Analyzing the Identity of a Metal Carbonate using Stoichiometry Introduction: In this lab, I determined the identity of an unknown carbonate. Based off of the mass percentage of this unknown carbonate, I was able to determine whether it was sodium carbonate, potassium carbonate, or rubidium carbonate. The purpose of this lab was to determine the mass percentage of said unknown carbonate in the given precipitate. The method used to determine this was called gravimetric analysis. There are many gravimetric analysis methods, but the one used in this lab, measured mass via precipitation. The ion, analyte, Strontium Carbonate was precipitated in a Büchner funnel apparatus with 150 mL of deionized water and a calculated amount of Strontium Chloride …show more content…
The mass percentage of unknown carbonate was 57.58%, this was ±0.98% more than the calculated percent mass of sodium carbonate. There was more metal carbonate in my sample than predicted in my brain. This calculation means that I discovered sodium carbonate. I didn’t predict so many students would discover potassium carbonate. We all had different amounts of the unknown carbonate between 1.0g ±0.1g-3.0g±0.1g. The slope of the graph is positive except for a few outliers. I know that my data was an outlier. To answer the question, did it matter how much unknown compound was investigated with respect to percent of carbonate in the sample? I believe it did matter how much of the unknown carbonate was investigated. We were instructed to measure about 1.0g or about 3.0g not exact amounts. I feel like exact measured amounts would have reduced the amount of outliers in the graph. The class data changed my confidence in identifying the unknown because twelve out of twenty students discovered the same compound, where as my discovery was an …show more content…
I enjoyed using the Büchner funnel apparatus. If I had more time in lab I would love to use Büchner funnel apparatus on different compounds. It makes me wonder how else can we utilize this piece of equipment. I’d also like to use more acidic chemicals to find different precipitates and practice properly disposing of acidic chemicals again, as executed in a previous lab. In CH 221 we practice stoichiometry. The stoichiometry used in this lab can be executed in the clicker questions, textbook problems, on sapling, or on ALEKS. Stoichiometry helps with conversions. We also practice stoichiometry to find the mass percentage, percent yield, theoretical yield, and mass of an element. All of these conversions can be used in this lab. The work executed in this lab can relate to that of producing washing soda. Washing soda is used in a laboratory setting and used in a household setting. To produce the proper amount for household use a factory needs to be able to synthesize sodium carbonate; measure out sodium carbonate, and correctly mass sodium carbonate. This process can help for the proper amount per single item distribution. Each item has its own said mass of sodium carbonate mixed with another cleaning
The purpose of the lab is to acquire the percent composition of zinc and copper. The procedure included obtaining a post 1983 penny and washing it with soap and water. Using a triangular file, we made an X on the penny. Then, we cleaned the top and bottom of the penny with steel wool until it was shiny. We rinsed the penny in acetone and dried it with paper towel.
In order to begin this experiment, first one must find the balanced chemical equation for the reaction which occurs between the aluminum and copper (II) chloride. This balanced equation being 2Al(s)+3CuCl2 (aq)3Cu(s)+2AlCl3 (aq). After finding this equation, one must use the process of stoichiometry in order to find how many grams of aluminum are needed in order to produce 0.15 grams of copper. In this experiment, the purpose was to produce between 0.1 and 0.2 grams of copper, so one should attempt to produce 0.15 grams of copper seeing as it is the average of those two numbers. The first step in the stoichiometric process which one has to complete is finding how many grams of copper are in one mole of copper.
To begin with, is the experimental process used to determine the identity of the rock. In doing so one will need to discover the density of the rock. By measuring the rock sample with grams per milliliter is a way used to figure out the density. In starting one will need to measure the mass of the rock using grams. Then using a set milliliter amount of a liquid substance, such as water, one will place the rock sample inside.
