Percent Of Water In Hydrate Lab Report

The purpose of the experiment was to determine the molar mass of three unknown solutes by measuring the freezing point depression of a solution containing each of the unknowns. When a solute is dissolved in a solvent, the vapor pressure of the solvent is lowered. As a result of the decrease in the vapor pressure, the colligative properties change.

The purpose of this experiment was to determine the percent by mass in a hydrated salt, as well as to learn to handle laboratory apparatus without touching it. The hydrated salt, calcium carbonate, was heated with high temperature to release water molecules. Gravimetric analysis was used in this experiment to determine the percent by mass of water in a hydrated salt. The hypothesis of this experiment was accepted on the basis that the percent by mass of volatile water in the hydrated salt would be fewer than 30%. The percent by mass was determined by the mass of water loss devised by the mass of hydrated salt multiplied by total capacity

The percent of water can be determined in a hydrate by first determining the mass of the hydrate Copper (II) Sulfate penta-hydrate. The substance will be a deep blue color when it is a hydrate. By heating the substance, water is evaporated, removing the water from the hydrate, making it anyhydrous through a simple decomposition reaction. Evaporation is completed when the substance turns from a blue to a white/ grey color. The mass of the water in a hydrate is determined by subtracting the mass of the hydrate from the mass of the anhydrate. The mass of the water is then divided by the mass of the hydrate, and multiplied by one hundred, resulting in the percent of water in the hydrate, which is 36.35%. The percent error is determined by subtracting

The purpose of this experiment was to find the amount and percent of water in a hydrated salt. Hydrated salt contains water molecules chemically bound to it while anhydrous salt is a substance that fire can readily remove the water molecules. Using the gravimetric analysis method, the sample salt was measured, heated, and was measured again after cooled down to room temperature; this procedure allowed the weight of water lost in the salt to be obtained. After two trials, the water in the salt was determined to have an average percent water of -35.04%. The null hypothesis, the heat will not affect the amount of water lost, was rejected ,however, the alternative hypothesis, the heat will allow for there to be water lost, was not supported.

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 powdered cobalt oxalate hydrate was weighed to about 0.3 g and placed in a pre-weighed crucible. The crucible and the cobalt oxalate were then heated until the cobalt oxalate decomposed into a stable, black solid, or Co3O4. Once the crucible was sufficiently

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

One error for the experiment could be from the popping and spritzing of the hydrate when heating it, which could have caused a slight decrease in mass of the whole compound. Since some of the hydrate spritz out of the crucible, the mass decreased. This directly affected the results of the experiment because the mass that spritzed out of the crucible caused the mass to be weighted differently when remeasured on the scale. Another error of the experiment was that there was some hydrate that stuck to the stirring rod when stirring the hydrate.

Subtract the mass of the dehydrated compound from the mass of the hydrate to find out the mass of the water that was in the hydrate. (20.037g)

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.

4H2O A 140.5g sample of NiSO4 . xH2O is heated until no further decrease in mass. The mass of the anhydrous salt is 77.5g. What is the formula of the hydrate? Calculate the percentage of

The purpose of the experiment is to determine how many moles of water or in each of the anhydrous compound, which will in the end let us determine the formula for the hydrate.

The hydrate prevention and control strategies may be implemented in isolation or in combination depending on the specific circumstance of the field, especially as it relates to practicability and economics [10] [A6]. The available different methods of hydrate inhibition are described briefly in this section.

Hydrates are inorganic salts that lose water when heated and undergo a color change. The loss of water makes the hydrate anhydrous. Copper chloride hydrate is what will be used in the lab to show what happens when a hydrate is heated which can lead to a better understanding of hydrates as a whole and their reactions that can be observed to discover things about the environment they’re in. After the removal of water from the copper chloride hydrate, the solid ,when reacted with aluminum and filtered, produces elemental copper. Through the understanding of the process used to achieve this, the experiment can be recreated for another hydrate to isolate a different element. By determining the amount of water and copper in the

    Three measures of unknown sulfate salt was measured and placed into three beakers with water, Hydrochloric acid, and Barium Chloride. The solutions were heated until Barium Sulfate precipitated and the solutions were filtered through ashless filter papers to collect the precipitate. The filter papers were combusted until only the BaSO4 remained, and the precipitates were massed in crucibles. Using the masses of the BaSO4, the percent composition of the sulfate was calculated and used to determine the percent composition of the sulfate in the original unknown sulfate salt. The experiment performed was required for gravimetric analysis of sulfate to find the mass percent of sulfate in the original unknown sample. Without gravimetric analysis, it would be much more difficult to determine the mass composition of sulfate ions in the unknown sample.  The ending result showed that the percent composition of sulfate in the unknown salt was approximately between 23.65% and 29.55%.