Gatorade and Sports Beverages

In chemistry, the definition of energy in an open or closed system is the capacity to do work on the surroundings or supply heat to the surroundings.

Climbers, hikers, cyclists, and athletes of all ages and abilities are open systems that supply energy through chemical reactions. Everything else -- the air surrounding them, the ground upon which they walk or run, we call the surroundings. All of us transfer matter between system and surroundings as we inhale, perspire, exhale, and take a drink of water.

The Law of Conversation of Energy, derived from centuries of observation and measurement indicates that energy cannot be created or destroyed. But energy need not stay in one place. Energy can be converted from one form to another and can be created in one place and show up in another.

We would ask him to dribble and jump and shoot just one shot from the foul line. Let us ask him not to drink or eat anything. And for these few seconds he is also not to breathe. We wish for this demonstration, to make him a closed system. In theory, no matter is exchanged from this system. No oxygen taken in, no water or carbon dioxide exhaled.

Let us suppose further that we could measure in detail the energy released by Michael Jordan during the brief workout. We would find he had generated a certain amount of heat as he exercised. We could determine the work he had done as he ran and jumped and shot the ball. Michael Jordan, as a closed system, would have created energy through chemical reactions, and released that energy to his surroundings -- the empty arena.

Our measurements would show that the exact amount of energy released from the closed system had flowed into the surroundings. The arena would be slightly warmer from the heat from transferred to the cooler air during Jordan's exercise. The basketball floor of the arena and the ball in the arena would be at higher energy because of the work Jordan transferred as he bounced the ball and ran and jumped along the floor.

The requirements of the Law of Conservation of Energy would be met. The system plus its surroundings -- Michael Jordan, the ball and the arena -- would contain precisely the same amount of energy after the exercise as they had before.

If the energy produced in a system such as Michael Jordan were converted solely to heat, little would be accomplished. No soaring dunks, no rebounds, no buzzer-beating three-point shots. All the energy would pass from system to surroundings as heat. The warm system causes the gas molecules in the surroundings to increase speed. The impact of collisions of these molecules spreads the extra energy throughout the surroundings. This thermal motion determines the temperature.

But energy has a strange duality. It appears as heat but can also appear as work. Work is the energy used to move an object against an opposing force. Michael Jordan's jump shot is work -- he moves his body upward against the force of gravity. His dribble is work. He pushes the ball toward the floor faster than gravity would take it -- fast enough so the ball can compress the air trapped inside and rebound against gravity back to his hand. The compressing ball as it strikes the floor, the rebounding ball as it returns to Michael Jordan's hand, all of this is work. Energy used to overcome gravity, to deform the shape of the air-filled ball.

When we express energy as the sum of heat and work, we are making a very specific claim concerning these two properties. They are related. The relationship between heat and work is a close one, so close the amount of heat and the amount of work must be expressed with numerical values having the same units. Within limits, energy may be controlled to appear as heat (as we use electric power to dry clothes in a dryer) or work (the same electric power rotating the drum in the same dryer).

Common practice in General Chemistry uses heat as the marker for the energy evolved from or required to carry out chemical reactions. In our discussion of exercise, the ability and limitations of work are important. We will focus on heat and work as the quantitative products of chemical reactions.

Briefly, we define the amount of heat and/or work using two units, the Joule (J), and the calorie. The Joule and the calorie are related as follows:

Both units represent quite small increments of energy. We must add 1 calorie of heat to increase the temperature of 1g of water 1 degree Celsius. Our bodies expend about 1J of work with a single heartbeat. For convenience sake, both the Joule and calorie are often expressed in multiples of 1000. We speak of the kilojoule (kj):

Thus we must add 4.184 kJ of heat to raise the temperature of 100g of water 10 degrees Celsius.

A 59 kg woman playing basketball will work off approximately 8 kcal/minute (2000kJ/hr). In these two examples we quantify work and heat interchangeably. Heat and work can each be expressed in kJ or kcal.

Deep in the history of our sciences, nutritionists began assigning a caloric value to foods as energy sources. The familiar food calorie, the Cal, is equivalent to 1 kcal. When we present the startling claim that a 59kg basketball player is expending 8kcal/minute, we are really indicating 8Cal, 480Cal/hour of energy. To begin to relate energy expended as work in exercise to the chemistry of energy production, we must know how much energy is available from the various sources. Carbohydrates yield approximately 4kcal/g as they are used by the body for fuel. Our 59kg basketball player, using 480Cal (=480kcal) would consume the equivalent of :

Earlier we stated that exercise physiologists suggest replenishment of about 45g/hr of carbohydrate is the maximum amount that is useful for an immediate energy boost. At a rate of consumption of 120g/hr for this basketball player we learn that there is apparently no way to restore fully the energy drain during exercise.

