Determining the Empirical Formula of Magnesium Oxide and the Chemical Formula of a Hydrate

The second experiment deals with hydrate analysis. Hydrates are crystallized salts that have water as part of their structure [1]. Each metal cation in the hydrate has a specific number of water molecules that it bonds with. Therefore, the molecular formula of a hydrate will always be one mole of the compound and an “x” number of moles of water [8]. For example, in Iron (III) chloride hexahydrate crystal, FeCl3 · 6H2O, there is one Iron(III) chloride formula unit per six water molecules. The objective of this experiment is to calculate the mass percent of water in copper (II) sulfate hydrate, CuSO4∙ xH2O. Dehydration refers to the strong heating of the hydrates which causes the water bonds to break apart. The salt product is known as the anhydrate, as in without water [6]. By performing this method, the mass of the salt product and the water can be determined. With this knowledge, the ratio may be calculated and the number of water molecules can be identified.  In order to find the mass percent of water in copper (II) sulfate the mass of water must be divided by the mass of the hydrate and then multiplied by 100%.  By knowing the mass of water and copper (II) sulfate, the percentage of water in copper (II) sulfate can be determined. Knowing this information, it can be predicted that the percentage composition of water in the hydrate is approximately around 60%.

Experiment 1: Determining the Empirical Formula of Magnesium Oxide

Materials and Method for the Experiment  

*The procedure for this experiment can be found on pages 212 and 213 in the McGraw-Hill Ryerson textbook.*

Some changes to the experiment that might change the results are:

Used less amount of Magnesium than stated

Experimental Results:

Table 1:  Qualitative Observations of the Magnesium

Table 2: Data obtained while reacting magnesium in air

 *This charts shows the results of the experiment “Determining the Empirical Formula of Magnesium Oxide”, the data is the weighing of the mass in each stage of the experiment. This data is obtained through, the weighing of the different materials.

Calculations (Mass by difference):

Mass of Magnesium = (Mass of crucible, lid and magnesium) – (Mass of clean, empty crucible and lid)

Mass of Magnesium = 46.29grams – 46.25grams

Mass of Magnesium = 0.04grams

Uncertainty: 

General equation when calculating numbers with uncertainty:

While addition or subtracting (let e represent uncertainty) [5]:

Absolute uncertainty= e1 + e2 + e3…

Uncertainty in mass of magnesium = (Uncertainty in mass of crucible, lid and magnesium) + (Uncertainty in mass of clean, empty crucible and lid)

Uncertainty in mass of magnesium = 0.01 + 0.01

         = 0.02 grams

The masses and uncertainties of oxygen can be determined in a similar manner but by subtracting the mass of magnesium from the mass of magnesium oxide instead.

Table 3: Quantitative observations of the Reactants and Product

To determine the empirical formula for Magnesium Oxide:

Convert all masses into moles

                        Mass of Element *  = Number of moles

0.04± 0.02 grams Mg          ×           = 1.65 × 10^-3 ± 0.02 mol

                        0.16g ± 0.02 grams O         ×                 = 0.01 ± 0.02mol

Determine the lowest whole number ratio of all molar quantities:

                General equation:  

               Mg=        1.65 × 10^-3 ± 0.02 mol   = 0.17

                                                     0.01 ± 0.02mol

                                

                                         O= 0.01 ± 0.02mol = 1

± 0.02mol

* In the case above, ignore the uncertainties as the goal is to find only whole number ratios.

Use the following conversion table to convert all mole ratios to whole number ratios [3]:

Therefore,         Mg= 0.17 * 6 = 1

                        O = 1 * 6 = 6

In conclusion, the empirical formula of Magnesium Oxide is MgO6.

Calculating the percent error:

Percentage Error (PE) =  × 100%

Percentage Error (PE) =   × 100% 

Percentage Error (PE) =  × 100% = -66%

Since the percentage error is negative, take the absolute value of the percentage. Therefore, the percentage error is 66 %.

