Volumetric Titrimetry Lab

We know that that the end point of the titration is reached when, after drop after careful drop of NaOH, the solution in the flask retains its pale pink color while swirling for about 30

4. To utilize the titration results to calculate the molarity of the hydrochloric acid and the

The purpose of this experiment is to determine an unknown concentration of acid (hydrochloric acid) with a standard solution of a base (sodium carbonate) using titration method.

for percent acetic acid in vinegar, it can be concluded that the percentage of acetic acid in the

2. In Part I of this experiment, acetic acid is titrated with NaOH. The net ionic equation for acetic acid reacting with NaOH is CH3COOH+ NaOH =NaC2H3O2+H2O. The equivalence point is when the moles of the titrant and other solution are equal.. You detect the equivalence point by obtaining the point on the graph where the steep pH occurs. In titrating acetic acid with NaOH, the pH is greater than 7 at the equivalence point because NaOH is a strong base so it results in a higher pH, due to the OH- ions in the solution.

The density had to be determined by massing a known volume of the solution and dividing the mass by the volume. The percent by weight of acetic acid in the solution needed to be found by dividing the grams of solute from the grams of total solution. This required manipulation of the equations for density and molarity to reach the desired value. With this information, it was possible to compare both the molarities and the percent composition by weight of the vinegars to store bought

When filling up the burette it is important that a funnel is used, however as the solution reaches the 0 mark it is ideal that the funnel be removed and a pipette used instead to reach the 0 mark, this is to achieve greater precision. During the experiment, it is important to swirl the flask continuously with one hand

ii. The second part of the titration series involves titration of NaOH with Hydrochloric acid (HCL). Again, three reps of titration and a blank titration have to be completed. A volumetric pipet is used to measure 10.00mL of HCL into three labeled conical flasks. Then the flasks are filled with deionized water until about the 50mL mark. A buret is

The purpose of the lab was met. The percent weight and concentration of acetic acid in the vinegar solution and its Ka, the concentration of phosphoric acid in the Unknown 86 container at each equivalent point, and each Ka value at each equivalent point are all listed below

After the buret was cleaned it was filled up with MnO4- to 5.75 mL. While the buret was being set up, a small sample of ferrous ammonium sulfate was being obtained and weighed. Also 10 mL of both DI water and 3M H2SO4 were obtained and added to an Erlenmeyer flask After these three steps were complete the .441 grams of ferrous ammonium sulfate that was obtained was added to an Erlenmeyer flask and dissolved in the solution of DI water and H2SO4. Once all the ferrous ammonium sulfate was dissolved, the Erlenmeyer flask was placed underneath the buret. The titration then began and 1 mL increments of MnO4- were added to the dissolved ferrous ammonium sulfate. Once 13-15 mL of MnO4- had been added, the solution began to turn pink.

Used a scoopula to take out baking soda and filled three 30 mL test tube with 0.2 g, 0.4 g, 0.6 g, 0.8 g and 1 g using an electronic balance. Put the test tubes in the rack. 4. Filled a tub with water 30 cm and filled water to the 100 mL mark on the graduated cylinder. 5.

Add the sodium hydroxide drop by drop until the end of the point is reached. The sodium hydroxide should be added while maintaining a gentle swirling motion of the flask. The solution should be a very light shade of pink when the titration is finished. Record the final burette reading. Repeat the titration three more

Experiment to investigate the amount of sodium hydroxide needed to neutralize the solution of vinegar

First, three titration curves and three second derivative curves were created to determine the average pH at the half-equivalence point from the acetic acid titrations. Titration curves were used as visuals to portray buffer capacity. The graphs and a table, Table 1, that showcased the values collected were created and included below. The flat region, the middle part, of Figures 1, 2 and 3, showed the zone at which the addition of a base or acid did not cause changes in pH. Once surpassed, the pH increased rapidly when a small amount of base, NaOH, was added to the buffer solution. Using the figures below and

2cm of a solution was tested and added 2 cm of 10% of potassium hydroxide solution and the test tube was shaked.