Analysis:

Raw Data:

Hydrate Magnesium SulfateBefore (g)/ ±0.001 g

After (g)/ ±0.001g

∆ Temperature (°C) ± 0.2 °C

Trial 1 12.857 9.548 0.8Trial 2 15.446 9.409 1.6

Anhydrous Magnesium SulfateBefore (g)/ ±0.001 g

After (g)/ ±0.001g

∆ Temperature (°C) ± 0.2 °C

Trial 1 12.512 9.375 8.2Trial 2 13.837 9.899 9.8

Processed Data:

Hydrated Magnesium SulfateMass (g)/ ±0.002g ∆ Temperature (°C) ± 0.2

°CTrial 1 3.309 0.8Trial 2 6.037 1.6

Anhydrous Magnesium SulfateMass (g)/ ±0.002g ∆ Temperature (°C) ± 0.2

°CTrial 1 3.137 8.2Trial 2 3.938 9.8

Heat produced in an endothermic equation (hydrated magnesium sulfate):MgSO4 .7H 2O(s)+aq→MgSO4 (aq )

Enthalpy=∆ H reaction

n(MgSO .7H2O )

Trial 1:Mass of water = 45g ± 1g (2.22%)Specific heat capacity of water = 4.18 J/g

∆T= 0.8 °C ± 0.2°C (25%)

n( MgSO47H2O )=m(MgSO4 .7H2O )

RAM(MgSO4 .7H2O )m(MgSO4 .7H2O )=3.309 g ± 0.001 g (0.03%)

n( MgSO4 . 7H2O)=3 .309 g ± 0. 001 g (0 .03% )

24 .31+32 . 07+16 .00×4 +7(1 .01×2+16 . 00)=0 .0134 moles±0 . 03 %

Enthalpy=45 g(±2 .22 % )×4 . 18Jg−1K−1×0 .8 K (±25 % )0 . 0134(±0 . 03 % )

=11200 Jmol−1

¿11. 2kJmol−1±27 . 25 %

Trial 2:Mass of water = 45g ± 1g (2.22%)Specific heat capacity of water = 4.18 J/g∆T= 1.6°C ± 0.2°C (12.5%)

m(MgSO4 .7H2O )=6 .037g±0 .001g(0 .02% )

n( MgSO4 . 7H2O)=6 . 037 g ± 0 .001 g (0 .02% )

24 .31+32 . 07+16 .00×4 +7(1 .01×2+16 . 00)=0 .0245moles±0. 02 %

Enthalpy=45 g(±2 .22 % )×4 . 18Jg−1K−1×1. 6 K (±12. 5 % )0 . 0245(±0 .02 % )

=12300Jmol−1

¿12 .3kJmol−1±14 . 74 %

Average Enthalpy of Hydrated Magnesium Sulfate:

=

11. 2kJ±3.1kJ (27 .25 % )+12 .3 kJ±1. 8kJ (14 .74 % )2

=11. 8kJmol−1±2 .5kJmol−1

Heat produced in an exothermic equation (magnesium sulfate):MgSO4(s )+aq→MgSO4 (aq)

Enthalpy=∆ H reaction

n(MgSO) .

Trial 1:

Mass of water = 45g ± 1g (2.22%)Specific heat capacity of water = 4.18 J/g∆T= 8.2°C ± 0.2°C (2.4%)

n( MgSO4 )=m(MgSO4 )

RAM(MgSO4 )m(MgSO4 )=3 . 137g±0.001g (0 .03% )

n( MgSO4 )=3 .137g±0 . 001g (0. 03% )120 . 38

=0 . 0261moles±0 . 03 %

Enthalpy=−45 g(±2.22 % )×4 .18 Jg−1K−1×8. 2 K (±2. 4 % )

0.0261(±0 .03 % )=−59100 J

mol−1

¿−59. 1kJmol−1±4 .65 %

Trial 2:

Mass of water = 45g ± 1g (2.22%)Specific heat capacity of water = 4.18 J/g∆T= 9.8°C ± 0.2°C (2.04%)

m(MgSO4 )=3 .938g±0 .001g (0 .03% )

n( MgSO4 )=3 . 938g±0. 001g (0 .03% )120 . 38

=0. 0327moles±0. 03 %

Enthalpy=−45 g(±2 . 22% )×4 . 18Jg−1K−1×9 . 8K (±2 .04 % )0 . 0327(±0 . 03 % )

=−56400J±4 .29 %

¿−56 . 4kJ±4 .29 %

Average Enthalpy of Anhydrous Magnesium Sulfate:

=

{−59 . 1kJ±2.7kJ (4 .65 % )}+{−56 .4 kJ±2 . 4kJ (4 . 29 % )}2

=−57 .8 kJmol−1±2 . 6kJmol−1

Hess's law states that the enthalpy change for any chemical reaction is independent of the route provided the starting conditions and final conditions, and reactants and products, are the same.

Law of conservation of energy:the enthalpy change in a complete cycle is 0 .

∆ H ( a)=∆ H (c )−∆ H (b )∴∆ H (a )=∆ H ( c )−∆ H (b)

∆ H (c )=enthalpy of anhydrousmagnesiumsulphate .∆ H (b )=enthalpy of hydrated magnesiumsulphate .

∆ H ( a)=(−57 .8 kJmol−1±2 . 6kJmol−1 )−(11. 8kJmol−1±2.5kJmol−1 )∆ H ( a)=−69 .6 kJ±5.1kJ (7. 32 % )

∆ H ( a)=−69 .6 kJmol−1±5 . 1kJmol−1 (7 .32% )The actual value for the ∆H of this reaction is −104 kJmol−1

. The difference with the literature value is fairly large, being:

{1−−69 .6kJmol−1 (±7 .32 %)

−104kJmol−1}∗100=33 .1%±7 .32 %

.

It deviates 33 .1%±7 .32 % from the accepted value for this reaction, which is a considerable amount.

The most likely reason, for being an error, in this experiment is the experiment not being performed under standard conditions. The conditions for the standard conditions for enthalpy changes are:

100 kPa

concentrations of 1 mol dm−3

for all substances all the substances in their standard states.

Evaluation:

The purity of the hydrated and anhydrous magnesium sulfate is most likely the major source of error. The anhydrous magnesium sulfate also could possibly have absorbed moisture from the air while being stored or taken out during the experiment. This will result in a smaller magnitude for both the enthalpy values of the hydrated magnesium sulfate and anhydrous magnesium sulfate, explaining the considerable amount of deviation. Thus giving a value (-69.6kJ) relatively far (33 .1 %±7 . 32 %) from its literature value.

Conducting the experiment in a dry environment can reduce this error. Storing the anhydrous magnesium sulfate in a dry environment, using pure magnesium sulfate, can also reduce the error.

Other sources of error could be from inaccurate extrapolation on the graph (it was done by eye), and a delay in the addition of the substance (hydrated/anhydrous magnesium sulfate) to water.

Extrapolating roughly, by eye, may not be so accurate. Resulting in a possibly greater or smaller temperature change than the actual value. This could be improved by finding the equation of a slope to accurately calculate the maximum temperature by extrapolating.

The substances, contained in a test tube with a small mouth, were added to water relatively slowly, over a few seconds. The delay in the addition of the substance (hydrated/anhydrous magnesium sulfate) to water will result in a reaction experiencing smaller temperature change than possible. It is possible because the reaction seemed to occur spontaneously when magnesium sulfate was added. Although the magnetic stirrer was used to help the speed of the reaction, the slow addition of the substances can be a factor affecting the reliability. This can be improved using a test tube with a bigger mouth, which would help the substances fall into the water relatively faster.