Stoichiometry Lab Report

Chapter 3: Stoichiometry
3: Stoichiometry
5: Thermochemistry
8: Covalent Bonding and Molecular Structure
15: Chemical Equilibrium
16: Acids and Bases

3.2 Stoichiometry and Compound Formulas
3.1 The Mole and Molar Mass
3.2 Stoichiometry and Compound Formulas
3.3 Stoichiometry and Chemical Reactions
3.4 Stoichiometry and Limiting Reactants
3.5 Chemical Analysis
Chapter Summary
Chapter Summary Assignment
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Thermodynamic Data
3.2e Hydrated Compounds

A hydrated ionic compound is an ionic compound that has a well-defined amount of water trapped within the crystalline solid. The water associated with the compound is called the water of
Part I: Properties of Hydrates
1.Place about 0.1 g of the following compounds in each one test tube:
CuSO45 H2O, CoCl26 H2O, NiCl26 H2O, and KAl(SO4)212 H2O.
2. Heat each test tube gently over a Bunsen burner flame and record your observations in your notebook.
3. After the sample has cooled, add a few drops of deionized water. What happens and what can be concluded?
Part II: Formula of a Hydrate
You and your partner will perform two trials of dehydration of a copper (II) sulfate hydrate.
During the course of the experiment, handle the crucible and lid only with crucible tongs as shown here or as demonstrated by your instructor.

Clean two crucibles with soap and water. Rinse the crucibles with distilled water and dry them with a paper towel. Check your crucible for cracks.
Heating your crucible first without the hydrate.
Prepare two set-ups as shown below using a clay triangle on a ring stand. Place each crucible on a clay triangle and heat the crucibles until red hot or for five minutes. Once the heating is complete, place the crucible on a clean wire gauze and let it cool to room temperature. Determine the mass of the crucible and lid to the nearest
Some chemicals, when exposed to water in the atmosphere, will reversibly either adsorb it onto their surface or include it in their structure forming a complex in which water generally bonds with the cation in ionic substances. The water present in the latter case is called water of hydration or water of crystallization. Common examples of minerals that exist as hydrates are gypsum (CaSO4•2H2O), Borax (Na3B4O7•10H2O) and Epsom salts (MgSO4•7H2O). Hydrates generally contain water in stoichiometric amounts; hydrates’ formulae are represented using the formula of the anhydrous (non-water) component of the complex followed by a dot then the water (H2O) preceded by a number corresponding to the ratio of H2O moles per mole of the anhydrous component present. They are typically named by stating the name of the anhydrous component followed by the Greek prefix specifying the number of moles of water present then the word hydrate (example: MgSO4•7H2O: magnesium sulfate heptahydrate).
Properties of Hydrates
It is generally possible to remove the water of hydration by heating the hydrate. Le Chatelier’s principle predicts that an addition of heat to an endothermic reaction (heat is a “reactant”) will shift the reaction to the right (product side). Heating will shift the equation of dehydration below to the right since it is an endothermic reaction. The residue obtained after heating, called the anhydrous