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
Reference Tools
Periodic Table
Molarity Calculator
Molar Mass Calculator
Unit Converter
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
hydration. A hydrated compound formula includes the term •nH2O, where n is the number of moles of water incorporated into the solid per mole of ionic compound. Prefixes are used in naming hydrated compounds to indicate the number of waters of hydration.
Many solids used in the laboratory are hydrated. For example, reagent-grade copper(II) sulfate is usually provided as the hydrated compound CuSO4•5H2O, copper(II) sulfate pentahydrate. Some common hydrated compounds and their uses are shown in Table 3.2.1. Notice that the molar mass of a hydrated compound includes the mass of the water of hydration.
Table 3.2.1: Some Common Hydrated Ionic Compounds
Molecular formula
Name
Molar mass (g/mol)
Common Name
Uses
Na2CO3•10H2O
Sodium carbonate decahydrate
286.14
Washing soda
Water softener
Na2S2O3•5H2O
Sodium thiosulfate pentahydrate
248.18
Hypo
Photography
MgSO4•7H2O
Magnesium sulfate heptahydrate
246.47
Epsom salt
Dyeing and tanning
CaSO4•2H2O
Calcium sulfate dihydrate
172.17
Gypsum
Wallboard
CaSO4•½H2O
Calcium sulfate hemihydrate
145.15
Plaster of Paris
Casts, molds
CuSO4•5H2O
Copper(II) sulfate pentahydrate
249.68
Blue vitriol
Algicide, root killer

Heating a hydrated compound releases the water in the crystalline solid. For example, heating the compound CuCl2•2H2O releases two moles of water (in the form of water vapor) per mole of hydrated compound. As shown in the following example problem, we can determine the formula of a hydrated compound by performing this experiment in the laboratory.
Example Problem: Determine the formula of a hydrated compound.
A 32.86 g sample of a hydrate of CoCl2 was heated thoroughly in a porcelain crucible until its weight remained constant. After heating, 17.93 g of the dehydrated compound remained in the crucible. What is the formula of the hydrate?

You are asked to determine the formula of an ionic hydrated compound.
You are given the mass of the hydrated compound and the mass of the compound when it has been dehydrated.
First, determine the mass of water lost when the hydrated compound was heated.
32.86 g – 17.93 g = 14.93 g H2O
Next, calculate moles of water and moles of the dehydrated compound.

Finally, determine the simplest whole-number ratio of water to dehydrated compound.

The chemical formula for the hydrated compound is CoCl2•6H2O.

3.2.5T: Tutorial Assignment
Due: 1/1/20 11:59 PM
Score: 0

3.2.5: Mastery Assignment
Due: 1/1/20 11:59 PM
Score: 0

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Experiment 10 Hydrates
PRIOR READING
Be sure to read and understand the material in chapters 2 (page 68), section 2.2, and the mole concept in chapter 3.
INTRODUCTION
Hydrates are crystalline solids that contain a fixed number of water molecules as an integral part of their crystalline structure. The number of water molecules bound per metal ion is often characteristic of that particular metal ion. One of the more common hydrates is copper(II) sulfate pentahydrate, which contains 5 moles of water per 1 mol of copper(II) sulfate, written as CuSO45H2O. It is used as a catalytic precursor, fungicide, and as a source of copper in chemical manufacturing processes. Epsom salt is magnesium sulfate heptahydrate, MgSO47H2O. Epsom salt is used to reduce inflammation when applied externally.
Many hydrates can be transformed to the anhydrous compound when heated strongly. For example, copper sulfate pentahydrate can be converted into anhydrous copper sulfate. This change can be followed visually. The blue crystalline copper sulfate pentahydrate is converted when heated to a white, powdery, anhydrous salt, according to:
CuSO45 H2O  CuSO4 + 5 H2O blue white
Or generally hydrated salt anhydrous salt + water vapor
It is also possible to reverse the above process, as shown in the equation below:
CuSO4 + 5 H2O  CuSO45 H2O white blue
If water is added to the white anhydrous copper sulfate, a blue color is obtained indicating that the blue pentahydrate is regenerated. The property of reversibility can be used to distinguish true hydrates from other compounds that produce water when heated.
Since many hydrates contain water in a stoichiometric quantity, it is possible to determine the molar ratio of water to salt. This is exactly what you will be doing in Part II of today’s experiment. We will begin by heating an accurately weighed sample of the hydrate to drive out the water. The compound formed is now anhydrous. By determining the mass of the anhydrous sample and subtracting this mass from that of the hydrate, we can determine the amount of water in the original substance.
In part I of this experiment, we will look at properties of several crystalline hydrates.
Some anhydrous salts are capable of becoming hydrated on exposure to the moisture in their surroundings. These salts are called hygroscopic and can be used as chemical drying agents or desiccants. Some salts are some excellent desiccants and are able to absorb so much moisture from their surroundings that they can eventually dissolve themselves. These salts are called deliquescent.

