Experiment 5: ColloidsAlmario, Ayana, Monge, June FrancisChemistry 14.1, HEG2 Jenica Marie Lorica MadridejosMarch , 2014

ABSTRACT

A hydrate is an ionic compound that contains water molecules in its structure. To determine the formula of a hydrate experimentally, calculate the mole-to-mole ratio of the water portion compared to the anhydrate portion. An anhydrate is the substance that remains after the water is removed from a hydrate. This experiment aims to determine the formula of the hydrate of copper sulfate crystals (CuSO4) by means of heating and extracting the water present in the compound.

KEYWORDS: INTRODUCTION

Hydrate is a term used in inorganic chemistry and organic chemistry to indicate that a substance contains water. Hydrates have a variety of practical applications. Their ability to gain or lose their waters of hydration makes them versatile. One formula unit of a hydrate contains one formula unit of an anhydride bonded to a fixed number of water molecules. Careful heating removes the water so that the ratio of water molecules to anhydride formula units can be determined, provided the molar mass of the anhydride is known. A hydrate is represented by the formula of the anhydride followed by a raised dot which represents the "weak" bond between the anhydride and the number of water molecules, i.e. CuSO4•5H2O. The stability of hydrates is generally determined by the nature of the compounds, their temperature, and the relative humidity (if they are exposed to air).

EXPERIMENTAL

A clean and dry test tube and cover with a cork stopper were secured and placed in a 100-mL beaker. The 100-mL beaker and stoppered test tube were weighed using an analytical balance. 1 g of copper sulfate crystals was placed in the test tube and was stoppered. The 1 g of copper sulfate crystals together with the stoppered test tube and 100-mL was weighed, if the increase in weight is less than one gram, add more copper sulfate crystals to the test tube and weigh again. The copper sulfate crystals were heated over a blue Bunsen flame until all of the crystals have decomposed or in other words the disintegration of the crystals into a gray powder. The heated copper sulfate crystals was cool to room temperature then was weighed together with the 100-mL beaker.

RESULTS

Formation of moisture occurred while heating the copper sulfate crystals. Gray powder was obtained in the experiment as a result of heating copper

sulfatepentahydrate. Table below shows the results acquired from the experiment.

Table 1. Experimental Data

Weight of beaker + Test tube + Stopper + Crystals before heating

65.99 g

Weight of beaker + Test tube + Stopper 64.99 g

Weight of beaker + Test tube + Stopper + After heating

65.62 g

Weight of residue (CuSO4) 0.63 g

Loss in weight upon heating (H2O ¿ 0.37 g

Formula weight of CuSO4 159.602 g/mol

Formula weight of H 2O 18.0148 g/mol

Moles of CuSO4 0.00395 mol

Moles of H 2O 0.02050 mol

Simplest ratio of moles of H 2O 1:5

Formula of Hydrate CuSO4. 5 H2O

Calculations

Weight of CuSO4= 65.99g– 64.99 = 1 g

Weight of H 2O = W of CuSO4 – W of residue

= 1g - 0.63g

= 0.37g

mole ofCuSO4=0.63gCuSO4 x1mol

159.602gCuSO4¿

=0.00395mol¿

mole of H 2O=0.37 g H2O x1mol

18.0148g H 2O =

0.02050 mol

Experiment 5: Colloids Page 1Chem 14.1

Ratio ofCuSO4∧H2O=0.02050mol H 2O0.00395molCuSO4

=5.18

DISCUSSION

Hydrate compounds are associated with certain number of molecules. Determining the formula of a hydrate through experiment involves heating the hydrate compound to remove the water molecules present. Heating the hydrate compound would retain a residue called anhydrate. Anhydrous compounds can be used as water absorbers.

Copper sulfate crystals are bright bluish ionic compound with rough texture and are also salts. It is an example of hydrate which contains five molecules of H 2O, which can be termed as copper sulfate pentahydrate. Copper sulfate is most often used as fungicide and insecticides and therefore is poisonous.

In the experiment, a beaker, test tube and a stopper were weighed. After weighing the apparatuses, 1g of copper sulfate crystals was added in the test tube and was covered with stopper and weighed. Next, the test tube containing the copper sulfate crystals was heated while moving the test tube up and down to obtain fine results. The copper sulfate will then change its color to darker shade of blue and moisture along the side of test tube could be seen. Continue heating the crystals until the compound turned gray to be sure that all the water was removed. After the process, cover the test tube with the stopper used in the previous step and let it cool down. This is to avoid contamination of the substance heated that might cause an unwanted error like gaining or losing weight. Then proceed to weighing the test tube together with the beaker and check results. The weight of the product is lighter since the water molecules were already removed from the compound. Calculate the results by getting the mole of the water removed dividing it by the residue, copper sulfate, after heating. The answer would be the ratio of the copper to hydate.

