General:

    In general a chemical reaction involves bond breakage and bond formation.  If we assume that all the energy associated with this process is in the form of heat then, a good approximation of the energy released or needed can be obtained by:

This works well when all the reactants and products are in the gaseous phase.  However even, here it is only an approximation since, the bond energies used to determine the heat are but average bond energy values. 

    The majority of the reactions that are done in the lab are of the open beaker variety, or in thermodynamic terms, reactions that are carried out at constant pressure.   The heat transferred under these conditions is referred to as Enthalpy, DH. If we carry out the at standard temperature and pressure then it is referred to as DH^o.  One of the nice things about DH^o for a reaction is  that we often do not have to experimentally measure it, but can determine it from known heats of formation (DH^o_f ) of the reactants and products in the reaction.

But where do these values come from?

Calorimetry:

    The experimental measurement of the heat evolved in a chemical reaction is referred to as calorimetry.  The equipment used is often very sophisticated depending on the degree of accuracy required of the value. It can also be relatively simple if all one wished to do is to get a good estimate of the value.  No matter the devise, a common theme in most calorimetry experiments is to have the components react and,  allow the heat evolved to be transferred to or from another medium, in many cases water.  What you measure is the temperature change in the medium, i.e. water and so long as the temperature change does not involve a phase change in the water then the heat transferred to the water can be determined by

    Knowing the quantity of material used in the reaction one can then determine the heated evolved in whatever units one requires. 

    Setting the q_reaction equal to the q_water is an over simplification, in that the heat evolved from the reaction is not only transferred to the water, but to the calorimeter and all the other objects that come in close proximity.  We can at least take our determination one step further in that, we can take into a account the heated transfer to or from the calorimeter.

The calorimeter constant (associated with any calorimeter) can be viewed as the mass of the calorimeter x the heat capacity of the calorimeter.  Thus we can obtain a better value with

    Now if we really wanted a better value then we need to take into account the heat that was transferred to the surrounding air and any other objects in proximity to the reaction.  Rather than do this we often design our calorimeters such that this is minimized.

    Two very simple calorimeters are depicted .  The first is a simple aluminum can calorimeter.  Aluminum is very light and its specific heat capacity (0.902 J/g.K) is very small. The amount of heat used to warm up the calorimeter to the temperature of the water inside the aluminum can is very small and can to a first approximation be ignored.

    Those four things!, protruding from the can are in fact paper clips.  These are used to simply balance the aluminum can calorimeter on a ring stand.

    The second calorimeter depicted is often referred to as a coffee-cup calorimeter.  It is just two Styrofoam cups placed inside one another.  This type of calorimeter has the advantage of minimizing the amount of heat that is lost or gained by the water to the surrounding environment.

The Thermodynamic Teaser:

    This you will be given when you come to lab!