Background

The enthalpy of hydration is the energy change when 1 mol of an anhydrous compound turns to 1 mol of a hydrated compound. The energy wchange when the stoichiometric amount of water is incorporated into the ionic crystal lattice.

This reaction is not possible to carry out stoichiometrically, but the energy change can be determined indirectly using Hess' law, providing that both forms of the salt are soluble.

Magnesium sulfate exists in both hydrated and anhydrous forms. The anhydrous form is easily prepared by placing the hydrated crystals in an oven at 200ºC for several hours.

Dehydration of magnesium sulfate crystals

MgSO₄.7H₂O(s) → MgSO₄(s) + 7H₂O(g)

Both anhydrous magnesium sulfate and magnesium sulfate heptahydrate dissolve in water making magnesium sulfate solution.

Dissolution of magnesium sulfate

MgSO₄.7H₂O(s) + (aq) → MgSO₄(aq) ← MgSO₄ + (aq)

Hence, to determine the enthalpy change of hydration or dehydration a simple Hess' cycle can be used.

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Requirements

Chemicals

• Anhydrous magnesium sulfate, MgSO₄
• Magnesium sulfate heptahydrate, MgSO₄.7H₂O
• Water

Apparatus

• Polystyrene beaker and lid
• Glass beaker, 250ml
• Measuring cylinder, 50ml
• Weighing boat
• Spatula
• Stirring rod
• Electronic balance
• Thermometer or temperature probe

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Procedure

• Measure about 50g water into a pre-weighed calorimeter (polystyrene beaker).
• Record the temperature and the mass.
• Accurately weigh about 5g of hydrated magnesium sulfate, MgSO₄.7H₂O.
• Dissolve the salt as rapidly as possible with stirring and record the highest (or lowest) temperature attained.
• Repeat using about 2g of anhydrous magnesium sulfate, MgSO₄, accurately weighed.
• Ensure that you measure the enthalpy change for both salts at least three times

Note: The anhydrous magnesium sulfate is kept in a dessicating jar to prevent absorption of water from the atmosphere.

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