MATERIALS
Nickel(II) chloride hexahydrate
Cobalt(II) chloride hexahydrate
Copper(II) sulfate pentahydrate
Potassium aluminum sulfate dodecahydrate
Porcelain crucible and lid
Unknown hydrate sample for percentage determination
EXPERIMENTAL PROCEDURE
Record all data and observations directly into your notebook.
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
0.001g.
Obtain an unknown hydrate from your Instructor. Record the number of the unknown in your notebook. Place about 1 to 2 g (to the nearest 0.001g) of the unknown hydrate into the crucible. Weigh the crucible, lid, and unknown to the precision indicated above.
Place the crucible and its lid onto the clay triangle. Arrange the lid so it is slightly ajar and begin heating very gently for about five minutes. If heated too strongly, some of the sample may decompose. Beware of spattering during the heating process. If spattering occurs put the lid on the crucible and remove the flame. Continue heating for another 10 minutes. Cover the crucible with the lid, cool to room temperature, determine the mass of the anhydrous salt in the crucible, and report it in your notebook.
Repeat the heating-cooling cycle a second time, this time heating the crucible for another 5 minutes. Determine the mass of the anhydrous salt. Continue the heating-cooling cycle until two successive weighing agree within 0.040 grams.

Calculate the mass percent water in the hydrate based upon the mass of the hydrate and that of the anhydrous salt. Then determine the molar composition of the unknown compound.
WASTE DISPOSAL
Dispose of any used and unused chemicals in the proper waste container located in the hood. Solids are placed in the container labeled “Solid Waste” and all other solutions in the beaker.
Prelaboratory Assignment: Hydrates
1. In an experiment, 2.3754 g of copper(II) sulfate pentahydrate is heated to drive off all the water of crystallization.
a) Write a balanced equation.
b) Determine the mass of anhydrous salt that remains. Show your calculations.
2. Cobalt (II) chloride is commonly obtained from chemical supply houses as a hydrate with the formula CoCl2 6H2O. What is the percent water by weight in this hydrate?
3. A hydrate of sodium phosphate, Na3PO4, contains 49.7% water by weight.
a) How many grams of water and how many grams of anhydrous Na3PO4 are in 1000. grams of this sample?
b) In this same 1000.-gram sample, how many moles of water and how many moles of anhydrous Na3PO4, are present?
c) What is the formula of the hydrate?

Data and Analysis: Hydrates
Part I: Properties of Hydrates
Record and analyze any observations .
Create a chart and list any observations (color, appearance before and after heating, etc.)
Distinguish between the two terms absorbed and hydrated water.
Explain how do you would test a colorless crystalline compound to determine if it was a hydrate?
Part II: Formula of a Hydrate
Compare data with your partner and discuss any discrepancies and deviations in our results and observations.
How would the following observations influence the percent by weight of water in the original salt mixture? Will your value be too large, too small, or not affected? Please explain your answer.
a) Your anhydrous substance was slightly warm when you weighed it.
b) You did not heat long enough and the hydrate was not completely converted to the anhydrous substance.
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Objectives
• Identification of hydrates in a group of compounds
• Investigation of the properties of hydrates
• Determination of the number moles of water of hydration in a hydrate

Background
Water, the most common chemical on earth, can be found in the atmosphere as water vapor.
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
compound, will have a different structure and texture and may have a different color than the hydrate. Example:
CuSO4•5H2O (s)
⎯→⎯Δ
CuSO4 (s)
+
5 H2O (g) deep blue ashy white
CuSO4 (s) ashy white
⎯⎯⎯→⎯)(2lOH
CuSO4 (aq) deep blue