How do we measure energy changes in a system? Two approaches seem possible. Suppose we could measure heat under conditions in which no work was performed on or by the system. Or suppose we could measure the work done by or on the system in the absence of any change in heat. Since energy released is the sum of heat evolved and work done, if we hold either one to a zero value, quantification of the other will lead to a value for energy.

Heat in the absence of work is the easier to measure. We determine the amount of heat flowing in or out of a system by measuring system temperature and multiplying that change in temperature by a property called the heat capacity of the system.

Heat Evolved(kJ) = Temperature Rise(degrees Celsius)xHeat Capacity(kJ/degrees Celsius)

Recall the United Center and Michael Jordan's brief exercise. Jordan as system, arena as surroundings, and the amount of energy released by Jordan was absorbed by the arena. Conservation of energy requires that energy moved from system will be absorbed by surroundings with no appearance or disappearance of energy. Practical determination of heat capacity cannot use Michael Jordan or the United Center. We use instead, a calorimeter -- an insulated vessel that contains system and surroundings.

An Early Laboratory Calorimeter
Photographed in the Tower at Georgia Tech, Atlanta, GA, 1904

We choose to determine energy in a chemical process by measuring heat in the absence of work. We measure the heat of the reaction, the heat evolved in the absence of work. Since work requires movement, action, action against a restraining force, we must find conditions in which there is no movement or action. We must establish a zero-work situation so that all energy will appear as heat. In a rigid, sealed container; any chemical process carried out within the vessel will do no work. The walls will not move. No movement. No work. The opposite proposition -- that we might measure energy by determining work in the absence of heat -- turns out to be impossible to carry out. Where there is work, there is always heat. The reasons are beyond the scope of our studies. But this fact has important implications.

It implies that machines, even machines like Michael Jordan, cannot produce work at 100% efficiency. Machines produce work and energy. The trick is to maximize the heat or the work, maximize the efficiency of conversion of fuels to the form of energy we wish to use. Maximizing heat is much easier to do. We have seen how we can achieve virtually 100% appearance of energy as heat.

We learned the foundation of thermochemistry rests on the ability to link the amount of heat released or required to the chemical equation for the specific chemical change. We call this heat, measured under zero-work conditions -- the change in enthalpy. As Robert Cade, Dana Shires and associates considered fuels such as carbohydrates in exercise, they used the change in enthalpy of the fuel-consuming chemical reactions to arrive at the energy that would be available. The amount of energy available from a given amount of fuel does not vary with how slow or fast we exercise. Enthalpy itself is a state property. All materials have enthalpy as an element of their nature. It is the change in this property through chemical processes which concerns us. The oxidation of 0.5g of glucose yields a certain amount of energy regardless of of how slow or fast the reaction takes place.

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Exercise in a Changing Environment

In the real world (during exercise, for example), energy, work and heat, are produced in complex, changing systems. Measurements will be imprecise unless we can establish a level playing field for energy measurement. We learned to call these predetermined conditions the standard state.

Accurate determinations of energy that can be compared from time to time and place to place depend on measurements recalculated to 25 degrees Celsius, with each of the components in its natural form at that temperature. To develop a standard for chemical comparison, we deal with standard enthalpy changes. For the exothermic reaction:

the standard enthalpy is -2.8MJ/mol. The molecular weight of glucose is 180, the oxidation 1g yields about 16kJ of energy, about 4kcal. Oxidation of proteins also yields about 4kcal/g.

Fats yield much higher energy per unit mass than carbohydrates. We understand this fact as a result of the structure of the fats in comparison to carbohydrates. If the process of oxidation takes a carbon and hydrogen-containing molecule to carbon dioxide and water while producing energy, is it not reasonable that we might get the greatest amount of energy from molecules containing little or no oxygen prior to oxidation? Wouldn't we guess that a molecule such as glucose, already containing 53% oxygen, would produce less energy than fats that might contain 11% oxygen?

Consider the oxidation of a beef fat component, tristearin, C57H110O6, molecular weight 890.

The standard enthalpy for tristearin oxidation is about -34MJ/mol, roughly 9kcal/g as compared to 4kcal/g for carbohydrates and proteins.