In our group our data was pretty good we had two mess ups on scale 2 (.06kg, 1N) and (.14kg, 2N). These points were a lot bigger then the points with scale 1 and scale 3, and when I made the graph those two points were farther away and more of set of the line of best fit. Also on scale 2 and 1 I had the same weight for 2 different masses I had 2N for both .14kg and .16kg this also happened on scale 1 and 3 on trail 6 and 5 on scale 1 I had (.14kg, 1.8N) and on scale 3 I had (.16kg, 1.8N). I think these mistakes happened because I read the scale wrong.
Isotopes of the same type will have a much more uniform weight. Despite these sources of error, the experiment was successful in representing the variations of the different isotopes of an
I. Purpose: To experimentally determine the mass and the mole content of a measured sample. II. Materials: The materials used in this experiment a 50-mL beaker, 12 samples, a balance and paper towels. III.
In this lab there were five different stations. For the first station we had to determine an unknown mass and the percent difference. To find the unknown mass we set up the equation Fleft*dleft = Fright*dright. We then substituted in the values (26.05 N * 41cm = 34cm * x N) and solved for Fright to get (320.5g). To determine the percent difference we used the formula Abs[((Value 1 - Value 2) / average of 1 & 2) * 100], substituted the values (Abs[((320.5 - 315.8) /
The initial experiment had results that were slightly different than I had predicted. In my hypothesis, I stated that coacervates would start forming when the pH was 5, but be most abundant when the pH was 4.5. Prior to the experiment, I did not think the coacervates would change in size, instead, I thought that as the solution
3. In this experiment, the percent yield was 90%. This number implies that there was little error in this experiment. However, this result could have been caused by certain external factors.
The actual data is the result on our experiment vs theoretical, which is based on the calculations above. I have also learned to pay more attention to draining out all of the product completely before continuing to test the experiment, as any small drop of contaminant can veer our results into a different
But the difference was no bigger than 0.08, and after the values were rounded the same empirical formula was deduced. So the experiment can be concluded as successful. Evaluation: The method used was simple and easy to follow; however, it did not include how much oxygen was needed to react completely. Also it didn 't mention what magnesium oxide looked like after it finished reacting, so it was a guesswork of determining whether the reaction was finished or not.
Introduction: In this lab, of water in a hydrate, or a substance whose crystalline structure is bound to water molecules by weak bonds, is determined by heating up a small sample of it. By heating, the water of hydration, or bound water, is removed, leaving only what is called an anhydrous compound. Based on the percent water in the hydrate, it can be classified as one of three types: BaCl2O ⋅ 2H20, with a percent water of about 14.57%, CuSO4
Aim: To find out the relationship between the greater concentration of sodium thiosulfate when mixed with hydrochloric acid and the time it takes for the reaction (the time it takes for the solution to turn cloudy) to take place and to show the effect on the rate of reaction when the concentration of one of the reactants change. Introduction: The theory of this experiment is that sodium thiosulfate and hydrochloric acid reach together to produce sulfur as one of its products. Sulfur is a yellow precipitate so, the solution will turn to yellow color while the reaction is occurring and it will continue until it will slowly turn completely opaque. The reaction of the experiment happens with this formula: “Na2 S2 O3 + HCL =
IV. Data and observations Mass of beaker (g) 174.01 Mass of beaker + NaOH pellets (g) 174.54 Mass of NaOH pellets 0.53 TRIAL 1 TRIAL 2 Mass of potassium acid phtalate (KHP) (g) 0.15 0.15 final buret reading (ml) 30.75
Synopsis This experiment is the determination of Calcium Carbonate (CaCO3) content in toothpaste with the use of back titration while demonstrating quantitative transfer of solids and liquids. A accurately weighed quantity of toothpaste was dissolved in excess volumes of HCl. This solution is then titrated with NaOH to find the volume of the excess HCl. The volume of HCl reacted, which is found by substracting the volume of given HCl with the volume of excess HCl reacted, can be further manipulated with mole fractions to find the mass of CaCO3 and thus the CaCO3 content in toothpastes.