Hydrate Lab Report

Hydration of Norbornene to Norborneol. Objectives: Used acid catalyzation to synthesize the norborneol product, then purified it with sublimation, performed a capillary tube melting point determination and an IR spectrum to characterize the product, and finally determined the stereoselectivity of the reaction. Overall Reaction: 0 out of 10 Reaction Mechanism:.. Stoichiometry Table - "Stoichi Compound Amount Density (g/mL) Molar Mass mMol Molar Equivalence Norbornene 0.3010 g 1.005 94.15 3.187 0.115 Sulfuric Acid (18 M) 1.00 mL 1.84 98.08 18 N/a because Catalyst Water 0.5 mL 1.00 18.02 27.67 8.68 Limiting Reagent: Norbornene Calculations: MMol of Norbornene: 0.3010 g (1 mol)/( = 8.68 Procedure: Procedure Observations: Add a stir bar into

The probable identification of the unknown crystalline hydrate given in class, which was unknown A, is ZnSO4 • 7H2O (Zinc Sulfate Heptahydrate). After completing all the calculations, the percent water of the Unknown hydrate A was 41.997% based on the information gathered during the experiment. By dividing the mass of water lost by the mass of the hydrate, 41.997% was calculated as the percent water. The percent water for CuSO4 • 5H2O and ZnSO4 • 7H2O was also calculated, the conclusion of which hydrate is the unknown was based on the comparison of all three percentages.

Pre-Lab: A) Hydrate- a compound, typically a crystalline one, in which water molecules are chemically bound to another compound or an element. B) Water of hydration (crystallization)- water that is chemically combined with a substance to form a hydrate and can be expelled (as by heating) without essentially altering the composition of the substance. C) Dehydration- the loss or removal of water from something.

The purpose of this lab is to determine the formula of an unknown hydrate. To achieve this, we heated a hydrate over a Bunsen burner to drive out the water. As a result, the anhydrate is left and the data is used to calculate the mole ratio between the amount of anhydride and water. Then the mole ratios are used to calculate the hydration number, which was 4.8, but was rounded to 5 in the formula. The accepted formula is 〖CuSO〗_4∙5H_2 O and the percent of error was 4%.

Water was the added to the anhydrous compound in the crucible. Observations through the entire experiment were recorded along with data. The components of the experiment were cleaned up, and calculation followed. The data obtained was used to calculate: the moles of anhydrous salt, the moles of water removed from the sample, the ratio of moles of anhydrous salt versus the moles of water, the percent water in the hydrate, and the empirical formula of the hydrate. These calculated answers were then compared to the published answers.

We were also provided with 6 molar following reagents:  H2SO4  HCl  NH3  NaOH Experimental Methods: We adopted these steps to find out the formula of unknown hydrate: Qualitative analysis: The unknown hydrate (in solution form) was reacted with all given known aqueous solution and reagents separately in separate test tubes and results were noted down in Table No. 1.1. Quantitative Analysis: The number of water molecules were determined by following steps: o Determined the mass of water that has left the compound.

Fulfilling the purpose, after concluding this experiment, the formula of an unknown hydrate was able to be identified. In order to obtain the data needed to develop a response to the purpose, multiple measurements were collected. The mass of an evaporating dish was measured without copper (II) sulfate hydrate (before and after heating), then measured with copper (II) sulfate hydrate before heating and with anhydrous after heating. Initially, the mass of an evaporating dish was recorded, then the dish was heated and the mass was again recorded. Through this additional step of heating, it insured that no excess water remained in the dish which, as a result, would alter the data. Copper (II) sulfate hydrate was added to the dish and the mass was

The goal of this experiment was to determine the empirical formula for a hydrate of magnesium sulfate and water. The technique that was used was measure the mass of the hydrate and then apply heat to evaporate the water. Then determine the mass of water that was in the hydrate and the mass of the remaining magnesium sulfate. The equation for the hydrate is determined by calculating the mole to mole ratio of the water and the anhydrous. The resulting formula will be formated as: MgSO4*_H2O

The purpose of this lab was to explore the characteristics of hydrates. Hydrates are solid ionic compounds that contain water that is chemically bound to the crystal. In doing this lab, the percentage of water contained in various hydrates, if dehydration is a reversible or irreversible change, and the mathematical relationship between starting mass and mass lost. As there is no simple way to predict the amount of water molecules in a hydrated compound, it must be determined by experimenting. This experiment involves heating said hydrates so the water molecules evaporated from the solid compound. In the different

The mass percent of water was determined using the mass of water and dividing it by the total mass of the hydrate and then multiplying that answer by 100%. The number of moles of water in a hydrate was determined by taking the mass of the water released and dividing it by the molar mass of water. The number of moles of water and the number of moles of the hydrate was used to calculate the ratio of moles of water to moles of the sample. This ratio was then used to write the new and balanced equation of the dehydration process. The sample was then rehydrated to the original state and the percent of the hydrate recovered was calculated by using the mass of the rehydrated sample by

In the experiment, we calculated the percentage of water in 1.5 grams of copper (II) sulfate hydrate by heating it over a Bunsen burner, causing the compound to dehydrate. We then could compare the mass of the hydrate before and after the experiment to determine how much water had been vaporized. Originally, the hydrate inside of the crucible had a total mass of 9.60 grams, but after being heated, the total mass dropped to 9.04 grams. With the crucible alone weighing 8.10 grams, that means that 0.56 grams or 37.33% of water that previously composed the hydrate had completely evaporated from the compound.

This purpose was reached. The water percentage in the hydrate was 52.3% and the number of water molecules in the hydrate was 7. This lab showed that a hydrate can have the water of hydration removed from the anhydrous solution by heating the original hydrate. This is important because during chemical reaction, the experimenter must know if an anhydrous solution is hygroscopic or not, to make sure to incorporate the water molecules that would be present during a reaction. For example, the hydrate sodium sulfate is used as a drying agent for removing traces of water from organic solutions, however, if the hydrate gets too hot, it will release water, and do the opposite effect than it was meant to do. The lab also showed that rounding measurements may produce the theoretical answer, but maybe no the true experimental

The occurrence of large natural clathrate hydrate deposits on the oceanic sea floors and the possibility that these gas hydrates could be mined as an energy source of hydrocarbon gas or used to sequester CO2 gas are still attracting considerable interest.

In Rosemary Jolly's class, students performed the hot plate procedure in order to evaporate the water from the unknown hydrate.  The equipment that the students used were two 100-mL beakers, an analytical balance, a hot plate, and a clean glass rod to stir the substance.  Students obtained about 1 or 2 g of the unknown hydrate into one of their 100-mL beakers.  They determined the combined mass of the sample and the beaker.  After doing this, they placed the beaker onto a hot plate that was on a medium setting.  Using the glass rod, students stirred their beakers in order for all of the unknown hydrate to melt until a dry powder appeared.  Once the dry powder appeared in the beakers, the students took their beakers off of the hot plate in order to cool to room temperature.  They placed the beakers on an analytical balance to record the mass of the beaker and residue.  This information was used later in the experiment to find the number of moles of water per formula weight unit of magnesium sulfate.  Students repeated the procedure with their second 100-mL beaker.  Once the second trial was completed, they used their data to determine the average number of moles of water present in the magnesium sulfate

Table 2: Consists of color extract taken from a red cabbage for a natural indicator. The pH reading that was measured by using the pH meter and the result of the pH reading to determine whether the solution was acidic or basic.