Hydrate Lab Report

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%.

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 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

Background: According to George B. Kauffman of Britannica, a hydrate is “any compound containing water in the form of H2O molecules, usually, but not always, with a definite content of water by weight.” Within this, there are two aspects of the compound; the “water of hydration” and the anhydrous salt. When a hydrate is heated, the “water of hydration” leaves the compound as vapor due to the unsteady bonds between the water and salt molecules. The anhydrous salt ends up becoming the only piece that is leftover. From the two, the percent water in a hydrate and the mass of the salt can be determined from the loss of the “water of hydration.” As a result, the formula of the compound can be found because the masses of the salt and the water can be used to determine the mole ratio, which are both small, whole numbers.

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.

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

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

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 doing this experiment is to be able to determine the formula of a copper sulfate hydrate heating it up

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 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

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.

Introduction A hydrate is a chemical compound, where water molecules are chemically bonded to another compound or element, it has a specific number of water molecules, a compound that contains water as part of its crystal. Hydrates are are characterized by the number of moles of water molecules per mole of salt. When there is a chemical formula (CuSO4 . 5H2O) this means for every one formula unit (CuSO4) there is a certain number of water molecules (in this case 5 water molecules).

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.

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.