chem lab

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Introduction:

In the following experiment, the characteristics of hydrates were analyzed through a series of tests that utilized the crystalline structure of a hydrate. The purpose of these tests was to find the various water compositions of the hydrates, and evaluate the process of dehydration within these substances, which requires the investigation of starting mass and mass lost. To test these variables, the different hydrates were exposed to a source of heat that dehydrated the samples overtime, which were then weighed to calculate how the mass had changed. These characteristics of hydrates are seen throughout everyday life, in products such as Borax (used for cleaning and cosmetic purposes), Epsom salt (a common healing agent), and Gypsum (used for the manufacturing of cement). These products would not remain effective if the water was removed from the crystalline structure of the hydrates and would then become classified as anhydrates.

Procedure and Observations:

To begin the testing of the hydrates, a pea-sized sample of each of the 4 substances were placed into separate test tubes and set into the test tube rack. The physical appearance of each substance in its initial form was observed and documented in the 1st row of Table 1. The copper sulfate was observed as small blue shards with white mixed throughout; ferrous sulfate was a pastel green color with a sand-like consistency; cobalt sulfate was a ruby color the size of tiny pebbles, and the sugar appeared to be a fine grain consistency, the color of white. Using a striker, the Bunsen burner was lit, and each test tube was held over the flame individually for approximately one minute, held at an angle using a test tube holder. During the heating of each sample, observations were gathered based on the changing appearance of the substance and were recorded in the 2nd row of Table 1. The copper sulfate turned to a dark shade initially, but then slowly changed to a white color that started from the outside edge and moved inward. Ferrous sulfate started to turn white during the first stage of heating, then gradually turned to a mossy green color. The cobalt sulfate had the quickest change and turned to a royal blue color with a sizzling noise. Lastly, the sugar turned to a dark brown color after bubbling and fizzing in the test tube. After the heating of one sample, the test tube was placed into the test tube rack to cool for 2-3 minutes and was done for the following 3 samples.

Once the samples were given an adequate amount of time to cool, the appearance of the dehydrated salts and sugar was recorded in the 3rd row of Table 1. The consistency of the copper sulfate had turned to a dry and chalky substance, with a light gray pigment. Similarly, the cobalt sulfate was chalky and had a ring of light blue, royal blue and specks of purple. Ferrous sulfate had turned to medium sized granules that were a mixture of a golden brown and mossy green color. The sugar sample had completely solidified and was a dark brown, almost black color. Lastly, 2-3 drops of water were added to each test tube and the physical changes following this addition was recorded in the 4th row of Table 1. Following the addition of water, the copper, cobalt and ferrous sulfate samples regained some of their initial characteristics. The copper turned to the original blue color, the ferrous had some of the pale green pigment return and the cobalt turned to a dark purple, transitioning to its original ruby pigment. The sugar, however, did not change when water was added to the dehydrated sample, and instead the water rested on top of the sugar layer.

Part 2 of the data collection began with massing the empty crucible, which ranged from 15.7802g to 17.535g, and this value was recorded in data table 2. The scale was zeroed to discard the weight of the crucible, and 1-3 grams of copper sulfate was added to the crucible. The calculated weight of copper sulfate in grams was recorded in data table 2 and ranged from 0.510g to 2.041g throughout the five trials. The crucible was then taken off the scale so it could be zeroed, and the total mass of the crucible and sample was weighed and reported in Table 2. This value ranged from 17.403g to 19.576g, varying in different trials. An iron ring was attached to the ring stand and the wire triangle was placed on the ring, which then held the crucible on top of it. The Bunsen burner was placed below the crucible and was lit using a striker. The flame was low enough to not overheat the crucible or its contents, but it did heat the salt enough to change the color and appearance. These observations are documented in Table 3, explaining how the original light blue eventually turned to a like grey/khaki color at the end of the heating period. Once the appearance of the salt had stopped changing, the sample was heated for another 5 minutes. After this time, the crucible was removed from heat and allowed to cool. The final step to this process was to weigh the crucible and salt sample together and report this value in table 2. The values for this measurement ranged from 16.755g to 18.939g in the five different trials.

The third and final part of data collection for this experiment started by measuring a pea-sized amount of Epsom salt into a test tube and recording the initial appearance of the salt in data table 4. The initial characteristics included clear/white shards that were long and spiked but varied in size. Performing the same procedures that were done in part 1 of data collection, the Epsom salt sample in the test tube was held over the Bunsen burner flame for approximately 1 minute. During this time, any noticeable observations were recorded in data table 4. There were no significant changes during the heating period but there were signs of water vapor building up on the sides of the test tube. After letting the salt cool for several minutes, the final observations of the Epsom salt were recorded in table 4. The Epsom salt remained the same color throughout the experiment and no major changes ever occurred, unlike the other 4 substances in Part 1.

