The Calorimetric Determination of Enthalpy and Entropy Changes of the Thermal Decomposition of Sodium Sulphate Decahydrate
Abstract
Sodium sulphate decahydrate thermally decomposes to anhydrous sodium sulphate at 50⁰ C and cools to form anhydrous sodium sulphate. Using Hess's Law the enthalpy and entropy changes of this process can be determined in order to deduce the spontaneity of the reaction. The transition temperature was determined by melting the hydrated crystals at 50⁰ C and allowing them to cool until a constant temperature was reached. The molar enthalpy and molar entropy changes were determined using calorimetric principles. Different weights of both decahydrate and anhydrous crystals were added to water and the temperature changes …show more content…
Materials and Methods
The transition temperature of sodium sulphate decahydrate was determined by allowing 5.01g of solid crystals to melt at a temperature of 50 ⁰C. The substance was allowed to cool and the temperature monitored until a constant temperature was reached. A plot of temperature versus time allowed the transition temperature to be determined.
The water equivalent of the calorimeter was determined using distilled water. Cool water was placed into the calorimeter and its temperature recorded, while an equal weight of water was heated to 35.3 ⁰C. The heated water was immediately poured into the water in the calorimeter and the temperature was recorded. The water equivalent of the calorimeter could then be calculated.
Approximately 100 g of cool distilled water was placed into the calorimeter and the initial temperature recorded. Solid sodium sulphate decahydrate crystals of 4.00 g were added to the water in the calorimeter and stirred until it dissolved. The final temperature of the water was then recorded. This process was repeated using 1.03 g and 2.07 g of sodium sulphate decahydrate, and 1.01 g, 1.50 g and 2.02 g of anhydrous sodium …show more content…
These values were then converted into enthalpy values in joules per mole. The enthalpy values for the hydrated crystals show a trend of positive values while those of the anhydrous crystals show a trend of negative values. The total enthalpy, however, is positive. Enthalpy is related to the spontaneity of a process. The entropy of the process was determined from the transition enthalpy and temperature values. Based on the positive enthalpy and entropy values, the reaction appears nonspontaneous.
The Gibbs free energy, however, is a better indicator or spontaneity. In order to calculate Gibbs free energy, As the temperature increases the Gibbs free energy decreases. This shows a movement towards disorder with increase of temperature. The reaction is nonspontaneous at low temperature. At the transition temperature of 31 ⁰C the free energy is low but still positive. The reaction is not spontaneous at this point as it is slowing down to allow the anhydrous crystals to form. At 50 ⁰C the reaction is highly negative and spontaneous.
In conclusion, the transition enthalpy and entropy values tend to show a nonspontaneous reaction. The Gibbs free energy, however, shows that the reaction is spontaneous at high temperatures when the decahydrate crystals melt. The reaction proceeds to become endergonic as the temperature
Abstract
Sodium sulphate decahydrate thermally decomposes to anhydrous sodium sulphate at 50⁰ C and cools to form anhydrous sodium sulphate. Using Hess's Law the enthalpy and entropy changes of this process can be determined in order to deduce the spontaneity of the reaction. The transition temperature was determined by melting the hydrated crystals at 50⁰ C and allowing them to cool until a constant temperature was reached. The molar enthalpy and molar entropy changes were determined using calorimetric principles. Different weights of both decahydrate and anhydrous crystals were added to water and the temperature changes …show more content…
Materials and Methods
The transition temperature of sodium sulphate decahydrate was determined by allowing 5.01g of solid crystals to melt at a temperature of 50 ⁰C. The substance was allowed to cool and the temperature monitored until a constant temperature was reached. A plot of temperature versus time allowed the transition temperature to be determined.
The water equivalent of the calorimeter was determined using distilled water. Cool water was placed into the calorimeter and its temperature recorded, while an equal weight of water was heated to 35.3 ⁰C. The heated water was immediately poured into the water in the calorimeter and the temperature was recorded. The water equivalent of the calorimeter could then be calculated.
Approximately 100 g of cool distilled water was placed into the calorimeter and the initial temperature recorded. Solid sodium sulphate decahydrate crystals of 4.00 g were added to the water in the calorimeter and stirred until it dissolved. The final temperature of the water was then recorded. This process was repeated using 1.03 g and 2.07 g of sodium sulphate decahydrate, and 1.01 g, 1.50 g and 2.02 g of anhydrous sodium …show more content…
These values were then converted into enthalpy values in joules per mole. The enthalpy values for the hydrated crystals show a trend of positive values while those of the anhydrous crystals show a trend of negative values. The total enthalpy, however, is positive. Enthalpy is related to the spontaneity of a process. The entropy of the process was determined from the transition enthalpy and temperature values. Based on the positive enthalpy and entropy values, the reaction appears nonspontaneous.
The Gibbs free energy, however, is a better indicator or spontaneity. In order to calculate Gibbs free energy, As the temperature increases the Gibbs free energy decreases. This shows a movement towards disorder with increase of temperature. The reaction is nonspontaneous at low temperature. At the transition temperature of 31 ⁰C the free energy is low but still positive. The reaction is not spontaneous at this point as it is slowing down to allow the anhydrous crystals to form. At 50 ⁰C the reaction is highly negative and spontaneous.
In conclusion, the transition enthalpy and entropy values tend to show a nonspontaneous reaction. The Gibbs free energy, however, shows that the reaction is spontaneous at high temperatures when the decahydrate crystals melt. The reaction proceeds to become endergonic as the temperature