Energy balance.

The body's energy inputs must balance the sum of its energy outputs and the energy stored. When the

body takes in more energy than it expends, the person is in positive energy balance and will gain weight. Healthy children are in positive energy balance during growth periods. Conversely, when energy intake is

less than expenditure, this negative energy balance leads to weight loss. A person in negative energy balance preferentially uses stored fat as fuel.

A person can gain or lose weight by manipulating energy intake or output. An optimal strategy to encourage weight loss involves both increasing energy output as well as reducing energy intake. In most people, a substantial decrease in energy intake alone leads to inadequate nutrient intake, which can compromise bodily function.

Because of the Inefficiency of Chemical Reactions, the Body's Total Free Energy Falls

The second law of thermodynamics states that chemical transformations always result in a loss of free energy. The total internal energy (E) of the human body is the sum of the disposable or "good" energy (i.e.,

the Gibbs free energy, G) plus the unavailable or "wasted" energy (i.e., the product of absolute temperature, T, and entropy, S):

Imagine that you eat some glucose, which increases your total internal energy by a small amount (∆E).Some of this energy will be stored as glycogen (∆G), and some will waste as heat (T · ∆S). According toEquation 1, as long as the temperature is constant, the change in total internal energy will have twocomponents:

Thus, some of the increased total energy (∆E) will be stored as glycogen (∆G). However, because of theinefficiencies of the chemical reactions that convert glucose to glycogen, some of the ∆E is wasted as heat(T · ∆S). Another way of stating the second law is that T · ∆S can never be zero or negative, and chemicalreactions can never be 100% efficient.

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If we add no energy to the body (i.e., ∆E is zero), the body's total free energy must decline (i.e., ∆G isnegative). This decline in G matches precisely the rise in T · S, reflecting inefficiencies inherent in chemicaltransformations. Consider, for example, what would happen if you took 1 mole of glucose (180 g), put it intoa bomb calorimeter with O2, and completely burned the glucose to CO2 and H2O. This combustion would

yield 686 kcal in the form of heat but conserve no usable energy. Now consider what happens if your body burns this same 1 mole of glucose. In contrast to the bomb calorimeter, your mitochondria not only oxidizeglucose to CO2 and H2O but also conserve part of the free energy in the form of adenosine triphosphate

(ATP). Each of the many chemical-conversion steps from glucose to CO2 and H2O makes available a small

amount of the total energy contained in glucose. Converting 1 mole of adenosine di-phosphate (ADP) and inorganic phosphate (Pi) to 1 mole of ATP under the conditions prevailing in a cell consumes approximately

11.5 kcal/mole. Therefore, if a particular step in glucose oxidation releases at least 11.5 kcal/mole, it can be coupled to ATP synthesis. The conversion of the lower-energy ADP to the higher-energy ATP traps energy in the system, conserving it for later use. The cellular oxidation of 1 mole of glucose conserves approximately 400 kcal of the potential 686 kcal/mole; the remaining 286 kcal/mole are liberated as heat. Because the body cannot reconvert this heat to energy, the body's total free energy falls.

Free Energy, Conserved as High-Energy Bonds in ATP, is Released in a Controlled Manner to Provide the

Energy for Cellular Functions

ATP consists of a nitrogenous ring (adenine), a five-carbon sugar (ribose), and three phosphate groups (Fig. 57-2). The last two phosphates are connected to the rest of the molecule by high-energy bonds. The same is true for a related nucleotide, guanosine triphosphate (GTP). If we compare the free energies of phosphate bonds of various molecules, we see that the "high-energy" phosphate bonds of ATP lie toward the middle of the free-energy scale. Thus, in the presence of Pi, ADP can accept energy from compounds

that are higher on the free-energy scale (e.g., phosphocreatine), whereas ATP can release energy in the formation of compounds that are lower on the free-energy scale (e.g., glucose-6-phosphate [G6P]). Thus, ATP can store energy derived from energy-releasing reactions and release energy needed to drive other chemical reactions.

Figure 57-2 Hydrolysis of ATP to ADP, inorganic phosphate, and H+. ADP, adenosine diphosphate; ATP, adenosine triphosphate.

Examples of the chemical reactions fueled by converting ATP to ADP and Pi include the formation of

peptide bonds in protein synthesis, bridge formation between actin and myosin during muscle contraction,

and pumping Ca2+ against its concentration gradient during muscle relaxation. Although the conversion of

ATP to ADP and Pi is often referred to as a "hydrolysis reaction," because the traditional representation is

ATP + H2O → ADP + Pi + H+, the reaction actually occurs in two steps. The first typically involves transfer

of a part of the phosphate group to an intermediate molecule, thus increasing the intermediate's free-energy content. The second step involves displacing the phosphate moiety, releasing Pi and energy. In contrast,

true hydrolysis reactions merely release heat, which cannot be trapped to drive chemical processes.

Printed from STUDENT CONSULT: Medical Physiology (on 28 August 2006)© 2006 Elsevier