= -16.4769
≈ -16 kJ mol-1
For CaO:
Temperature change = 9ºC
We have to follow exactly the same process and calculations as with calcium carbonate, only we used a different amount of CaO, and the temperature change was different.
So, using the formula
E=ΔT x mass surroundings x specific heat capacity of surroundings
We get:
E= (-9) x 51 x 4.2
= -1927.8J
Then, finding amount of CaO we used, we use the following formula:
No moles = Mass ÷ Mr
Getting:
No moles = 1.39 ÷56.08
= 0.025 mol.
Then, having converted the energy intake into kJ (-1.9278 kJ), we put the data into this formula:
ΔH= (E ÷ 1000) ÷ No moles
Getting:
ΔH = (-1927.8 ÷ 1000) ÷0.025
= -1.9278 ÷ 0.025
= -77.112
≈ -77 kJ mol-1
Now we have found both measured enthalpy changes, we can work out the enthalpy change for the thermal decomposition of calcium carbonate. This is done by following Hess’s Law, which states that “The enthalpy change in a reaction is the same, independent of the number of steps taken, provided the reactants and products end up in the same states that they would ordinarily be”. It is this principle that is behind this experiment – since we are dissolving both chemicals in HCl, and starting with solid calcium carbonate, calcium oxide and carbon dioxide, we can work out the enthalpy change using this law. This is best done by constructing a triangle, as shown below:
CaCO3 ΔH3 CaO + CO2
ΔH1 ΔH2
CaCl2
We draw out the triangle, and can see that, in order to get from calcium carbonate to calcium oxide, we must go in the same direction as the ΔH1 arrow, and against the direction of the ΔH2 arrow. This means that ΔH3 = ΔH1- ΔH2
So, substituting in our values for those of ΔH1 and ΔH2, we get:
ΔH3 = ΔH1- ΔH2
ΔH3 = -16 + 77
ΔH3 = +61 kJ mol-1
According to the experimentally determined values, ΔH for the thermal decomposition of Calcium carbonate is +61 kJ mol-1.
Evaluation:
There were a large number of problems with this experiment, the most significant source of error being the heat loss to the apparatus. The reactions took place in a glass beaker. Glass has a relatively high heat capacity, and so a large amount of heat which is either taken in or given out is lost to the apparatus. This is obviously a major problem, especially with the reaction of calcium carbonate with HCl, as this does not appear to have a very large change in temperature, so every slight bit of absorption by the apparatus will have a large effect on the temperature change, and consequently the enthalpy change. Th other piece of apparatus that will have absorbed a lot of heat is the thermometer itself, as it is also made of glass, and so also has a high heat capacity, leading to similar problems as with the beaker. Another inaccuracy in the experiment was the time at which the temperature was taken – it was taken at the end of the reaction, which will have resulted in a temperature change which will have been less than it should have been. This can be counteracted by taking the peak temperature reached. The problem of heat loss to the apparatus could also be reduced by performing the experiment in an insulating container – for example; a polystyrene cup would dramatically reduce the amount of heat lost. As I mentioned above, apart from the beaker, the thermometer itself also absorbs a large amount of the thermal energy. This could be avoided by using a temperature-measuring device (such as a thermistor) with a low heat capacity, so that minimal heat is lost to it. Another major problem of the experiment would be the heat that is lost not to the apparatus, but to the surroundings. This is a problem, as it means that the reading on the thermometer even if it did have a low heat capacity, would still not be the same as the actual change, as some is lost to the surroundings. We also had a major problem with the time at which the temperature was taken – we were instructed to take the temperature at the end of the reaction, by which time it had fallen significantly as the solution had cooled whilst the last remnants of the reaction were taking place. A way to prevent this would be to take the peak temperature on the thermometer, as this is the maximum measurable value, so must be the maximum measurable enthalpy change. There was yet another problem with the scale of the thermometer – we were using 0-100ºC thermometers to measure relatively small changes - 2ºC for calcium carbonate. This is a problem, as it is harder to read off accurately the temperature change when using a big scale and experiencing a small change – there is a large possible percentage error. This obviously has a large impact on our calculations, as the temperature change must be recorded as accurately as possible in order to get a correct representation of the enthalpy change. A method of combating this would be to use a 0-50ºC thermometer, as this has smaller gradations on it, and so it is easier to read off the values. It may be even better to use a digital thermometer, which has been professionally calibrated to a high degree of accuracy – this minimises the human error in taking a reading from the thermometer.
There were also a number of other problems with the experiment, but these were of lower significance than those mentioned above. The most easily noticed one of these is the measuring cylinder – we were instructed to use a 250cm3 measuring cylinder to measure out 50cm3 of 2 mol dm-3 HCl. This could have produced an error, as it is using a large cylinder to measure out a relatively small amount of acid. However, whilst meaning that this variable was not entirely under control during the experiment, as the acid was in excess, it is not especially crucial that exactly 50cm3 was used – it does not affect the readings taken. Another variable was the concentration of the acid. This could have affected the results, by the concentration not being accurate. If this were the case, however, then the error would have been systematic, and so will have appeared in both sets of results, as the acid was drawn from the same bottle. Another one of these small errors is the accuracy of the balances – the readings given may not have been entirely accurate. However, this slight inaccuracy (±0.004g) pales in the face of the other experimental problems, so it can be safely disregarded.
These problems were not the only ones with the experiment – the major factor being that it was only performed once, with no repeats – these values were taken to be correct, with no comparisons made. This could easily be rectified by performing a suitable number of repetitions – for example, 4 repetitions could be made, and an average taken. This would vastly improve the reliability of the end results, as the average would more accurately reflect the true temperature change.
Overall, there were a large number of problems with this experiment, and correspondingly, there are a large number of things that I would like to change if I were to be able to repeat this experiment. The experiment was successful I that results were obtained, but I suspect that these results are vastly different to the actual values.