Gatorade
Chemical Reactions - Our Source of Energy
Energy in Systems and Surroundings - Law of Conservation of Energy
Dr. Matt Hermes

Sports Science

Rehydration

Energy  Replacement

Quiz

Problems

Micro/Macro
and Symbolic
Representation

Chemistry uses macroscopic, large scale observations to help describe and understand matter at the unseeable, molecular level.

And then we represent both the microscopic and macroscopic with often complex symbolic representation.

Here we encounter an equation for the oxidation of glucose. The macroscopic effect is the energy and heat we generate. We cannot see the microscopic transition of each glucose molecule into energy and reaction products. The equation:

C6H12O6(s) + 6O2(g) --> 6CO2(g) + 6H2O(g)

is our symbolic representation of BOTH the macroscopic energy effects we experience AND the microscopic , molecular phenomenon.

Energy is a concept.  Most definitions of the word energy fail to provide its exact meaning when applied to scientific matters.  In science the word energy is a concept that expresses two measurable properties, heat and work.  Here is the relationship of energy, heat and work: Energy Released=Work Done + Heat Released The Law of Conservation 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. Remember that energy, in an open system, can do work on the surroundings or supply heat to the surroundings.

Chemical Concepts
Perhaps we can find no more important chemical concepts than those that relate energy, work and heat:

If the energy produced in a system were converted solely to heat as it is in the gas flames heating our homes, no work is accomplished. No mountains climbed in the Rockies, no soaring three-point shots in crowded basketball arenas. All the energy passes 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.

Energy's Strange Duality - Heat and Work

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. Basketball's three-point jump shot is work. Remember how Michael Jordan moved his body upward against the force of gravity. His dribble was work. He pushed the ball toward the floor faster than gravity would take it -- fast enough so the ball could 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 is used to overcome gravity, to deform the shape of the air-filled ball.

Quantifying Heat and Work

1 cal = 4.184J

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):

1 kJ = 1000J

and the kilocalorie (kcal).

1 kcal = 1000 cal

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

Chemical Reactions and the Production of Energy Heat and Work

We learned the foundation of thermochemistry rests on the ability to link the amount of energy released or required to the chemical equation for the specific chemical change. We often experiment under conditions where no work is done on or by the system.  The heat, evolved or required is the change in enthalpy. We use 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 burn the fuel such as in 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.  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:

C6H12O6(s) + 6O2(g) --> 6CO2(g) + 6H2O(g)

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:.

2C57H110O6 + 163O2 --> 114CO2(g) + 110H2O(g)

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