Experiment 2: Determining the Chemical Formula and the Mass percent of water in a Hydrate

Materials Needed:

400mL beaker

Tongs

Scoopula

Electronic balance

Glass rod

Hot pad

3g to 5g hydrated copper (II) sulfate

*The procedure for this experiment can be found on pages 212 and 213 in the McGraw-hill Ryerson textbook.*

 Some changes to the experiment that might change the results are:

Used less amount of hydrated copper (II) sulfate (CuSO4)

Placed temperature at medium low

Table 4: Qualitative Observations

Table 5: Determining the Chemical Formula of a Hydrate

*This chart shows the numerical results of the experiment “Determining the Chemical Formula and the mass percent of water in a Hydrate”. It provides the mass recorded in each stage of the experiment.

Calculations (Mass by difference):

Mass of Hydrate copper (II) sulfate = (Mass of beaker or evaporating dish + hydrate copper (II) sulfate) – (Mass of empty beaker or evaporating dish)

Mass of Hydrate copper (II) Sulfate = 100.20grams -98.64grams

Mass of Hydrate copper (II) Sulfate = 1.56 grams        

                 The mass of anhydrate copper (II) sulfate and H2O can be determined in a similar manner. *Refer to the uncertainty equation on Page 3 to calculate all uncertainties.

Table 6: Masses of Reactants and Product:

To determine the chemical formula of a Hydrate:

Convert all masses in to moles

Mass of Element *  = Number of moles

0.73 ± 0.02 grams CuSO4 ×  = 4.57 × 10^-3 ± 0.02 mol

0.83 ± 0.02 grams H2O ×  = 0.046 ± 0.02 mol

Determine the lowest whole number ratio of all molar quantities:

Lowest ratio=         

CuSO4= = 1

    H20= = 4.91 = 5

*In this case, disregard uncertainties as it accounts for only a minor difference to the mole ratio.

Therefore, the chemical formula for Hydrate Copper (II) Sulfate is CuSO4 × 5 H2O 

Determining the mass percent of water in a Hydrate of Copper (II) sulfate:

       Molar mass of water (H2O)____   x 100%                               

Molar mass of CuSO4 x 5 H2O

Molar mass of water= 5 * [(2 * Mass of Hydrogen) + Mass of oxygen]

         = 5 * [(2 * 1.01) + 16.00]

         = 90.1 grams

Molar mass of CuSO4 = Mass of Copper + Mass of Sulfur + (4 * Mass of oxygen)

             = 63.55 + 32.06 + (4 * 16.00)

             = 249.71 grams

Therefore,

       Mass percent of water =     90.1 grams_   x 100%    = 36%                           

                                                             249.71 grams

Bibliography

         Anne , M. (2011). Molecular and empirical formula. Retrieved from http://chemistry.about.com/od/molecularformulas/a/Molecular-Formula-And-Empirical-Formula.htm

         Chieh, C. (2008). Chemical formulas: Key concepts. Retrieved from http://www.science.uwaterloo.ca/~cchieh/cact/c120/formula.html

         Clancy, Christina. McGraw-Hill Ryerson chemistry 11. Whitby, Ont.: McGraw-Hill Ryerson, 2001. Print.

         Denker, J. (2010). Measurements. Retrieved from http://www.av8n.com/physics/uncertainty.htm

         Forman, A. (2007). Naming hydrates. Retrieved from http://www.digitalillusions.ca/applewoodscience/LessonsOnLiine/Chemistry/SCH3U0/nomenclature/hydrates.html

         Gwen, S. (2010). Moles. Retrieved from http://www.kentchemistry.com/links/bonding/empirical.htm         

         Kauffman, G. (2011). Hydrates. Retrieved from http://www.britannica.com/EBchecked/topic/278148/hydrate

         Regina, F. (2009). Uncertainties and Significant figures. Retrieved from http://nebula.deanza.edu:16080/~norona/4a_labs_files/uncertaintyandsignificantfig.pdf.

         Taylor, J. (2006). Introduction to error analysis. Retrieved from http://www.wellesley.edu/Chemistry/Chem105manual/Appendices/uncertainty_sigfigs.html

                

[3]

Figure 1: This figure shows the setup for the experiment. This image was taken from Investigation 6-A in the McGraw-Hill Ryerson textbook.  

Figure 2: This image shows the setup for the experiment. The experimental results should look something like this. The substance loses some of its color [3].