5.4: Procedure

Construct a simple calorimeter by nesting two styrofoam cups together. Make a hole in the cardboard lid just large enough to insert the thermometer. Thoroughly rinse and dry the calorimeter. Place the calorimeter on a piece of cardboard to insulate it from the laboratory bench.

Obtain approximately 110 mL of 1.00 M hydrochloric acid in a 250 mL beaker. Carefully measure 50.0 mL of the 1.00 M HCl solution and add it to the calorimeter. Put the lid on the calorimeter, but do not insert the thermometer. Obtain approximately 110 mL of 1.00 M sodium hydroxide in a 250 mL beaker. Carefully measure 50.0 mL of the 1.00 M base solution in a clean, dry graduated cylinder. Measure the temperature of the base solution. Rinse and dry the thermometer and then measure the temperature of the acid solution. Accurate temperature measurements are crucial for good results in this experiment.

Lift the calorimeter lid and rapidly add all the base solution. Avoid splattering the solution. Insert the thermometer into the lid and gently stir the solution. Be sure the bulb of the thermometer is completely immersed in the solution. Watch the temperature rise on the thermometer and record the maximum temperature. Do not read the thermometer before the temperature stops rising. As the temperature of the mixture rises, heat from the reaction is lost to the calorimeter cup and its surroundings; therefore, the maximum observed temperature change is somewhat less than that actually attained by the reaction mixture. You will not need to correct for such heat losses in this experiment. Empty, rinse, and dry the calorimeter.

Repeat the above procedure for a second trial and average the results of the two trials. For computational purposes you may assume the specific heat capacity and density of your solution are approximately that of water, or 4.184 J/g ^oC and 1.00 g/mL, respectively.

DISPOSAL and CLEANUP: All solutions can be washed down the sink drain with plenty of water. Thoroughly rinse the calorimeter and dry with a paper towel.

DATA

Write the complete, balanced chemical equation for the reaction that occurs.

Write the net ionic equation for this reaction.

Enthalpy of Neutralization Trial 1 Trial 2

1. Molarity of HCl __________ __________ M

2. Molarity of NaOH __________ __________ M

3. Volume of acid solution __________ __________ mL

4. Volume of base solution __________ __________ mL

5. Total volume after mixing __________ __________ mL

6. Density of solution (see previous page) __________ __________ g/mL

7. Total mass of solution __________ __________ g

8. Temperature of acid solution (T_A) __________ __________ ^oC

9. Temperature of base solution (T_B) __________ __________ ^oC

10. Average temperature before mixing __________ __________ ^oC

11. Maximum temperature after mixing __________ __________ ^oC

12. ΔT = T_{after mixing} – T_{before mixing} __________ __________ ^oC

13. Heat gained by calorimeter cup = C_{cal cup} x ΔT __________ __________ J

(C_{cal cup} = 20. J/^oC)

14. Heat gained by solution = m_soln x c_soln x ΔT __________ __________ J

15. Total heat of the surroundings = Q_{cal cup} + Q_soln __________ __________ J

16. Heat of the reaction __________ __________ J

17. Moles of acid used __________ __________ mol HCl

18. Moles of base used __________ __________ mol NaOH

19. Moles of H₂O produced ___________ __________ mol H₂O

20. Molar enthalpy of reaction = ΔH = total heat/mol H₂O ___________ __________ kJ/mole^*

21. ΔH average __________ kJ/mole (include sign)

DISCUSSION

Is this reaction endothermic or exothermic? Draw an energy diagram for the reaction.

What measured (not calculated or given) quantity limits the number of significant figures in the molar enthalpy?

Given the accepted value of the molar enthalpy for this reaction is -55.84 kJ/mol, what is the percent error for your experiment?