MATERIALS AND METHODS:
Before the acetic acid solution could be titrated with sodium hydroxide (NaOH), the actual concentration of NaOH needed to be determined. By way of standardization, the actual concentration of NaOH was established, to account for the fact that the solid is not pure and for its tendency to react with carbon dioxide in the air.
A 50 mL burette (±0.01 mL, Kimax) was rinsed thoroughly, twice with reverse osmosis water, and then twice with approximately 5 mL of ~0.25 M NaOH solution (Fisher Scientific, Certified ACS Pellets, S318-3). A 125 mL Erlenmeyer flask (Fischer) was obtained and 0.999 g of potassium hydrogen phthalate (KHP, C8H5KO4) (Acros Organics, 99+%, Code: 417955000, Lot: A0358893) was added. A precision balance
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Three drops of the indicator, phenolphthalein, was added to the flask containing the dissolved KHP. The burette was then filled with ~0.25 M NaOH, using a glass funnel, and a small amount was drained in order to remove the remaining air from the tip. The initial volume of the burette was recorded to the nearest hundredth milliliter. The flask was placed under the burette and slowly, the NaOH was drained into the KHP solution, with constant mixing. The burette continued to be drained until the KHP solution changed from a clear color, to a persistent light pink, indicating the end …show more content…
Exactly 10.00 mL of vinegar (Stop and Shop Distilled White Vinegar, All Natural, 5% Acidity) was added to a 125 mL Erlenmeyer flask, using a 50 mL burette. Three drops of phenolphthalein indicator were mixed into the flask. The initial volume of the burette containing the NaOH was recorded to the nearest hundredth milliliter. The flask containing the vinegar was placed under the 50 mL burette and slowly, the NaOH was drained into the flask, with constant mixing. The burette continued to be drained until the acetic acid solution changed from a clear color, to a persistent light pink, indicating the end point of the titration. The final volume of the burette was then recorded to the nearest hundredth milliliter and the difference between the initial and final volumes was calculated; 31.35 mL or 0.03135 L. This volume in liters was multiplied by the concentration of the NaOH solution, 0.2838 M to determine the number of moles of acetic acid solution. This value was then divided by the volume of acetic acid solution originally added to the flask, 10.00 mL or 0.01 L, to obtain the molarity of the acetic acid solution; 0.8903 M. This process was replicated exactly a second and third time. The second titration required 31.39 mL of NaOH to be added, resulting in an acetic acid molarity of 0.8915 M. The third titration required 31.44 mL of
...ost likely to be battery acid. If it is water, it has a Ph level of around 7. For vinegar, the Ph level is approximately 2.4 - 3.4. Thus, once testing the liquid compare it with the Ph levels above to discover the mystery solution.
20.0cm3 of 0.10M ethanoic acid was pipetted into a conical flask. 3. 0.10M sodium hydroxide solution was titrated using phenolphthalein as indicator, until the solution was just turned pink. 4. A further 20.0cm3 of the same ethanoic acid solution was added to the flask and was mixed thoroughly.
neutralize 35ml of our base. Once we weighed out the KHP we then dissolved it
Titration of sodium carbonate and sodium hydroxide I was asked to find the concentration of a solution of sodium
Acid-Base Titration I. Abstract The purpose of the laboratory experiment was to determine equivalence. points, pKa, and pKb points for a strong acid, HCl, titrated with a. strong base, NaOH using a drop by drop approach in order to determine. completely accurate data. The data for this laboratory experiment is as follows.
* Pipette 25cm3 * 2 x 500cm3 beaker * Conical flask 250cm3 * Burette * White tile * Burette stand * Stand * Indicator * 300cm3 of Hydrochloric acid- standard solution (concentration of 0.05M) * Distilled water * Filter Paper * Stirring rod * Funnel Method: * Add 1g of Ca(OH)2 to 300cm3 of distilled water in a 500cm3 beaker. Keep stirring the solution till the solid stops dissolving. This leaves a saturated solution. * Filter off the excess solid into another 500cm3 beaker using a damp filter paper (distilled water). * Repeat the filtration of the solution till there is no solid left.
I rinsed the burette by opening the tap and allowing the HCl to flow, which released any air bubbles and cleaned the tip of the burette. After adding 5 drops of phenolphthalein indicator, the solution turned pink.
Some improvements to the experiment might be using Na Acetate or Na Citrate as buffers instead of KHPO4. The pH ranges are 4.5-5.5 and 4.7-5.5, respectively. This range falls closer to the ideal pH of 5, then KHPO4 (pH
10.0 g of cyclohexanol and 2 mL of conc.(85%) phosphoric acid were placed in a 50 mL ST round bottomed flask and the two were mixed by swirling.
Chemistry: Acid-Base Titration. Purpose: The objective of this experiment were: a) to review the concept of simple acid-base reactions; b) to review the stoichiometric calculations involved in chemical reactions; c) to review the basic lab procedure of titration and introduce the student to the concept of a primary standard and the process of standardization; d) to review the calculations involving chemical solutions; e) to help the student improve his/her lab technique Theory: Titration was used to study acid-base neutralization reaction quantitatively. In acid-base titration experiment, a solution of accurately KHP concentration was added gradually to another solution of NaOH concentration until the chemical reaction between the two solutions was completed. The equivalence point was the point at which the acid was completely reacted with or neutralized by the base.
vii. This would allow the determination of the percentage of citric acid in the lemon juice specifically, rather than the total acidity. The results of this could have been compared to those of the titration, and the contribution of citric acid to the overall initial acidity could have been determined.
Neutralization Experiment AIM:- To investigate how heat is given out in neutralizing sodium hydroxide (NaOH) using different concentrations of Hydrochloric Acid. Background Information:- Substances that neutralize acids are called alkalis. An acid is a substance that forms hydrogen ions (H+ ) when placed in water. It can also be described as a proton donor as it provides H+ ions. An example of an acid is hydrochloric acid (HCl), Sulphuric acid (H2SO4) etc.
The sample was subjected to steam distillation as illustrated in Figure 1. A total of 50ml of distillate was collected while recording the temperature for every 5.0 ml of distillate. The distillate was transferred into a 250ml Erlenmeyer flask and 3.0 g of NaCl was added. The flask was cooled and the content was transferred into a 250-ml separatory funnel. Then 25.0ml of hexane was added and the mixture was shaken for 5 minutes with occasional venting. The aqueous layer was discarded and the organic layer was left inside. About 25.0ml of 10% NaOH was then added and the mixture was shaken as before. The aqueous layer was collected and then cooled in an ice bath. It was then acidified with enough 6.00 M HCl while the pH is being monitored with red litmus paper. Another 25.0 ml of hexane was added and the mixture was shaken as before. The hexane extract was saved and a small amount of anhydrous sodium sulfate was added. The mixture was then swirled for a couple of minutes then filtered. A small amount of the final extracted was tested separately with 1% FeCl3 and Bayer’s reagent.
The purpose of this experiment is to use our knowledge from previous experiments to determine the exact concentration of a 0.1M sodium hydroxide solution by titration (Lab Guide pg.141).
AIM - To analyse some fruit and vegetable juices for the contents present in them. APPARATUS - Test tubes, burner, litmus paper, beaker, tripod stand, conical flasks, burette, pipette. CHEMICALS REQUIRED - 1. Fehling's solution A 2. Fehling's solution B 3.