53902.8/1000=53.9028
ENTHALPY CHANGE = -53.90 kJ mol¯¹
Enthalpy change of HNO3:
Q = M x C x change in T
= (25 + 21) x 4.2 x 7.2
= 46 x 4.2 x 7.2
= 1391.04 J
No of moles = C x V
= 1 x 25/1000
= 0.025
1/0.025 = 40
40 x 1391.04 = 55641.6
55641.6/1000=55.64kJ
ENTHALPY CHANGE = -55.64 kJ mol¯¹
Enthalpy change of H2SO4:
Q = M x C x change in T
= (25 + 24) x 4.2 x 8.1
= 49 x 4.2 x 8.1
= 1666.98 J
No of moles = C x V
= 1 x 25/1000
= 0.025
1/0.025 = 40
40 x 1666.98 = 133358.4
133358.4/1000=133.36kJ
ENTHALPY CHANGE of 2 moles = 133.36 kJ
ENTHALPY CHANGE = 133.36/2 = -66.68 kJ mol¯¹
(e)
Comparison
From my results I can see that the enthalpy changes for neutralisation for each acid have fairly similar values. The difference in values can be explained by sources of error. So therefore, I can conclude that the enthalpy of neutralisation of an acid with an alkali does not depend on the particular acid or alkalis providing both are strong. My results clearly show this as all acids and alkalis were strong and gave approximately similar values. However, you have to use more moles of weak acid to achieve the same extent of neutralisation compared to using a strong acid, assuming both acids are of the same basicity. Hence, for a weak acid, the enthalpy change per mole of weak acid is less because you used more weak acid than strong acid to achieve the same extent of neutralisation.
(f)
Evaluation
The values for Hydrochloric acid and Nitric acid were -53.90 kJmol¯¹ and -55.64 kJmol¯¹ respectively. These were relatively close to the expected value of -57 kJmol¯¹ however; the value for Sulphuric acid was not. The enthalpy change of sulphuric acid was of -66.68 kJmol¯¹ and was far from the value I was expecting.
This may have been due to the concentration of sulphuric acid not being made accurately enough or the apparatus being contaminated. The 2 things that were likely to be contaminated was firstly the pipette with which I measured the 25cm³ volume of acid, or the polythene cup in which all 3 reactions had taken place. The cause of this contamination might have been due to the cup not being rinsed properly after a titration, or the acid inside pipette not rinsing the pipette through accurately enough. The fact that the same cup was used for all three titrations might have been a cause of contamination, which resulted in the attainment of the results shown above. The hydrochloric, nitric or sulphuric acid could easily have remained when I was carrying out the titration for the other one. This would have allowed the temperature to have changed differently affecting the final enthalpy change value I obtained. Also, there was no clear indication on the method given, for me to rinse out the polythene cup with the next acid or to wash it out with tap or distilled water.
Even though I gained good results, there were also many sources of errors. One source of error could have been that while reading the temperature off of the thermometer, I could have read it wrong or the mercury in the thermometer was broken up. I also may have held the pipette from the bulb which would have affected the volume that it holds and therefore alter my results. Another error is that there may have been many inaccurate burette readings. This would clearly affect the temperature change of the reaction as it would mean that more of the reaction has taken place at the amount of liquid stated, and so the rate would be faster than it should be at that point. As a result, the values for the enthalpy change for one mole of neutralisation would be different, making my results inaccurate.
To further improve my experiment I could read the thermometer and burette more accurately, allowing a much smaller margin of error. Also I would deter from holding the pipette from the bulb, as this would seriously affect the volume that it holds, and in turn make my experiment void and of little use, as it wouldn’t be entirely accurate.
