Enthalpy changes

ΔH2 The Reaction between CaO + HCL

J = m.c. ΔT is used to calculate the energy produced using   heat capacity of HCL, and 50ml of HCL with the temperature change in the reaction.

J = m.c. ΔT        50 x 4.2 x 10.5 = -2205 J

                        50 x 4.2 x 10    = -2100 J

                        50 x 4.2 x 9      = -1890 J

                        AVG                 = -2065 J

moles =    mass                        1.4    = 0.025

            molar mass                         56

ΔH2  = -82.6           units = kJ/mol

ΔH3

-18.22 kJ/mol - - 82.6 kJ/mol = 64.38 kJ/mol

Alternatively using Hess’s cycle

        

So, ΔH1 -18.22 kJ/mol - ΔH2 -82.6 kJ/mol = ΔH3 64.38 kJ/mol

Note- values are negative as they are an exothermic reaction.

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Evaluation 

In this experiment the goal was to determine the enthalpy changes of a reaction. Although I have calculated the change I believe that there errors made during the reaction which corrupted the final results. The enthalpy change was supposed to be around 178 kJ/mole; however, my enthalpy change was 64 kJ/mole. In this evaluation I will determine what errors caused my results to be so dissimilar.

Procedural errors

In the experiment there were two main procedural errors, the first one was that there was considerable heat loss while the experiment was being carried out. This can have a dramatic effect on the enthalpy changes of a reaction. If, say 1°C of heat loss was saved lost on each of the above reactions then the resultant enthalpy change would have been very different as we can see by doing the calculations below

ΔH1

J = m.c. ΔT        50 x 4.2 x 3.5 = -735 J

                        50 x 4.2 x 3    = -630 J  

                        50 x 4.2 x 3    = -630 J  

                        AVG               = -665 J

665  = -26.6 kJ/mole

0.025

If we do the same for the ΔH2 reaction then we get -91 kJ/mole.

If we then use Hess’s cycle to find ΔH3 then we get 117.6 kJ/mole.

This shows just how 1°C of heat loss can affect the results of the reaction.

To combat this next time I will conceal the experiment in a polystyrene cup. Polystyrene has excellent heat absorbing qualities, and would ensure that little heat is lost during the reaction. I could also put a lid on the beaker during the reaction.

Another error was that there was not enough volume of solution for the bulb of the thermometer to be completely submerged. This meant that the thermometer could not record the change in temperature to sufficient accuracy. As I have previously shown the temperature can greatly affect the results so this makes the experiment less reliable. To solve this next time I would simply use more volume of solution so that the bulb is totally covered. Alternatively I could have used a smaller beaker.

Out of the two experiments I got the most accurate was the CaO experiment. This is because of the difficulty of measuring the temperature change. In the CaCO3 reaction, the temp change was only 2.5°C at the most, while in the CaO experiment the greatest change was 10.5°C. This means that a lower resolution of the temperature would affect the CaO reaction less as the change would be of a lower percentage of the recorded change. As well as this we had sufficient weighing scales to measure the amounts of CaO and CaCO3 powder, so any changes would be negligible.

Errors in measurement

An important reason why the results were so different to the prediction was because of the errors in measurement and the fact that some of the measurements made were not very accurate.

The first error would be that the measuring cylinder used to measure the acid was too big. We used a 250 ml measuring cylinder when all we needed to use was 50 ml of HCl. This means that we would have an error of up to 3ml while measuring the HCl. If we take this into account in our calculations then the results become very different.

ΔH1

J = m.c. ΔT        53 x 4.2 x 2.5 = -556.5 J

                        53 x 4.2 x 2    = -445.2 J

                        53 x 4.2 x 2    = -445.2 J

                        AVG               = -482.3 J

482  = -19.28kJ/mole

0.025

If we do the same for the ΔH2 reaction then we get -87.5 kJ/mole.

If we then use Hess’s cycle to find ΔH3 then we get 68.22kJ/mole.

This shows just how 3ml more of HCL could affect the reaction. It could affect it even more as well. If there was a =3ml error on the first reaction and a -3ml on the second then we get 77.6kJ/mole. To stop this from happening again I could use a 50ml pipette to measure the solution to a much smaller resolution

Another error in measurements was that the thermometer’s resolution wasn’t small enough. We could only detect a 0.5°C change in temperature. As we have seen before a temperature error can drastically change the results. To avert this problem next time, I cold use a loggit temperature probe. This is an electronic device that measure temperature to 2 d.p. This would make my results much more reliable.