Hydrate Lab

Kimberly Graziano & Hyunjae Kim

Purpose

Experimental Question: How can we experimentally determine the formula of an unknown hydrate, A?

Testable Prediction: Our unknown hydrate may be a hydrate of copper(II) sulfate, magnesium sulfate, iron(III) chloride, or iron(III) nitrate. Observing our nitrate, it has a white crystalline structure, representing that similar to table salt. Since copper (II) sulfate is usually a bright blue due to Cu2+ presence, we can exclude that option from our prediction. Iron (III) chloride usually has a bright yellow appearance. Iron (III) sulfate has a purple tint to it, and has a crystalline structure. This means we can exclude these three options from our prediction. Magnesium sulfate, the only left option, is white in appearance which makes it a possible identification for our hydrate.  We believe our hydrate was magnesium sulfate, because the unknown hydrate was more closely related in physical appearance to that of magnesium sulfate, compared to the the three other options.

        Furthermore, in order to determine the exact name of the hydrate, we must find out the ratio between the anhydrate and water that are associated with the hydrate. By multiplying the mass of the anhydrate, which is magnesium sulfate in the experiment, with its molar mass, the number of moles present at the end can be determined. The number of water moles can also be known by repeating the same procedure, but with the molar mass of water instead. Once the numbers of moles of two substances are known, the ratio can be computed by dividing them. The ratios between molecules are in integers, but as this is an experiment, it will be more likely to acquire the ratio in decimal points. Once we know how much water is needed for each magnesium sulfate, we can then name the substance in MgSO4 x H2O, where x represents the ratio. This hydrate was previously mentioned in class to be magnesium sulfate heptahydrate. Thus, the ratio between water and magnesium sulfate will be close to being 7:1.

Data & Analysis

Initial Data: 

Calculated Data:

Observations:

The hydrate (A) was white and crystal-like at the beginning. The substance looked similar to table salt as the particles were clearly visible. As the hydrate was heated with a bunsen burner, the consistency of the substance changed from being crystal-like to chalk-like. It seemed as if the particles were partially melting. This indicated that more water evaporated progressively and the sizes of the particles decreased accordingly. Also, some of the hydrate popped out from the beaker due to the heating. Other group’s reaction did not seem to cause some of their data to get lost. Furthermore, the strength of the heat had to be changed once from low to high as there was only 15 minutes total to heat the substance.

Calculations:

In order to identify the unknown substance, it was necessary to research the known ratios of all four choices and then compare the calculated ratio for each of the four to that of the unknown. The two closest ratios of the same substance was the identity of the anhydrate.  

Known Formulas of Hydrates:

1.) Copper (II) sulfate pentahydrate = (CuSO4·5H2O)

                Ratio (water to anhydrate): 5 to 1

2.) Magnesium sulfate heptahydrate = (MgSO4·7H2O)

                Ratio (water to anhydrate): 7 to 1

3.) Iron (III) chloride tetrahydrate= (FeCl3·4H2O)

                Ratio (water to anhydrate): 4 to 1

4.) Iron (III) nitrate nonahydrate= (Fe(NO3)3·9H2O)

                Ratio (water to anhydrate): 9 to 1

Experimentally Calculated Hydrates:

If the anhydrate of the hydrate were...

1.) Copper (II) sulfate

- Determining the Number of moles

molar mass of CuSO4 = 159.61g/mol

1.48g CuSO4 x 1 mol CuSO4 / 159.61g mol-1 CuSO4 = 0.009273 mol CuSO4

molar mass of H2O = 18.02g/mol

1.47g H2O x 1 mol H2O / 18.02g mol-1 H2O = 0.08158 mol H2O

- Ratio of Water to Anhydrate

number of moles H2O / number of moles CuSO4

= 0.08158 mol / 0.009273 mol = 8.80 mol H2O / 1 mol CuSO4     (3 significant figures)