In determining the formula of hydrate, inconsistencies are unavoidable since there can be many errors while doing the process. The result might not be exact but could still be closer to the theoretical value of the substance. These are some possible errors encountered in the course of the experiment: The anhydrous salt could have been exposed to air prior to measurement, and reabsorbed some moisture, thus disrupting measurements and making the measurement inaccurate, not all of the water may have evaporated during heating, thus making the moles of water smaller, the moles of CuSO4 larger, and will give you a smaller number of water molecules in the hydrate, , some of the hydrous salt may have

spattered, thus removing a portion of the hydrous salt from the beaker or test tube, thus  the mass of water would appear to be larger than it actually is, which would lead to a higher ratio in the hydrate and human error is always in effect, given that the laboratory does not function under ideal conditions. As such, there is always the possibility of inaccuracies with measurement, perception of measurement, inaccuracies of equipment, and other such errors.

CONCLUSION AND RECOMMENDATIONS

The goal of this experiment is to determine the percentage of water by mass in a hydrate, and to calculate the ratio of salt to water in a hydrated salt. To achieve this, a known mass of hydrated salt was heated, removing the water from the hydrated salt thus forming the anhydrous salt. Given that the mass of the hydrated salt is known, it is also given that the mass lost is equivalent to the mass of the water. By comparing the mass lost to the mass of anhydrous salt left behind, we can calculate the percentage of water within the hydrated salt. We can also (via stoichiometric and molar ratios), calculate the ratio of salt to water. 

Given the data presented above, findings showed the experimental percentage of water within the hydrated salt to be: 37%. The accepted percentage of water within hydrated copper sulfate is: 36.1%. As such, the percent error within this measurement is 2.49% 

The experimental stoichiometric ratio between copper sulfate and water was found to be: 1:5.18. The accepted stoichiometric ratio between copper sulfate and water is: 1:5. The percent error within this measurement is 3.6% 

The formula for hydrated copper sulfate is:

Possible improvements that could be made to this experiment in the future could include increasing the sample size, to produce a more average measurement. Another possibility would be to get more accurate equipment, to replicate the experiment within a dehumidified environment, and/or perform the experiment by being careful in handling the hydrous salt to reduce the possibility of the hydrous salt to be spattered, which will change the expected results.

Experiment 5: Colloids Page 2Chem 14.1

REFERENCES

Petrucci, R. H. General Chemistry. Pearson Educational Inc.

Determining the Formula of a Hydrate, Retrieved January 18, 2014 from: http://www.csun.edu/~jte35633/worksheets/Chemistry/11-6Hydrates.pdf

Copper(II) sulfate, Retrieve January 18, 2014 from: http://en.wikipedia.org/wiki/Copper(II)_sulfate

I hereby certify that I have given substantial contribution to this report.

___________________________Ayana Almario

___________________________June Francis Monge

GUIDE QUESTIONS

1. A piece of iron weighing 0.920 grams was heated until all iron was converted to an oxide. If the weight of the oxide obtained is 1.315 grams, what is the formula of the oxide?

W of Fe = 0.920gW of oxide = 1.315gW of O reacted = 0.395g

Fe = 0.920 g/55.85 g/mol = 0.01647 molesO = 0.395 g/16 g/mol = 0.02469 moles

Weight of O Reacted = W of O – W of Fe

= 1.315-0.920g

= 0.395

mole ofFe=0.920gCuSO4 x1mol

55.85 gFe¿

=¿¿ 0.01647 moles

mole ofO=0.395gOx 1mol16 gO = 0.02469 moles

Experiment 5: Colloids Page 3Chem 14.1

Ratio ofFe∧O=0.02469molesmolO0.01647mol Fe

=1.5 (multiply by 2 to get a whole number)

Formula: Fe2O3

2. A 1.000-gram sample of that contains 0.0128 moles of a hydrocarbon was burned in excess air to convert all the components to carbon dioxide and water. if 3.385 grams of carbon dioxide and 0.692 grams of water are produced, what is the empirical formula of the hydrocarbon? What is its molecular formula?

Calculate the mass of C in 3.385g CO2Molar mass CO2 = 44g/molMolar mass C = 12g/molMass of C = 12/44*3.385 = 0.923g

Mass H2 in 0.692g H2OMolar mass H2O = 18g/molMolar mass H2 = 2g/molMass of H = 2/18*0.692 = 0.077g H

C = 0.923/12 = 0.0769H = 0.077/1 = 0.077Empirical formula = CH

Molar mass CH = 13g/mol

The compound: 1.0g = 0.0128mol: 1mol = 1/0.0128 = 78.125g/mol

78.125/13 = 6.00

Molecular formula = (CH)6 = C6H6 = benzene.

Experiment 5: Colloids Page 4Chem 14.1