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Chemistry - Energy in Chemical Reactions

A chemical reaction resulting in a net release of energy is called exothermic. Conversely, a chemical reaction requiring a net input of energy to occur is called endothermic. The relationship between chemical reactions and energy exchange is correlated with the breaking or making of chemical bonds. Atoms bonded together represent a lower state of total energy than those same atoms existing separately, all other factors being equal. Thus, when separate atoms join together to form a molecule, they go from a high state of energy to a low state of energy, releasing the difference in energy in some form (heat, light, etc.). Conversely, an input of energy is required to break that chemical bond and force the atoms to separate.

An example of this is the strong bond between two atoms of hydrogen (H) and one atom of oxygen (O), to form water (H2O). When hydrogen and oxygen atoms bond together to form water, they release energy. This, by definition, is an exothermic reaction, but we know it better as combustion: hydrogen is flammable in the presence of oxygen.

A reversal of this reaction occurs when water is subjected to an electrical current, breaking water molecules up into hydrogen and oxygen gas molecules. This process of forced separation requires a substantial input of energy to accomplish, which by definition makes it an endothermic reaction. Specifically, the use of electricity to cause a chemical reaction is called electrolysis.

Energy storage and release is the purpose of the so-called “hydrogen economy” where hydrogen is a medium of energy distribution. The reasoning behind a hydrogen economy is that different sources of energy will be used to separate hydrogen from oxygen in water, then that hydrogen will be transported to points of use and consumed as a fuel, releasing energy. All the energy released by the hydrogen at the point of use comes from the energy sources tapped to separate the hydrogen from oxygen in water. Thus, the purpose of hydrogen in a hydrogen economy is to function as an energy storage and transport medium. The fundamental principle at work here is the energy stored in chemical bonds: invested in the separation of hydrogen from oxygen, and later returned in the re-combination of hydrogen and oxygen back into water.

The fact that hydrogen and oxygen as separate gases possess potential energy does not mean they are guaranteed to spontaneously combust when brought together. By analogy, just because rocks sitting on a hillside possess potential energy (by virtue of being elevated above the hill’s base) does not means all rocks in the world spontaneously roll downhill. Some rocks need a push to get started because they are caught on a ledge or resting in a hole. Likewise, many exothermic reactions require an initial investment of energy before they can proceed. In the case of hydrogen and oxygen, what is generally needed is a spark to initiate the reaction. This initial requirement of input energy is called the activation energy of the reaction.

Activation energy may be shown in graphical form. For an exothermic reaction, it appears as a “hill” that must be climbed before the total energy can fall to a lower (than original) level:

 

For an endothermic reaction, activation energy is much greater, a part of which never returns but is stored in the reaction products as potential energy:

 

A catalyst is a substance that works to minimize activation energy in a chemical reaction without being altered by the reaction itself. Catalysts are popularly used in industry to accelerate both exothermic and endothermic reactions, reducing the gross amount of energy that must be initially input to a process to make a reaction occur. A common example of a catalyst is the catalytic converter installed in the exhaust pipe of an automobile engine, helping to reduce oxidize unburnt fuel molecules and certain combustion products such as carbon monoxide (CO) to compounds which are not as polluting. Without a catalytic converter, the exhaust gas temperature is not hot enough to overcome the activation energy of these reactions, and so they will not occur (at least not at the rate necessary to make a significant difference). The presence of the catalyst allows the reactions to take place at standard exhaust temperatures.

The effect of a catalyst on activation energy may be shown by the following graphs, the dashedline curve showing the energy progression with a catalyst and the solid-line curve showing the reaction progressing without the benefit of a catalyst:

 

It should be noted that the presence of a catalyst has absolutely no effect on the net energy loss or gain resulting from a chemical reaction. With or without a catalyst, the difference in potential energy before and after a reaction will be the same. The only difference a catalyst makes to a chemical reaction is how much energy must be initially invested to spark the reaction. To use the example of hydrogen and oxygen gas once again, the presence of a catalyst does not cause the combustion of hydrogen and oxygen to release more energy. All the catalyst does is make it easier for the combustion to begin.

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written by isabelle, March 12, 2013
really helpful and interesting/fascinating information! thank you!

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