boiling point of anhydrous sodium sulfate

The boiling point of a substance depends on the temperature at which it changes from one state to another. For example, if sodium sulfate is free of water and is allowed to come in contact with aqueous solutions, it will absorb moisture from those solutions and turn into a decahydrate form.

Sodium sulfate is a chemically stable and unreactive substance and will not react with most oxidizing or reducing agents at normal temperatures. Consequently, it is an excellent dehydrating agent.

In the laboratory, it is commonly used to remove traces of water from organic solutions. Its boiling point varies widely, depending on the concentration of the solution.

Its solubility increases rapidly at 0degC and decreases at higher temperatures. It is dissolved in about 33g per 100g water at 30degC, and in 2.4 parts of water at 100degC.

Natural sources of sodium sulfate exist in many salt ponds and surface depressions where spring waters run over rock containing sodium sulfate minerals, such as magnesium sulfate (calcium) double salt or mirabilite. Approximately 3.3 billion tons of economic reserves are estimated to exist worldwide.

Sodium sulfate-bearing mineral deposits are geologically young and widespread in occurrence. They are mainly of post-glacial age and are found in seawater and several saline or alkaline lakes.

The primary source of industrial production is the Mannheim process, which uses sulfuric acid and sodium chloride in anhydrous form, or the Hargreaves process, which uses sulfur dioxide. Alternatively, it can be produced during hydrochloric acid production by neutralizing excess sulfuric acid with sodium hydroxide.