Many salts that have been crystallized from water solutions appear to be perfectly dry, yet when heated they discharge large quantities of water. An example can be hydrated copper (II) sulfate. Exactly this salt is used in the described experiment.
What is the number of moles of water of crystallization associated with one mole of copper (II) sulfate, in the hydrate CuSO4 * xH2O (s)? The independent variable in this experiment is sample thermal treatment ( heating and cooling) and the dependent variable is the number of moles of water of crystallization. A number of other variables can affect the final result of the experiment: whether the crucible was clean and dry before filling it with the salt, whether the sample was heated too strongly and cooled too long before weighting.
In the hydrate copper (II) sulfate one mole of salt is believed to be combined with five moles of water. Bearing this in mind the following hypothesis was made: If the hydrate is heated until there is no further loss in weight, five moles of water will be lost from one mole of salt.
When certain ionic solids crystallize from aqueous solutions, they combine with a definite amount of water, which becomes a part of the crystalline solid. Salts that contain water as part of their crystal structure are called hydrates (or hydrated salts) and the water in the crystal structure is called the water of hydration (or water of crystallization). When the water of hydration is removed from the hydrate, the salt that remains is said to be anhydrous. For a hydrate, the number of moles of water present per mole of salt is usually some simple, whole number. The formula for a hydrated salt is usually written as MN . x H2O, where MN means a salt, x indicates the number of water molecules that are associated with one molecule of salt and the dot means that water molecules are rather loosely attached to the salt. In this experiment, the mass of the water driven off, as well the mass of the salt remained was determined. Then using the molar mass of CuSO4, the amount in moles of water and the anhydrous salt was calculated. Finally given the mole ratio, the formula of the hydrated salt was founded. The essential formulae for this experiment were:n= m : M n- number of molesm- mass of a substance [g]M- molar mass [g/mol]% error= I(A%-E%) : A%I * 100%A%- actual percentage of H2O in a hydrateE% – experimental percentage of H2O, in a hydrateII – absolute value
spectaclescrucible (50 ml)crystals of hydrated copper (II) sulfatesupport standburnerbalance (accuracy 0.05g)stirrertongstest tubepipette
During heating the spectacles should be worn.Hot crucibles should be carried with tongs, not with hands.The crucible should be cooled before weighting. 1.The empty, clean and dry crucible was taken and weighted (the tare button was used before placing the crucible on the balance). The result was recorded in a raw data table.2.The crucible was filled with a little amount of a hydrated copper sulfate (in a way, that the salt covered with a thin layer the bottom of a crucible).3.The crucible with the hydrated salt was weighted without using the tare button and the result was recorded in a raw data table.4.The crucible was placed on a support stand. The burner was lighted (a very light, almost transparent flame was set) and placed under the crucible. (see fig.1)5.The crucible was heated for about 10 minutes and the salt was stirred for the whole time. The observations were recorded.6.The crucible was taken with the tongs from the burner and left to cool.7.When the crucible was cool enough to hold it in the hands, it was reweighed (the tare button was used) and the result was recorded in a row data table.8.The crucible was heated again for about 8 minutes, cooled and reweighed. The result was recorded.9.
The crucible was heated again for about 2 minutes, cooled and reweighed. The result was recorded. There was no further loss in weight so no further heating was needed.10.A little amount of remained salt was placed in a test tube. A few drops of water were added with a pipette. The observations were recorded.11.The place of work was cleaned up.Figure 1: Deployment of the burner, the support stand and the crucible, which should be followed, so that the sample may be heated properly.
RESULTS DATA COLLECTION
The row data, which was collected during the experiment, is shown in Table 1. Additional observations made during the experiment are listed in Table 2. Table 1 : Changes in mass of the crucible from the moment it was empty until it was reweighed after third cooling.Number of weightingsRecorded mass [g] (+/- 0.05g)Notes129.34Mass of empty crucible231.05Mass of crucible and hydrate before heating330.46Mass of crucible and hydrate after first heating430.44Mass of crucible and hydrate after second heating530.44Mass of crucible and hydrate after third heatingTable 2 : Observations made from the time the crucible was filled with the hydrated salt until the end of the third heating.Number of heatingsObservations0Crystals are blue, approximately as big as sugar crystals.1Crystals are bluish, they are smaller2Crystals are white, they are very small (they form a powder)3No visible changes ( still white powder )Figure 2: Observations made after adding a few drops of water to the anhydrous salt.
1.Initial mass of the hydrated salt CuSO4 * xH2O (s) was calculated:mass of crucible and hydrate before heating – mass of empty crucible = mass of hydrate31.05 – 29.34=1.71 (g) (+/- 0.05g)2.The mass, and hence the number of moles of anhydrous salt that remained was calculated, using the molar mass of the anhydrous salt used:m= 30.44- 29.34 = 1.1 (g) (+/- 0.05 g)MCuSO4 = 63.5 + 32 + 4 * 16 = 159.5 (g/mol) (+/- 0.5 g/mol)n= 1.1 : 159.5 = 0.00690 (mol) (rounded to 3sf)3.The mass and hence the number of moles of water lost from the salt was calculated:m= 1.71- 1.1 = 0.61 (g) (0.05g)MH2O= 2 + 16 = 18 (g/mol) (+/- 0.5 g/mol)n= 0.61 : 18= 0.0339 (mol) (rounded to 3sf)4.
