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Enthalpy of Hydration Lab

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Introduction

Marina Horta Enthalpy of Hydration Lab October 29/09 Abstract; The purpose of this experiment was to calculate the enthalpy of hydration for Magnesium Sulfate. Using the enthalpy of dissolution and Hess's Law the results of this lab were calculated to be -17197.18 kj mol-1. Using the graphs as a reference the final calculation can be determined to be on average correct. Introduction; The formation of a solution involves the interaction of solute with solvent molecules. When a solute is dissolved there are three energy process involved; The first is the separation of the solvent molecules which requires energy and so is an endothermic reaction, the second is also and endothermic reaction, the separation of the solute ions or molecules, and last is the attractive force between the solute molecules and the solvent, releasing energy ,exothermic. ...read more.

Middle

4 H2 (g) + 3C (s) +103.8 kJ 4 H2 (g) + 2 O2 (g) 4 H2O (g) -968 kJ 3 C (s) + 3 O2 (g) 3 CO2 (g) -1180.5 kJ which says that the enthalpy change is independent from the path taken and only the final and initial values matter. The values for each set will either result in a negative indicating an exothermic reaction or positive indicating an endothermic. The enthalpy of hydration can then be determined using the following equation; Materials; Styrofoam cups with lid, 50 ml volumetric pipette, stir plate and stir bar, ring stand and ring clamp, temperature probe, Xplorer GLX Data logger Distilled water, Anhydrous Magnesium Sulfate, Magnesium Sulfate hepathydrate Safety; MgSo4 is a hygroscopic and irritant MgSo4 * 7H2O is an irritant Calculations; Anhydrous trial 1 (3.49 g) ...read more.

Conclusion

The second two experimental graphs are negative for the anhydrous and show an exothermic reaction where the breaking of bonds between the solute and solvent releases energy and the graphs shows an increase in temperature with the addition of the solute. In comparison with other students however the results of this lab were too low, any error that can be attributed is most likely due to errors in calculations because the graphs used seem to be somewhat correct. Any experimental error that occurred would have been that not all of the solute dissolved, not allowing for the graph to be accurate. Leaving the anhydrous uncovered for any amount of time could also have added to error in calculations if any moisture was absorbed ( at this time humidity levels were higher) causing added weight. ...read more.

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