Calorimetry Lab

Determination of the Heat Exchanged in Chemical Reactions

Introduction:

Calorimetry is the science of determining heat and energy exchange in various situations and reactions. It is used for many things everyday including solid and liquid fuel testing, waste disposal, and explosive testing. In our lab, we will are applying calorimetry to determine the change in enthalpy of a weak acid-strong base reaction. My beginning question was: How can we apply Hess’ law and calorimetry to chemical equations to determine the heat exchanged in a reaction? We started by putting 100 mL of water into a coffee cup calorimeter (a polystyrene cup inside another polystyrene cup as an insulator), a magnetic stir bar was added and using the program LabProTM the

2H3C6H5O7(s) + 3Mg(OH)2 (s) → Mg3(C6H5O7)2 (aq) + 6H2O (l) In order to get to this end point, we used Hess’ Law which says that the enthalpy of a net chemical reaction is the sum of the enthalpy changes of each individual step. There were four steps, or reactions, where we measured the change in enthalpy.

Step 1: Mg(OH)2(s) Mg(OH)2(slurry)

Step 2:2H3C6H5O7(s) +Mg(OH)2(slurry)Mg3(C6H5O7)2 (aq) + 6H2O (l)

Step 3:H3C6H5O7(s)H3C6H5O7(aq)

Step 4:2H3C6H5O7(aq) + 3Mg(OH)2 (slurry) → Mg3(C6H5O7)2 (aq) + 6H2O (l) By determining the change in enthalpies for each step, we were able to add all of the individual enthalpies together to get the change in enthalpy for our goal reaction. The results can be seen in Figure 1. We were able to calculate qsol, qrxn, and ΔH using: qsol=mcT qsol+qrxn=0 Hrxn=qrxnmol

We know that energy is constant therefore any heat lost by our reaction is transferred to the surroundings. I was not able to locate any literature values for the change in enthalpy, despite looking very extensively.

Error Analysis: For this experiment I assumed that the specific heat capacity for all solutions was 3.853 J/g°C. Since the actual specific heat capacities could have been more or less than this, there is no specific direction or magnitude that this would skew the results. The second is that we assumed the density of our citric acid monohydrate solution was 1.06 g/mL. Again, depending on what the actual density of our solution was, it could skew the result. My two values for the change in enthalpy, -118.21 and -155.31 kJ, don’t necessarily agree with each other, but do not completely disprove one another. One possible experimental error is that the magnetic stir bar that was placed into the solutions may not have always been spinning. In the beginning trials, it was very inconsistent. This would cause the temperature change to be less extreme because not all of the solution would be