        Ratio: 8.80 to 1

2.) Magnesium sulfate

- Determining the Number of Moles

molar mass of MgSO4 = 120.36g/mol

1.48g MgSO4 x 1 mol MgSO4 / 120.36g mol-1 MgSO4 = 0.01230 mol MgSO4

molar mass of H2O = 18.02g/mol

1.47g H2O x 1 mol H2O / 18.02g mol-1 H2O = 0.08158 mol H2O

- Ratio of Water to Anhydrate

number of moles H2O / number of moles MgSO4

= 0.08158 mol / 0.01230 mol = 6.63 mol H2O / 1 mol MgSO4 

        Ratio: 6.63 to 1

3.) Iron (III) chloride

- Determining the Number of Moles

molar mass of FeCl3 = 162.20g/mol

1.48g FeCl3 x 1 mol FeCl3 / 162.20g mol-1 FeCl3 = 0.009125 mol FeCl3

molar mass of H2O = 18.02g/mol

1.47g H2O x 1 mol H2O / 18.02g mol-1 H2O = 0.08158 mol H2O

- Ratio of Water to Anhydrate

number of moles H2O / number of moles FeCl3

= 0.08158 mol / 0.009125 mol = 8.94 mol H2O / 1 mol FeCl3

        Ratio: 8.94 to 1

4.) Iron (III) nitrate

- Determining the Number of Moles

molar mass of Fe(NO3)3 = 241.86g/mol

1.48g Fe(NO3)3 x 1 mol Fe(NO3)3 / 241.86g mol-1 Fe(NO3)3 = 0.006120 mol Fe(NO3)3

molar mass of H2O = 18.02g/mol

1.47g H2O x 1 mol H2O / 18.02g mol-1 H2O = 0.08158 mol H2O

- Ratio of Water to Anhydrate

number of moles H2O / number of moles Fe(NO3)3

= 0.08158 mol / 0.006120 mol = 13.3 mol H2O / 1 mol Fe(NO3)3

        Ratio: 13.3 to 1

Unknown Anhydrate: Magnesium sulfate

The identity of the mysterious substance was magnesium sulfate. After comparing experimentally acquired ratios to the factual ratios for each substance, we determined that the ratios of magnesium sulfate was the closest one out of all four. Its experimental ratio was 6.63 to 1 and its expected ratio was 7:1. The ratios of other three substances were incongruous to each other. For example, the ratio we got from an experiment for iron (III) nitrate was 13.3:1 while it should have been 9:1, according to the information from the resource.

Identity of the Hydrate: MgSO47H2O Magnesium heptahydrate

Error Analysis:

1.) Percent Error

% Error = | (actual value - experimental value) / actual value | x 100%

          = | (6.63 - 7.00) / (6.63) | x 100% = 5.58% Error

The error being only 5.58%, the overall ratio of water to magnesium sulfate was somewhat accurate. However, there must be a few sources of errors that affected the data.

2.) Sources of error

Some sources of deviation of the data may include:

        a. An insufficient amount of time for waiting until all water of the hydrate evaporated. We could have not gotten rid of the water in the hydrate to begin with as 15 minutes of heating was perhaps too short. If the heating continued on for longer, more water could have evaporated to the air, leaving less amount of anhydrate left in the beaker. Less moles of magnesium sulfate in the beaker would have then increased the ratio as the number of water moles would have been divided by a smaller value.  

        b. A loss in the amount of hydrate due to some popping out of the beaker while heating. This phenomenon could have deviated the ratio by causing a loss in the amount of water and anhydrate. According to a smaller ratio compared to the expected ratio, more water was probably lost during this occurrence, which lowered the number of water moles. Because the number of moles of water was lower than what it could have been originally, the ratio of water to anhydrate was 6:63:1 rather than 7:1.

        c. Change in the strength of the heat while maintaining the same amount of time to heat. As we altered the strength of the flame from low to high without increasing the amount of time to wait until all the water can evaporate, there could have possibly been some water left in the beaker with magnesium sulfate that did not evaporate to completion. If we had either heated the beaker with a strong flame from the beginning or increased the amount of time of heating, the number of moles of water during calculation could have been larger. Then the larger number of moles of water divided by the smaller number of moles of anhydrate could have produced a higher ratio that is closer to 7:1 than what we got.

Conclusion

        From this lab, we are able to conclude that our prediction was strongly supported in both terms. First, the assumption that the hydrate is associated with magnesium sulfate due to its white appearance is proven to be correct. Then, the experimental ratio of water to magnesium sulfate being 6.63 to 1 with about 6% error strongly supports our hypothesis to a deeper level. As 6.63:1 is relatively close to 7:1, the expected ratio for this substance, we can thus conclude that the unknown hydrate is magnesium sulfate heptahydrate, MgSO4 7H2O. By using both quantitative and qualitative approaches, we can successfully predict the identity of the hydrate and its structure consisting of anhydrate and water.  

Discussion of Theory

3 new ideas introduced in this experiment...

1.) In this lab, we learned how to apply stoichiometry in a new way to determine a formula of a hydrate. Unfamiliar with hydrates, we were first oblivious to how one could experimentally come up with a correct formula. But as soon as we used previous knowledge of stoichiometry by using molar masses and numbers of moles, we were easily capable of depicting a reasonable empirical formula for the hydrate. Thus, at the end, we learned that there are countless numbers of applications of stoichiometry in chemistry.

2.) Furthermore, this lab illustrated a new term for the group - hydrate. Before this, we had heard of this scientific word briefly in textbooks and in class, but we were never sure of its exact definition. However, as we dehydrated the hydrate and discovered that a hydrate is made of some anhydrate and water with a certain ratio, we soon realized what a hydrate actually was. The exact definition of a hydrate - any substance that contains some amount of water molecules in its structures - was illustrated in a precise way in this experiment.

3.) The last idea we learned was how to apply the knowledge of colors of specific ions and solids. By knowing that ions such as Cu2+ and Fe3+ have their designated colors, we were able to eliminate three options for the anhydrate, FeCl3, Fe(No3)3, and CuSO4, as the hydrate appeared to be white due to the colorless magnesium.Thus, this knowledge of specific colors of ions led us to confidently conclude that the anhydrate was undoubtedly magnesium sulfate.