Data:

Copper Sulfate

Ferrous Sulfate

Cobalt Sulfate

Sugar

Initial

Small light blue shards with some white pigment mixed throughout.

Round shaped, tiny pebbles the color of pastel green.

Pebbles the size of table salt; ruby pigment with darker shades mixed throughout.

White small, fine grains; whiter in clumps but clear individually.

During

Darker shade of blue appears from inside then moves to the outside edge. Outsides gradually turn white and move inward.

Carmel brown to white on the outside edge of test tube. Brown color close to the flame. Overall the sample turns a mossy green color with specks of brown.

Quickly turns a dark royal blue color. There is sizzling and signs of liquid. Some of the salt has stuck to the side of the test tube.

Brown bubbling near the flame. Turns a brown/red color. Brown rises from the bottom of test tube to the top.

After

Dry and chalky consistency. Light grey color with no blue remaining. All the sample is stuck together in a clump.

Golden brown closest to flame. Inside ring remains white with the rest of the sample staying a mossy green.

Chalky with light blue center and royal blue edges. There are specks of purple on sides of tube.

The sample has solidified and is a dark brown/black color. The color on sides is a yellow pigment.

H2O

Addition

Original blue color has returned where the water droplets hit the sample.

More of the original pale green returns to the sample.

The sample turns a dark purple color and sizzles when water hits.

The water sits on the surface of substance and the sample does not change.

Table 1: Characteristics of Salt and Sugar Samples. This data table contains the qualitative observation from Part 1 of the data collection, before, during and after being exposed to heat, and after the addition of water.

Trial 1

Trial 2

Trial 3

Group Trial 4

Group Trial 5

Before Heating

Mass of Crucible and sample (g)

17.695

17.403

18.482

18.027

19.576

Mass of Crucible (g)

15.780

15.780

15.780

17.517

17.535

Mass of Sample (g)

1.9150

1.6230

2.7020

0.51000

2.0410

After Heating

Mass of Crucible and sample (g)

17.0120

16.755

17.453

17.858

18.939

Mass of Crucible (g)

15.780

15.780

15.780

17.517

17.535

Mass of Sample (g)

1.2380

0.97500

1.6730

0.3410

1.4040

Mass of water lost by dehydration (g)

0.68300

0.64800

1.0300

0.1690

0.6370

Table 2: Masses Before and After Heating Copper Sulfate. The following data table contains the quantitative measurements from Part 2 of data collection.

Time: (minute)

Observations:

1

Slightly darker color; white circle appearing around edge; white circle begins to expand inward

2

Circle forming on outside of white circle, new color is khaki; center is becoming whiter; slowly all is becoming white

3

Top Layer is white, bottom layer is blue; blue is becoming very pale; small circle of blue remaining; eventually all blue is gone, and the entire sample has turned a light grey/white color.

Table 3: Copper Sulfate Observations. This data table contains the qualitative observations of the copper sulfate while being heated.

Time:

Observations:

Initial

White shards varying in size; long and spikey form; clear color, but white in abundance.

Heating

No significant changes; moisture built up on sides of test tube; combined to make one clump.

After

Salt remains mostly the same; no significant change from initial; top of substance looks more crystalized.

Table 4: Epsom Salt Heating Observations. This data table contains the qualitative observations from Part 3 of data collection, when the Epsom salt is being heated.

Data Analysis and Calculations:

Graph 1: Mass of water lost vs. mass of sample. The following graph shows the relationship between the mass of the sample in data table 2 and the mass of water lost. This relationship can be expressed using the equation: (y = 0.371x – 0.0204).

Equation for Mass % of water:

((mass of water lost from dehydration) / (mass of sample)) * 100 = Mass % of water

(0.6830g / 1.915g) = 0.3560 * 100 = 35.67%

Trial #

Mass % of water

My Sample 1

35.67%

My Sample 2

39.93%

My Sample 3

38.12%

Group Sample 1

33.14%

Group Sample 2

31.31%

Table 5: Mass of % water calculations. This table contains each of the trials and the corresponding mass % of water.

