Enthalpy Change for the Decomposition of Ammonium Chloride PDF

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Summary

Lab reported earned an A. ...

Description

Enthalpy Change for the Decomposition of Ammonium Chloride Purpose To determine the calorimeter constant (J/°C). To determine the enthalpy of neutralization and dissolution (J/mol) for reactions.

Results Experiment 1 data 45.00 ml

Volume of cold water in the calorimeter Initial temperature of the cold water (°C)

21.50 °C Volume of hot water added to the calorimeter (ml) Initial temperature of the hot water (°C)

45.00 ml 60.00 °C

Final temperature in the calorimeter (°C) 34.90 °C Experiment 1 data analysis Mass of water calculation = density * volume: 45.00 ml * 1 g/ml = 45.00 g of water Q = m C*ΔT Q hot = 45.00 g * 4.184 J/g*°C * (34.90 °C- 60.00 °C) = - 4725 J Q cold = 45.00 g * 4.184 J/g*°C * (34.90 °C- 21.50 °C) = 2522 J Q cal = -(Q cold + Q hot) Q cal = -( 2522 J-4725 J) = 2203 J ΔT cal = 34.9 °C – 21.50 °C =13.4 °C C cal = Q cal /ΔT; 2203 J / 13.4 C = 164.4 J/ °C

Experiment 2 data Volume of HCl solution in the calorimeter (ml) Volume of NH3 solution added (ml) Final temperature in the calorimeter (°C) HCl concentration NH3 concentration Experiment 2 data analysis HCl + NH3 --> NH4Cl 2 M * 0.025 L of NH3 = 0.050 moles of NH3

25.00 ml 25.00 ml 29.00 °C 2.00 M 2.00 M

2 M * 0.025 L of HCl = 0.050 moles of HCl Q = m C*ΔT -50.00 g *4.184 J/g°C * (29.00 °C - 21.50 °C) = -1569 J = Qsol Heat change of calorimeter = 164.4 J/C * 7.500 °C =1233 J Q rxn= -(-1569 J +1233 J) -1233 J +1569 J =336.0 J / 0.0500 moles = - 6720 J/mol of HCl and NH3

Experiment 3 data Volume of cold water added in the calorimeter (ml)

25.00 ml

Initial temperature of the hot water (°C) 21.50 °C Mass of NH4Cl added to the calorimeter (g) Final temperature in the calorimeter (°C)

5.000 g 11.00 °C

Moles of NH4Cl= 5.000 g/ 53.49 g/mol = 0.09348 moles Q=mC*ΔT ΔH dissolution = -25.00 g * (11.00 °C-21.50 °C) * 4.184 J/g*°C =1098 J; 1098 J/ 0.09348 moles = 11750 J/mol

Discussion The enthalpy change for NH4Cl was determined by a series of three experiments. The first experiment was required to find the calorimeter constant which was 164.4 J/°C. The calorimeter constant reflected the calorimeter’s insulative abilities. The second experiment determined the enthalpy of formation of NH4CL, which demonstrated that it was an exothermic reaction. The final experiment explored the solubility of NH4Cl, and its enthalpy of dissolution was found to endothermic which was apparent by the negative ΔT value. The official ΔH dissolution for NH4CL was 14780 J/mol at rtp, the experimental value was 11750 J/mol.

Conclusion The experiment was successful as the calorimeter constant (J/°C), enthalpy of neutralization (J/mol) and enthalpy of dissolution (J/mol) were determined....