Balancing Chemical Energy

Chemical Energy obeys the same kind of relationships that we expect for all forms of energy. Controlling the production and use of Chemical Energy is something that all life on Earth must master. Different forms of energy may be interconverted, as in this example of the conversion of potential energy (mgh) into kinetic energy (mv²/2, panel a), and finally into heat (Q=mCDT, panel b):

Chemical energy is the energy stored in molecules by virtue of the arrangement of atoms with the molecule. When atoms are rearranged to create new molecules (i.e. a chemical reaction), energy is either released or consumed during the process. The amount of energy released or consumed depends on the relative stability of the reactants and products.

When Hydrogen and Oxygen gases are reacted, chemical energy stored internally in the reactants is released into the surroundings, usually in the form of heat. The internal energy of the system has decreased (downhill in an Energy Plot) so this is an Exoergic Reaction. Synthesis of Hydrogen and Oxygen gas by decomposition of water requires the imput of energy, is uphill on the graph, and is called an Endoergic Reaction.

Energy, in all its forms, must be conserved in a chemical transformation. This simple statement is sometimes referred to as the First Law of Thermodynamics and can be written:

DE = q + w

Since the system can exhange either heat or work to the surroundings, the internal energy change of the system must be the combination of the heat and work during the process, and thus the above equation. This equation implicitly defines the 'signs' of the values of heat, q, and work, w, in relation to the internal energy of the system and the surroundings. When the system is heated with a bunsen burner, for instance, the q for the process is positive. If the system expands and 'pushes' against the surroundings, like the piston in your auto engine, the w for the process is negative.

In general, when we perform a chemical reaction in the laboratory, the heat and work flow during the course of the process. Therefore, for processes that take place at contant external pressure (like things that happen in the room we are now in) it is sometimes better to think of a measure of the internal energy very closely related to E, called H the Enthalpy of the system.

The Heat flow during a process that is carried out a constant external pressure is exactly equal to the change in Enthalpy of the system.

Since the Energy and Enthalpy are very closely related, and the Enthalpy changes during a typical chemical reaction are much easier to measure, Enthalpy is a widely used quantity to gauge the 'internal chemical energy content' or 'ability to provide heat' for chemical substances. Processes that change the Enthalpy of the system are catogorized into those that release heat, Exothermic, and thoses that consume heat, Endothermic.

Another way of looking at exothermicity is that the change in state is downhill on an enthalpy diagram.

We have defined a bunch of terms that have common definitions but are actually strictly defined in thermodynamics. To help get things straight, the following Glossary of Thermodynamic terms may be useful.

The Universe is divided into...
The System: The part of the universe in which you are interested. No Thermodynamic measurements actually take place within the system, yet this is the part of the Universe that we learn about.
The Surroundings: Everything outside the system. This is where the manifestation of processes taking place in the system are felt.
The Boundary: The separartor between the system and the surroundings. This boundary contains no matter or volume, but Heat and Work or even Mass may flow across it during a Process.

The System is defined by its'
State Function: Any property of the system that does not depend on its History. Temperature, Pressure, color, Enthalpy are State Functions. Work and Heat are not.
Extensive Property: Any property that scales with the size of the system is extensive. Volume, mass, Enthalpy are extensive. You must report any value of an extensive property on a per mole basis.
Intensive Property: Any property that does not change when you cut the system in half is intensive. Temperature, color, and pressure are all intensive.

The System can undergo a...
Change of State: A transformation that is defined by just the initial State and Final State of the system. Work and heat flow during a Change in State.
The Path: a definition of all the intermediate states between the initial state and the final state in a change of state. Every change of state has an infinite number of paths. (A trip from Gainesville to Denver can be driven in many ways, or even flown, but the change in position (state) is the same).
A Process: Any external conditions that are maintained during a change of state may define the path and thus be called a process. Typical process occur at constant pressure (isobaric) constant temperature (isothermal) and constant volume (isochoric).

Work: Work is a form of Energy that flows across the Boundary during a Change in State that can be entirely converted into lifting a weight in a gravitational field in the surroundings. Work is directed energy, and is the type of energy that we wish to generate when we drive a car or take an elevator. The amount of work that we get from a given change in state depends on the path or process. Work is NOT a state function; it only occurs during a Change in State. Work can be entirely converted into Heat without loss, and thus both Heat and Work have units of Energy. The metric unit of work is the same as all other energy units, the Joule (twice the kinetic energy of a 1.00 kg object travelling at 1.00 m/s). We commonly see the related unit, the Watt which is the energy release of 1J per second, on light bulbs and stereo equipment.

Heat: Heat is a form of Energy that flows across the Boundary during a Change in State that can be entirely converted into changing the temperature of a mass of water in the surroundings. Heat is random (thermal) energy and is the type of energy that we would like to use when we make work. Heat is not a state function; it only occurs during a Change in State. Heat can only be partially converted to Work: There is always loss or 'waste' in the conversion of random energy to directed energy. The 'natural' metric unit of heat is the calorie, which is the amount of heat required to raise 1.00 cm³ = 1.00 g water by 1.00 Kelvin. The Calorie (actually the kilocalorie) is commonly used in quoting the 'fuel value' of foods.

Since Work can be completely converted into Heat, We can easily determine the 'mechanical equivalent of heat', i.e. the number of Joules in the calorie. This number is EXACT in the metric system.

4.184 Joules = 1 calorie

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PJ Brucat // University of Florida