The purpose of this lab was to explore hydrates when water is removed from a compound with heat. By heating the magnesium sulfate, the mass of the water can be determined as the difference between the original weight and the weight after being heated. Not only can the mass be determined but also the amount of moles, ratio of magnesium sulfate to water, experimental percent of water, theoretical percent of water, and etc. From the lab it can be concluded that water can be evaporated from the substance. It was discovered during the lab that the substance contained five water molecules. This can be calculated by finding the difference between the hydrated magnesium sulfate and dehydrated magnesium sulfate after being heated. By converting the
mass into moles and creating a ratio, the formula can be determined: MgSO4 + 5H2O. After this, the accepted or theoretical percent can also be determined by determining the molar mass of the water and dividing by the molar mass of the hydrate. In addition, the experimental percent of water in the hydrate can be determined with the mass of the water driven off divided by the original hydrate. Some errors were found in the lab, for example there was a 34.14% error between the experimental and theoretical coefficients of water in the water to hydrate ratio as well as a 23.95% error in the experimental water percentage to theoretical water percentage. This could have been caused by the loss of substance when the hydrate was being boiled and particles had escaped or water from the air getting into the substance. Another factor could be the limited number of the numbers shown on the scale. Factors such as these could have interfered in the process making the product less accurate. The lab’s purpose was the explore and test the properties of a ionic compound to determine the water of hydration. It can be inferred from this lab that ionic compounds often have water molecules attached to the substances.
The purpose for this lab was to use aluminum from a soda can to form a chemical compound known as hydrated potassium aluminum sulfate. In the lab aluminum waste were dissolved in KOH or potassium sulfide to form a complex alum. The solution was then filtered through gravity filtration to remove any solid material. 25 mLs of sulfuric acid was then added while gently boiling the solution resulting in crystals forming after cooling in an ice bath. The product was then collected and filter through vacuum filtration. Lastly, crystals were collected and weighed on a scale.
The isomerization procedure was done in order to create dimethyl fumarate from dimethyl maleate. Dimethyl maleate and dimethyl fumarate are cis and trans isomers, respectively. This procedure was done via a free radical mechanism using bromine. The analysis of carvones reaction was done in order to identify the smell and optical rotation of the carvone samples that were provided. The odor was determined by smelling the compound and the optical rotation was determined using a polarimeter.
Compress the safety bulb, hold it firmly against the end of the pipette. Then release the bulb and allow it to draw the liquid into the pipette.
11.) Subtract the mass of the evaporating dish from the mass of the evaporating dish and it's contents. Multiply that number by 10 to get the solubilty in grams per 100 cm3 of water.
We began this investigation by suiting up in lab aprons and goggles, we then gathered our materials, found a lab station and got to work. We decided to start with the magnesium in hydrochloric acid first, we measured out 198.5 L of HCl and put it in the foam-cup calorimeter and took initial temperature reading. We then selected a piece of magnesium ribbon and found its mass: 0.01g. This piece was placed in the calorimeter and the lid was shut immediately to prevent heat from escaping. We “swirled” the liquid mixture in the calorimeter to ensure a reaction, and waited for a temperature change. After a few moments, the final temperature was recorded and DT determined.
Afterwards, we conducted crystallization to evaporate the liquid in an attempt to detect the presence of a salt. Before stating which of the potential
Moisture is heavy, and thus it can change the results of the experiment, as we only want the weight of magnesium and the magnesium oxide.
Refer to Chemistry Lab # 2 – Investigating Changes. No changes have been made in this experiment. Methods = == ==
The first step that we took to accomplish our goal was to put on our safety goggles and choose a lab station to work at. We received one 400ml beaker, one polyethylene pipet, two test tubes with hole rubber stoppers, two small pieces of magnesium (Mg), one thermometer and a vial of hydrochloric acid (HCl). We took the 400ml beaker and filled it about 2/3 full of water (H20) that was 18 OC. Then we measured our pieces of Mg at 1.5 cm and determined that their mass was 1.36*10-2 g. We filled the pipet 2/3 full of HCl and poured it into one of the test tubes. Then, we covered the HCl with just enough H2O so that no H2O would be displaced when the stopper was inserted. After inserting the stopper, we placed the Mg strip into the hole, inverted the test tube and placed it in the 400ml beaker. HCl is heavier than H2O, so it floated from the tube, into the bottom of the beaker, reacting with the Mg along the way to produce hydrogen gas (H2). We then measured the volume of the H2, cleaned up our equipment and performed the experiment a second time.
In this lab, I determined the amount of heat exchanged in four different chemical reactions only using two different compounds and water. The two compounds used were Magnesium Hydroxide and Citric Acid. Both compounds were in there solid states in powder form. Magnesium Hydroxide was mixed with water and the change in heat was measured using a thermometer. The next reaction combined citric acid and magnesium hydroxide in water. The change in heat was measured as well. For the third reaction citric acid was placed in water to measure the change in heat. In the last reaction, citric acid was combined with water. The heat exchanged was again measured. It is obvious we were studying the calorimetry of each reaction. We used a calorimeter
A precipitation reaction can occur when two ionic compounds react and produce an insoluble solid. A precipitate is the result of this reaction. This experiment demonstrates how different compounds, react with each other; specifically relating to the solubility of the compounds involved. The independent variable, will be the changing of the various chemical solutions that were mixed in order to produce different results. Conversely the dependent variable will be the result of the independent variable, these include the precipitates formed, and the changes that can be observed after the experiment has been conducted. The controlled variable will be the measurement of ten droplets per test tube.
In a 100ml beaker 30mls of water was placed the temperature of the water was recorded. 1 teaspoon of Ammonium Nitrate was added to the water and stirred until dissolved. The temperature was then recorded again. This was to see the difference between the initial temperature and the final temperature.
In this experiment, the water of crystallization is removed from hydrated copper(II) sulfate. The mass of water is found by weighing before and after heating. This information is used to find x in the formula: CuSO4.xH2O. Note that x must be an integer (a whole number).
Conclusion This experiment was set out to find the effect of different temperatures of hydrochloric acid on the rate of reaction with magnesium. The information recorded was then interpreted and compared to the hypothesis. From this information, a conclusion can be made to show that the rate of reaction relates to temperature in the reaction between hydrochloric acid and magnesium. In conclusion, as proven in this experiment, the higher the temperature of hydrochloric acid, the faster the reaction it has with magnesium.
This experiment conducted used the Gravimetric Analysis method to determine the concentration of an unknown sulfate solution. The unknown sulfate solution was pipetted into a smaller beaker, acidified, heated, decanted through filtration, dried, cooled and weighed. What was found at the end of the result was Barium Sulfate precipitate which weighed 0.1783 grams. The concentration of the unknown sulfate solution was found to be 2.9385 g/L and its molarity was 0.0306M.