Determination Of Water In A Hydrate Lab Report

The lab performed required the use of quantitative and analytical analysis along with limiting reagent analysis. The reaction of Copper (II) Sulfate, CuSO4, mass of 7.0015g with 2.0095g Fe or iron powder produced a solid precipitate of copper while the solution remained the blue color. Through this the appropriate reaction had to be determined out of the two possibilities. Through the use of a vacuum filtration system the mass of Cu was found to be 2.1726g which meant that through limiting reagent analysis Fe was determined to be the limiting reagent and the chemical reaction was determined to be as following:-

For our second procedure our first step was to gather materials of a 2 cm long magnesium strip, a bunsen burner, and crucible tongs. After we gathered materials we turned the gas to the bunsen burner on, carefully lit the burner with our striker, and adjusted the flame. We then used the crucible tongs to pick up the magnesium strip and place it in the flame, keeping it a safe distance away. Lastly, we recorded and watched our reaction.

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

In this task the concentration of an unknown sample of copper sulphate using colorimetry was used to find the concentration. In this investigation copper sulphate was used which is CuSO4.5H20 as a formula. To make a standard solution which was 1M, the same clean equipment was used to make up the standard solution as used to make sodium carbonate. However there was one difference and that was that the hot distilled water was used to dissolve the copper sulphate crystals. There had to be enough hot water in order to dissolve the crystals into the beaker and then add cold distilled water to cool the solution.

The Cu Later lab experiment is designed to allow you to practice lab skills in implementing and performing a series of reactions. Specifically, four types of chemical reactions will occur: oxidation/reduction; double replacement; single replacement; and decomposition. You will begin with a known amount of copper metal, which, after progressing through several steps, is reproduced. In this experiment you will observe and record the various changes such as heat, color changes, and production that occur. This procedure is used to observe some chemical reactions of copper and its compounds while also performing the lab appropriately as to retain the copper as much as

Put approximately 9-10(g) copper ore into beaker. Use spatula to break up any large pieces. Next add 17ml H2SO4 (aq) (hydro sulfuric acid) to the beaker. Began mixing until all or most traces of blue dissipate; or the copper ore will no longer dissolve (should appear as a milky liquid). Next use pipette to and remove solution and divide solution into 2 individual test tubes then Place test tubes into centrifuge and run centrifuge for 1 minute. Remove from centrifuge machine Fill a cuvette with the clear solution from the test tube making sure not to disturb the sediment at the bottom. Note the solution should bluish in tint Final place the cuvette in the colorimeter. Then record data and calculate in results section.

At the end of the experiment when the lid was removed, it was found out that the blue colour of the copper (II) sulphate solution has faded away. It was turned to pale grey and there were some precipitates present. It was the zinc powder that was in excess to ensure that the copper (II) sulphate solution could react fully with the zinc powder.

The primary of this lab is about the hydrate. The hydrate is a compound formed by an ionic bond combined with the water molecule (known as “water of hydration”) attached to it. So for every formula unit, there would be some amount of water molecule combined with the ionic compound and would act as a single compound. Our theoretical hydrate for this lab was CuSO4 * 5H2O. To remove the water from the compound and get the mass of the ionic compound, you need to follow the dehydrating procedure by heating. That would separate the ionic bond (CuSO4) from the water molecule (5H2O) once it reaches certain decomposition temperature. And to physically determine whether or not the water of hydration is removed, see the color, in our case, turning white from blue. Once the water molecules have entirely evaporated, the mass left would be the weight of the ionic compound (mass of CuSO4 in grams).

To begin Lab 7 of Chem 115, a clean and dry porcelain crucible and its cover were obtained. Next, an iron ring was attached to a ring stand. A clay triangle was placed on top to the ring and a Bunsen burner was placed under the ring. Following the setup for the experiment, the crucible and its cover were placed on the clay triangle and were heated for about five minutes. After, the burner was turned off and the crucible and cover were left to cool to room temperature. Once the crucible and its cover had reached room temperature, tongs were used to move them to a wire gauze. Using the wire guaze, the crucible and its cover were transported to an analytical scale to weigh and record the mass of it. Next, a strip of magnesium was obtained and

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

During the heating period, a noticeable change took place in the sugar test tube – it began to caramelize, melting into a golden brown color, then finally a dark brownish/black hue; it remained constant at that color. The copper sulfate began losing color when applied to the heat, turning from a bright blue to white and remained white after heating. Once cooled, 2-3 drops of water were added to each tube; the sugar remained the same, while the copper sulfate immediately bubbled and transformed from white back to its original blue hue. Next 1-3 grams of the hydrated copper sulfate were added to a crucible, which was weighed prior to adding the sample, then heated gently with a Bunsen burner. Once the salt stopped changing color, it was heated for an additional 5 minutes, then the total mass was measured. Finally, a pea-sized amount of Epsom Salt was added to a test tube and heated for 1 minute without the tube becoming red

A hydrate is a chemical compound containing salt and water. Depending on the salt, there is a maximum amount of water molecules that can be absorbed by the salt. An anhydrous salt is when the hydrate loses water, which happens when the hydrate is heated

This experiment is based on determining the chemical formula for a hydrated compound containing copper, chloride, and water molecules in the crystal structure of the solid compound, using law of definite proportion. The general formula of the compound is CuxCly•zH2O, and aim is to determine chemical formula of this compound.

Purpose of this experiment was to find the amount and percent of water in a hydrated salt. Also, to successfully determine percent error and standard deviation. Hydrated salts are substances that occur naturally who usually contain an amount of water molecules chemically bonded to the compound. A few hydrated salts have weak bonds within the water molecules which allows heat to remove the water molecules creating an anhydrous salt. Hydrated salts that lose water molecules to the atmosphere without a heat source are known as efflorescent. Salts that readily absorb water are called deliquescent. An example of an anhydrous salt would be Magnesium Sulfate or also known as Epsom Salt. Epsom salt separates under heating and becomes

When the zinc was added to the copper (II) sulfate solution, the solution started to bubble. As the solution was stirred, it turned a cloudy blue. Small flecks of a brown solid were visible. As the solution became colorless, the brown solid settled to the bottom of the beaker. The solid formed was copper in its elemental state. The color faded from the solution as the copper ions slowly formed into solid copper. The copper was poured into a funnel with filter paper and washed three times with 25 mL