The Enthalpy Change of the Thermal Decomposition of Calcium Carbonate
Results:
For CaCO3: T1 = 17
T2 = 19
DT= 02
using 2.57g of CaCO3
For CaO: T1 = 18
T2 = 27
DT= 09
using 1.39g of CaO
Analysis:
In order to determine the enthalpy change for the thermal
decomposition of calcium carbonate, we must work out the enthalpy
changes for both the reactions of calcium carbonate and calcium oxide
with hydrochloric acid.
For CaCO3:
Temperature change = 2ºC
To find the enthalpy change of a reaction, we must first work out the
amount of energy taken in by the reaction. This is done by using the
following formula:
E=DT x mass surroundings x specific heat capacity of surroundings
For this calculation, we will assume that the specific heat capacity
of HCl is identical to that of water, and that the shc of water is
4.2J/ºC/g. We used 51cm3 of HCl, so the mass of this is taken to be
51g, as 1cm3 of water weighs 1g (and we are assuming that HCl(aq) has
the same density as water). So, putting this data into the equation,
we get:
E= (-2) x 51 x 4.2
= -428.4J
Then, in order to find the enthalpy change for this reaction, this
value should be converted into kJ, and divided by the number of moles
of the substance (in this case, calcium carbonate). To find the number
of moles used, we divide the mass used by the relative atomic mass of
the substance. So, we get:
No moles = Mass ¸Mr
= 2.57 ¸100.08
= 0.026 mol.
Then we convert our energy into kJ rather than joules, getting 0.4284
kJ, and dividing this by the amount of substance we used. So, this
comes out as:
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...sregarded.
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.
There were no significant error factors that may have affected the arrangement of the lab experiment. Everything went smoothly with relative ease.
For the sample calculations, let’s use the marshmallow as an example. Its initial mass was 0.66 grams and its final mass was 0.36 grams. To calculate the amount burned, subtract 0.36 from 0.66 to get 0.30 grams. (Mass burned = mi- mf). To find the marshmallow’s change in temperature, use the formula (ΔT =
So, to determine the change in enthalpy we will employ Hess’s Law of heat summation: It states that the value of DH for a reaction is the same whether it occurs directly or as a series of steps (LeMay et al, 1996). We will perform the two following reactions: Mg + 2HCl ® MgCl2 + H2 and. MgO + 2HCl ® MgCl2 + H2O, determine their enthalpy changes (DHs), and they will then be “added” to that of a given equation, the combustion of water, H2 + 1/2 O2 ® H2O
Okay, if our lithium weight is going to be 6.941 g/moL Then that means we have to take 24.6g of Lithium and multiply it by 1 mol of Lithium over 6.941 g of Lithium. This would equal to be 3.544 mol of Lithium. Then we have to take that 3.544 and multiply it by 1 mol of hydrogen gas over 2 mol of lithium. Which would then equal into 1.772 mol of hydrogen gas. We can then figure out that 1.772 is our “n”. The “T” is our 301 Kelvin, the “P” is our 1.01 atm and the “R” is our 0.0820 which would be the L atm over mol k. And we can’t forget about our “V” which would be V equals nRT over P which equals 1.772 mol divided by 0.0820 L atm over mol kelvin multiplied by 301 kelvin over 1.01 atm which equals to our final answer of: 43.33 of H2
Possible sources of error in this experiment include the inaccuracy of measurements, as correct measurements are vital for the experiment.
Introduction: The purpose of this lab was to cycle solid copper through a series of chemical forms and return it to its original form. A specific quantity of copper undergoes many types of reactions and goes through its whole cycle, then returns to its solid copper to be weighted. We observed 5 chemical reactions involving copper which are: Redox reaction (which includes all chemical reactions in which atoms have their oxidation state changed), double displacement reaction, precipitation reaction, decomposition reaction, and single displacement reaction. 4HNO3(aq) + Cu(s) --> Cu (NO3)2(aq) + 2H2O (l) + 2NO2(g) Oxidation reduction reaction Cu (NO3)2(aq) + 2 NaOH (aq) --> Cu (OH)2(s) + 2 NaNO3(aq) Precipitation Reaction Cu (OH)2(s) + heat --> CuO (s) + H2O (l) Decomposition reaction CuO (s) + H2SO Data Results: (mass of copper recovered / initial mass of copper) x 100 Mass of copper recovered: 0.21 Initial mass of copper: 0.52 (0.21/0.52)x100 =40.38%.
Hydrochloric acid + calcium carbonate arrow calcium chloride + carbon dioxide + water. HCl(aq) + CaCO3(s) arrow CaCl2(aq) + CO2(g) + H2O(l) Things that affect the reaction rate of this experiment are: 1. The temperature of the hydrochloric acid. 2.
The Effect of Temperature on the Rate of Reaction Between Hydrochloric Acid and Calcium Carbonate
Moles Volume HCl Volume Water 2 M 10 cm 3 0 cm 3 1.5 M 7.5 cm 3 2.5 cm 3 1 M 5 cm 3 5 cm 3 0.5 M 2.5 cm 3 7.5 cm 3
Text Box: CaCO3 + HCl = CaCl2 + CO2 + H2O calcium carbonate + hydrochloric acid = calcium chloride + carbon dioxide + water
* Note the mass down in the table at the end of the first page.
In this lab, I determined the amount of heat exchanged in four different chemical reactions only using two different compounds and water. The two compounds used were Magnesium Hydroxide and Citric Acid. Both compounds were in there solid states in powder form. Magnesium Hydroxide was mixed with water and the change in heat was measured using a thermometer. The next reaction combined citric acid and magnesium hydroxide in water. The change in heat was measured as well. For the third reaction citric acid was placed in water to measure the change in heat. In the last reaction, citric acid was combined with water. The heat exchanged was again measured. It is obvious we were studying the calorimetry of each reaction. We used a calorimeter
" This means that therefore the enthalpy change of a reaction can be measured by the calculation of 2 other reactions which relate directly to the reactants used in the first reaction and provided the same reaction conditions are used, the results will not be affected. We have the problem set by the experiment to determine the enthalpy change of the thermal decomposition of calcium carbonate. This is difficult because we cannot accurately measure how much thermal energy is taken from the surroundings and provided via thermal energy from a Bunsen flame into the reactants, due to its endothermic nature. Therefore, using the enthalpy changes obtained in reaction 1 and reaction 2 we can set up a Hess cycle.
is 40.c but his is a chemical enzyme so it will work best a little
There is also the potential of human error within this experiment for example finding the meniscus is important to get an accurate amount using the graduated pipettes and burettes. There is a possibility that at one point in the experiment a chemical was measured inaccurately affecting the results. To resolve this, the experiment should have been repeated three times.