Application of Hess's Law to determine the enthalpy change of hydration of Magnesium sulphate

2. Determine the enthalpy change of solution of MgSO4• 7H2O(S)

1. 50 cm3 of deionized water was poured from the measuring cylinder to the polystyrene cup again
2. Using a thermometer, the temperature of the water in the cup was measured
3. 0.025 mole of MgSO4• 7H2O(S) was weighed accurately by the balance instead of MgSO4(S)and was added into the foam cup
4. The solute was stirred to make sure all anhydrous magnesium sulphate (VI) were completely dissolved into the water as quickly as possible
5. The highest temperature of the solution was taken down
6. The molar enthalpy change of solution of MgSO4(S) was then calculated

( Assume:

The specific heat capacity of the solution in the foam cup = 4.2kJkg-1K-1 

The specific heat capacity of the foam cup = 1.3kJkg-1K-1                  )

Result:

Weight of the polystyrene cup = _________ g

Calculation: 

 1. The following assumptions were made:

     The specific heat capacity of the solution in the foam cup = 4.2kJkg-1K-1          

The specific heat capacity of the foam cup = 1.3kJkg-1K-1                                

         The thermal capacities of thermometer is negligible                        

     Density of the solution = 1gcm-3                                    

2.    Calculate the molar enthalpy change of hydration of MgSO4(S)

ΔH for the reaction: MgSO4(S)  + 7H2O(S) → MgSO4• 7H2O(S)

Using Hess’s Law ,

ΔH1 = ΔH2 – ΔH3

     ΔH2 :

E = m c ΔT

         = [ (      g ) ( 4.2kJkg-1K-1  ) + (      g ) ( 1.3kJkg-1K-1 ) ]

[ (     + 273 ) – (     + 273 ) ] K

         =        J

      Number of moles of MgSO4 = 0.025 mol

           ΔH2 = (     J ) / ( 0.025 mol )

                =       kJ/mol

   molar enthalpy change of solution of hydrous MgSO4(S) = -     kJ/mol

      ΔH3 :

E = m c ΔT

         = [ (      g ) ( 4.2kJkg-1K-1  ) + (      g ) ( 1.3kJkg-1K-1 ) ]

[ (     + 273 ) – (     + 273 ) ] K

         =        J

      Number of moles of MgSO4 = 0.025 mol

           ΔH3 = (     J ) / ( 0.025 mol )

                =       kJ/mol

    Molar enthalpy change of solution of MgSO4• 7H2O(S) = -      kJ/mol

Using Hess’s Law ,

ΔH1 = ΔH2 – ΔH3

      = [ (         ) – (         ) ] kJ/mol

     =           kJ/mol

The true enthalpy change of the hydration of magnesium sulphate

   = -104.0 kJ

   Thus , the percentage error

= [ (        ) – (        ) ] / (         )

              =         %

Discussion:

3. Point out some factors leading to the difference between the experimental value and the true value

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4. Explain why the enthalpy change of the hydration of magnesium sulphate cannot be measured directly in the laboratory

     The reason is that this reaction, hydration, is a very slow process which is    

  impossible to obtain directly in the laboratory                          

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5. Explain why it is not necessary to plot a temperature-time graph to determine the ΔT

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6. Draw an enthalpy level diagram for the reactions involved in the enthalpy cycle used