Calculating molar mass of hydrated copper sulfate:

CuSO4 = 159.0 g/mol

H2O = 18.01 g/mol

5 H2O = 90.05 g/mol

CuSO4 5 H20 = (159.0 g/mol) + (90.05 g/mol) = 249.1 g/mol

Calculating actual % mass of water in Copper Sulfate:

5 H20 / CuSO4 5 H20 = (90. 05 g/mol) / (249.1 g/mol) = 0.3615

0.3615 * 100 = 36.15%

Average experimental value of mass % of water:

(35.67%) + (39.93%) + (38.12%) + (33.14%) + (31.31%) = 178.17%

(178.17%) / 5 = 35.63%

Calculating percent error:

% Error = ((experimental – theoretical) / (theoretical)) *100

% Error = (35.63 – 36.15) / 36.15 = -0.01438

% Error = (-0.01438) * 100 = -1.44%

Error = | % error |

Error = | -1.44% | = 1.44%

The following calculations and graphs reflect the composition of hydrates that were tested in this lab. The graph is a visual representation to demonstrate the influence of the mass of a sample on the overall mass lost when the sample has been dehydrated. The high mass percentage of water in copper sulfate was not expected initially and was surprising when calculated. The expected outcome before the beginning of this lab was a water mass percentage of less than 10% based on the solid form of the salt, however, the theoretical and experimental values were both more than 1/3 of the entire mass of the copper sulfate.

Conclusion:

The purpose of this experiment was to find the various water compositions of hydrates and evaluate the process of dehydration within these substances. To do this, the samples of hydrates were exposed to a heat source and the water molecules were removed from the crystalline structure of the salt and sugar. Through the series of tests that were performed, the concept of dehydration as a reversible or irreversible process was observed, as well as the relationship between starting mass and mass lost during the dehydration process. Although the salts and sugar were all tested in the same conditions, not all responded the same way. This proved that not all the substances tested in the experiment possess the characteristics to be qualified as a true “hydrate”.

In part 1 of data collection, with the reintroduction of water, the only substance that returned to its original color was the copper sulfate. Ferrous and cobalt sulfate both regained some original pigment, and the sugar did not return to its original characteristics at all. According to the definition of a hydrate, the substance must be able to undergo a reversible dehydration process in order to be considered one. Because of this, copper sulfate can be confidently categorized as a hydrate, while sugar is most definitely not a hydrate based on the results of the lab. Unlike the 4 other substances, Epsom salt did not change in appearance when exposed to heat. Based on this observation, Epsom salt does not show a difference in its physical characteristics during the dehydration process, and therefore proves the method of observing color change to be ineffective for some hydrates. A better way to observe the release of water molecules would be to use the linear relationship between starting mass and mass lost to calculate the water mass percentage of the Epsom salt.

Further investigating the composition of copper sulfate, part 2 of data collection compared the mass of one sample to the mass of water lost. According to the law of conservation of mass, within a closed system the mass must remain constant, which means the mass lost within the substance was released into the air in a gaseous state. In graph 1, this is illustrated and shows a linear relationship with a positive slope demonstrating an increase in (y) when there is an increase in (x). Using these two variables, the percentage of mass was calculated and fluctuated within an 8.62% range depending on the trial, but when averaged calculated to be 35.63%. This value compared to the theoretical value of 36.15%, calculated a percent error of -1.44% or error of 1.44%. This low percentage means that the overall results of the experiment were accurate, and any sources of error were minor. Possible sources of error within the tests could be attributed to an inaccurate measurement reading from the scales, or an overheating/underheating of the samples. If a sample was underheated, this would result in a larger value for the mass of sample after heating because the dehydration process would be incomplete and would still contain H20 molecules.

In our tests to explore the simple characteristics of hydrates, it can be concluded that the reversible process of dehydration is an effective way to identify a hydrate. The changes that are observed are both quantitative and qualitative, which help better understand the ability of a hydrate to release the water molecules from their chemical structure when exposed to a heat source and reintroduce these water molecules when in contact with H20. Although these series of tests worked well with the samples that were tested, it became clear during the results of the Epsom salt test that not all hydrates change in appearance during the dehydration process. To further understand the properties of hydrates, a following experiment must be performed to understand how water molecules are arranged within varying crystalline structures, and to what extent this affects the physical changes during dehydration.

References:

Gillespie, Claire. “Hydrous Vs. Anhydrous.” Sciencing, 2 Mar. 2019, sciencing.com/hydrous-vs-anhydrous-5554365.html.

Mass of Water Lost vs. Mass of Sample

mass of water lost

1.915 1.623 2.702 0.51 2.0409999999999999 0.68300000000000005 0.64800000000000002 1.0229999999999999 0.16900000000000001 0.63700000000000001

Mass of Sample

Mass of Water Lost