The number of moles of water that are combined with one mole of the salt was calculated:0.0339 : 0.0069 = 4.91 (rounded to 3sf)5.Bearing in mind, that the number of moles should be an integer, the formula of the hydrate was found: 4.91= 5 → CuSO4 * 5H2O6.The percentage of water in the sample of hydrate was calculated:0.61 : 1.71 * 100% = 35.7 % (rounded to 3sf)7.The actual percentage of water in a hydrated copper sulfate was calculated: (5 * 18) : (159.5 + 5 * 18) * 100% = 36.1% (rounded to 3sf)8.The percent error in this experiment was calculated:I (36.1% – 35.7 %) : 36.1% I * 100%= 1.11% (rounded to 3sf) Table 3 summarizes all results of the calculations made during data analysis.Table 3 : Results of all calculations made during data analysis.Quantity measuredResultInitial mass of the hydrate1.71 (g) (+/- 0.05g)The mass of anhydrous salt1.1 (g) (+/- 0.05g)
The number of moles of anhydrous salt0.00690 (mol) (rounded to 3sf)The mass of water driven off0.61 (g) (+/- 0.05g)The number of moles of water driven off0.0339 (mol) (rounded to 3sf)The number of moles of water combined with one mole of the salt4.91 (rounded to 3sf) ~ 5Percentage of water in the sample of hydrate35.7 % (rounded to 3sf)Actual percentage of water in the hydrate36.1% (rounded to 3sf)Percent error in the experiment 1.11% (rounded to 3sf)
CONCLUSION AND EVALUATION
The data support the alternate hypothesis that in a hydrated copper (II) sulfate one mole of salt is combined with five moles of water. The possible explanation for the observations follows. When the salt is heated, the temperature of the hydrate is so high, that water molecules evaporate. In result, the color and size of salt’s crystals changes. This explains, why the salt’ s crystals were white and smaller after heating. When the anhydrous salt is added to the water, water molecules attach themselves to hydrate’s ions. When solution is evaporated, water molecules are so strongly attached to the ions, that they remain and become incorporated into the crystal structure. This explains, why after adding a few drops to the anhydrous copper sulfate, the solution became blue. This means also, that the process of discharging water from a hydrate is reversible (the anhydrous salt can bind water molecules again). Although the data support the hypothesis, the number of moles of water in the sample, which was obtained from the calculations, wasn’t exactly the same as the actual value ( it had to be approximated).
The possible experimental errors, which can explain this are: systematic error – the measurements of the balance was accurate only to the first decimal placerandom errors- the mass of an anhydrous salt could be changed, because a little amount of the salt might possibly stay on a stirrer; the crucible and the contents were allowed to cool too long/ to short before finding their mass. The crucible should be cooled before weighing, not only because of safety, but also because of fact, that during cooling water may also evaporate. On the other hand after the crucible is cool, the weighing should be carried out as rapidly as possible, because the anhydrous salt remained is very hygroscopic, so it can bind the water from air- this results in change of mass of the salt. In order to significantly improve the accuracy of the result no stirrers during heating should be used. Then one of the possible random errors would be eliminated, because change in mass of the anhydrous salt wouldn’t occur. Instead of stirring, the crucible should be heated more gently, so the salt wouldn’t burn.
If possible, use of more accurate balance should be considered. This would lower the uncertainties of mass results. When the crucible is cooled, student’s reaction should be immediate- the sample should be weighed very quickly, but on the other hand weighting should not be carried out too early. It’s important to mention that the described technique cannot be used to determine the water of crystallization present in all salts. Some of them are unsuitable, because they are too resistant. They do not decompose before reaching their melting point. In those cases special techniques are needed. For example those hydrates can be dried in a melted form and then crystallized from anhydrous solvent.
Furthermore, gas hydrates are also unsuitable for described technique, because while heating both water and gas would be lost. In some cases dissolved gases would be dangerous, when heated under higher pressure and mixed with the air, for example ethylene discharged from ethanol would explode. That is why dehydration of ethanol is carried out by adding to it strong acid desiccant, for example sulfuric acid. Products of this reaction are water and either diethyl ether or ethylene.
1.http://en.wikipedia.org/wiki/Hydrate2.http://pl.wikipedia.org/wiki/Hydrat3.http://goto.glocalnet.net/ibweb/homepagelink.htm4.www.unit5.org/chemistry/christjs/9hydratelab.doc5.”A portfolio of Investigations for I.B. Chemistry”, Australia 19996.Internal assessment guide- material from biology seminaries