“Epsom salts” is the common name of magnesium sulphate heptahydrate: MgSO4.7H2O. Its common name originated from the name of an English town - Epsom and Ewell - known for its mineral springs. Many chemists and supermarkets stock Epsom salts on their drug aisle as an over-the-counter treatment for sprains and sore muscles. Epsom salts do not pose a serious health risk unless ingested in large quantities and do not exhibit dangerous reactivity with other substances. They, therefore, provide a suitable “starter” chemical for aspiring young scientists to experiment with.
Other People Are Reading
When salts dissolve in water, the positively and negatively charged components - known as “ions” - separate. Epsom salts dissolve to form positively charged magnesium ions and negatively charged sulphate ions. This separation of ions is reversible. That is, if the water is removed, such as by evaporation, then the ions from the salt will recombine to form crystals. Drexel University provides a procedure for seventh- or eighth-grade students for crystallising Epsom salts on the end of a string.
Princeton University has published a method of preparing ammonium sulphate and magnesium hydroxide using Epsom salts. The technique simulates the extraction of magnesium from seawater and calls for the combination of Epsom salts, household ammonia, which produces ammonium sulphate, a water-soluble substance, and magnesium hydroxide, which is not water-soluble and precipitates as a solid. The experimenter then removes the magnesium hydroxide by filtration and recovers the ammonium sulphate by evaporating the water from the solution. Because the technique involves potentially hazardous chemicals, it is best suited to high school chemistry students.
Determining the Empirical Formula of Epsom salts
Many first-year college chemistry students conduct an experiment to determine the formula of a hydrated salt, such as magnesium sulphate heptahydrate. Los Angeles City College publishes such a procedure on their website. The procedure calls for weighing a salt sample before and after heating the sample to dryness. The difference between the two masses represents the amount of water contained in the sample. A simple calculation converts that mass of water to moles of water, and should reveal a 7:1 ratio of water to magnesium sulphate. Because the procedure involves open flames and requires a working knowledge of the mole concept, the experiment is most appropriate for advanced highschool or first-year college chemistry students.
- 20 of the funniest online reviews ever
- 14 Biggest lies people tell in online dating sites
- Hilarious things Google thinks you're trying to search for