# Calorimetry

Pages: 5 (975 words) Published: January 21, 2013
Experiment 1: Calorimetry

INTRODUCTION

In the calibration of the calorimeter, the net ionic equation used is H+(aq) + OH-(aq) H2O(l). The reaction released heat and is said to be exothermic. HCl is the limiting reactant of the reaction and o.oo5 moles of it was used. The heat generated by the reaction is 55.8 kJ. The sign of T of the reaction used for calibration is opposite to that of H.

In the determination of heats of reaction, the reaction of 15 mL 1 M CuSO4 + 0.05 g Zn produced a net ionic equation of CuSO4 + Zn ZnSO4 + Cu. The limiting reactant of the reaction is Zn and 0.05 moles of it was used. The sign of T of the reaction used for calibration is opposite to that of H.

WORKING EQUATIONS:
H+(aq) + OH-(aq) H2O(l) Hrxn= -55.8 KJ/mol Hrxn=qrxn / nLR
qcal=CcalT qcal=-qrxn
Ccal=mass x specific heat
Hrxn=npHop - nr Hor
qcal= [Ccal+ mass solid (ppt.) x specific heat of ppt.] x T

RESULTS AND DISCUSSION

The reaction of HCl and NaOH for calorimeter calibration brought about a 1°C change in temperature (from 29.5 °C to 30.5 °C). The results of the other groups’ calibration are found in TABLE A of the data sheet.

The displacement of one metal by another (1M CuSO4 and 0.5g Zn) resulted to an increase in the temperature of the system. The 30 °C temperature measured from 15 mL 1 M CuSO4 went up to 31.5 °C upon addition of a 0.05-gram Zinc dust. Therefore, it is an exothermic process. All other neutralizations and reactions in the experiment have negative enthalpy values, so they are also exothermic.

It is therefore concluded that the factors that determine the enthalpy values of reactions are the kinds of reaction that have occurred and the nature of reactants involved.

REFERENCES

 Silberberg, M.S. Principles of General Chemistry 2nd Ed. McGraw-Hill Companies Inc., New York. 2010

Petrucci, R.H., Herring, F.G., Madura, J.D., Bissonnette, C. General Chemistry: Principles and Modern Applications.Tenth Edition. Pearson Canada Inc., Toronto, Ontario, 2011.

General Chemistry II Laboratory Manual. Institute of Chemistry, University of the Philippines Diliman. June 2011

1. After obtaining the theoretical values of ΔHrxn, explain any discrepancy of the values to the experimental. Give some possible sources of errors.

Some calculated actual enthalpy value is of higher than the theoretical values. Also, some calculated actual enthalpy value is less than the theoretical values. This might be caused by inaccuracies in the concentrations of prepared solutions. If the concentration of the solution used is greater than the theoretical, then there would be a greater number of moles present in the solution. This would reflect to a greater enthalpy value.

2. In the procedure for the determination of ΔH, explain why it is important:

a. that the total volume of the resulting solution be 15 mL?

To uphold the Law of Mass Conservation in an isolated system where no matter (and energy as well) shall be exchanged between the system and its surroundings. Moreover, it shall be used in determining the Ccal, the calibration constant, which shall become the basis of calculations all throughout the experiment.

b. to know the exact concentrations of the reactants?

This shall be used to determine the number of moles of the limiting reactants (through stoichiometry). Since ΔHrxn = _qrxn_
nLR

c. to know the exact weight of the metal solids used?

It is necessary for the equation qcal = mcΔT, where m is the sum of the weights of the metal solid and the solution (from given density).

3. The neutralization of 200 mL of 0.5 M HA by sufficient NaOH evolves 6.0 kJ of heat.

a. Calculate the enthalpy change for the neutralization of 1 mole HA.

ΔHrxn= _qrxn_ =__- 6.0 kJ__ = - 60 kJ/mol
nLR0.100 moles

b. Is HA a weak acid or strong acid? Justify your answer using thermochemical...

References:  Silberberg, M.S. Principles of General Chemistry 2nd Ed. McGraw-Hill Companies Inc., New York.
2010