Dropper Lab Report

There are four major systematic errors in the laboratory, which are the fact that some of the Alka-Seltzer pills were chipped, some of the water splashed out of the beaker, we didn’t know when the reaction was actually finished, and the mass of the Alka-Seltzer might have been different.

In this experiment, we use a known concentration (Na2CO3) to determine an unknown concentration of acid (HCl). This reaction is neutralisation reaction, which we’ll get water and salt as the products.

The mole is a convenient unit for analyzing chemical reactions. Avogadro’s number is equal to the mole. The mass of a mole of any compound or element is the mass in grams that corresponds to the molecular formula, also known as the atomic mass. In this experiment, you will observe the reaction of iron nails with a solution of copper (II) chloride and determine the number of moles involved in the reaction. You will determine the number of moles of copper produced in the reaction of iron and copper (II) chloride, determine the number of moles of iron used up in the reaction of iron and copper (II) chloride, determine the ratio of moles of iron to moles of copper, and determine the number of atoms and formula units involved in

What are two possible sources of error for this experiment? Discuss how these errors would have changed your results. Specifically, how would the ratio of moles of anhydride to moles of water have changed?

Conclusion – The main idea of this experiment is that if you have the reactants of a double replacement reaction that you should be able to find out several things. First, you can figure out the products from switching the positive ion with the other positive ion. Once you have found the products you can determine the phase of matter they are in from Table F and H. Finally, you can also determine whether the reaction went to completion or not. The observations and data above leads me to believe that our experiment is valid. We were very careful and observant of the directions that needed to be taken. Once source of error could have been that certain substances were placed in the wrong wells, causing your data to stray from the correct information. One way this experiment could be applied to a real life situation is he residue in the bottom of a glass in your bathroom. There is stuff dissolved in water, and any trace of water left in a glass will eventually evaporate and leave the residue.

The results showed the molarity of the NaOH solution. This experiment was completed twice and a new average molarity

In 2 and 7 I added 50 mL of .1 M NaCl. I added sodium acetate to the rest of the beakers: 1 gram to 3 and 8, 5 grams to 4 and 9, and 10 grams to 5 and 10. I then filled the beakers that contained the solid sodium acetate with 50 ml of .10 M acetic acid. Specifics can be found on page 84 of the lab manual. Though the lab manual instructed to use a pipet, we did not have an accurate 1 mL pipet or a graduated pipet, so we instead prepared two graduated burets with 1 M Sodium Hydroxide and 1 M hydrochloric acid. Using a standardized pH probe with a Lab Pro to measure changes in pH, we added 1 mL of HCl at a time and recorded the changes. The same was done for the NaOH.

The organic layer on the top formed, and was also extracted out through vacuum filtration like the step before. The solid that was produced from the reaction was bright yellow in color. It then was set out to dry for a week, and worked on the next experiment time. After setting the solid aside, the original solution that was still in the separatory funnel, 14 mL of 10% NaOH was added. The aqueous layer was then again separated into a separate flask. This same step was repeated twice more. The aqueous solution that formed was then cooled in an ice-bath. Concentrated HCl was then added until a pH reading of 1 was reached making it an acidic solution. This took 13 mL of HCl to reach the adequate pH. Another precipitate formed from this process, and it was extracted through vacuum filtration. The solid precipitate that was formed was a cloudy white color. It also, was set out to dry for a week until the next experiment time.With the remaining solution in the separatory flask, CaCl2 (a drying agent, which helps dry the precipitate that was filtered out so that no extra liquid will be left over 5), was added. This solution was boiled over a steam bath until no bubbles were present, and left to dry

The litmus test might have contributed to a significant loss of the precipitate, but only one instead of multiple tests were performed to see if the reaction mixture was acidic enough. Given the carefully controlled experiments and the 90% yield, it can be deemed that the methods used to perform the experiments are full proof and performed as immaculately as possible.

To determine the concentrations in both parts, the change in pH is monitored for every added amount of 0.1035 M NaOH. The pH of the solution is taken using the Orion Model 420A pH reader. The data for Figures 1-4 can be found in Appendix. B.

Its concentration was found through a quantitative and qualitative titration testing. Because HCl is acidic, a 1.07M NaOH solution was used to titrate an aqueous solution of HCl (the unknown solution). The PASCO program gave the final results shown in Figure 1. A typical titration curve can be described as starting at the initial pH of the solution, then some acid/base is added, until a sharp curve either up or down takes the pH past a neutral of 7, and towards the pH level of whatever the added acid’s/base’s pH is (Clark6). The stronger the acid or base being titrated, the faster the change occurs, which means a sharper slope of the titration graph. Also, it is important to note the equivalence point on this titration graph. At the equivalence point “the correct amount of standard solution must be added to fully react with the unknown concentration” (Xavier3). In Figure 1, this equivalence point is shown by where the sharpest slope occurs in the curve, which is seen at about a pH of 5.0

Background A chemical equation that is in equilibrium will naturally balance itself out should either the products are the reactants be increased, or other stresses such as temperature and pressure are placed on it. However, as with a lever in physics, the fulcrum or balancing point is not always in the middle. In most cases, one side of the equation will be favored more. Ergo, the equation will try to balance to this point of equilibrium.

For lab eight, the molarity of the NaOH solution is 0.07823 M. For lab nine, the molarity of the NaOH solution is 0.4224 M. The molarity of the NaOH solution for both lab eight and nine was not accurate because the solution did not turn the color to pink. Some source of error in this lab is oxalic acid is not completely moved from the plastic lab tray to the beaker. There is a small amount of substance left in the plastic lab tray that cause the mass of the beaker with NaOH solution to have less mass than I measured previously. Another source of error is not all the base was delivered from the burette to the beaker with oxalic acid. Some base remained on the side of burette instead of fully going into the beaker. This leftover base would change

Possible sources of error could be related to how the collected pH of the solutions may not have been

For this experiment, a pH meter was used so this part of the experiment began with the calibration of the pH meter with specified buffers. The buret was then filled with the standard HCl solution and a set-up for titration was prepared. 200g of the carbonate-bicarbonate solid sample was weighed and dissolved in 100 mL of distilled water. The sample solution was then transferred into a 250-ml volumetric flask and was diluted to the 250-mL mark. The flask was inverted several times for uniform mixing. A 50-mL aliquot of the sample solution was measured and placed unto a beaker. 3 drops of the phenolphthalein indicator was added to the solution in the beaker. The electrode of the pH meter was then immersed in the beaker and the solution containing the carbonate-bicarbonate mixture was titrated with the standard HCl solution to the phenolphthalein endpoint. Readings of the pH were taken at an interval of 0.5 mL addition of the titrant. After the first endpoint is obtained, 3 drops of the methyl orange was added to the same solution and was titrated with the standard acid until the formation of an orange-colored solution. Readings of the pH were also taken at 0.5 mL addition of the titrant.