Acid and Base

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Experiment 4:

ACIDS AND BASES: PH Measurements and Macroscale Titration

CHM023L – A12
Group no. 6
Members:| Contributions:|
| Conclusion|
| Recommendation|
| Tables and figures with analysis|
| Principles, Equation|
| Abstract, tables|

Date Performed: February 28, 2012 - Tuesday
Date Submitted: March 6, 2012 - Tuesday
Submitted to:


This experiment introduces us the pH measurement and application of macroscale titration of aqueous solutions of acids and bases. The first part of the experiment was applying macroscale titration to calculate the mass of an HCl solution consumed that was initially done by pounding the tablets and getting a weight of 0.15 grams. The powdered form antacid was mixed with 100mL of 0.1M HCl in the erlenmeyer flask. It was stirred and heated until it boiled then was filtered. Two drops of phenolphthalein was added to the filtered solution. It was titrated with 0.1M NaOH until its color turned pink. The procedure was repeated for the second trial. The second part of the experiment was acquiring the pH of solutions using calorimetric measuring tools. First is by using pH paper. The pH paper was dipped into the ten drops of 0.1M HCl then read the pH using the chart. The procedure was repeated using 0.1M NaOH, 0.1M CH3COOH, and 0.1M NH4OH. Next, the pH of 30mL of 0.1M HCl, 0.1M NaOH, 0.1M CH3COOH, and 0.1M NH4OH(placed separately into four beakers) was measured using pH meter. And lastly, the pH of the solutions was calculated using the obtained values from the procedures. In the third part, the percent of acetic acid in vinegar was determined. The buret was rinsed with 0.2M NaOH and was later on filled with the same solution up to its mark of 0.0 ml. 5ml of commercial vinegar, 45ml distilled water and two drops of phenolphthalein was placed in the Erlenmeyer flask. The vinegar was titrated with 0.2 NaOH until its color turned to light pink and then the volume of NaOH used was recorded. The procedure was repeated for the second trial.


Vinegar, aspirin, Vitamin C, baking soda and ammonia are just few of the familiar substances that we use every day that can be considered as acids or bases. Acids have a sour taste, cause color changes in plant dyes, react with certain metals to produce hydrogen gas and react with carbonates and bicarbonates to produce carbon dioxide gas. On the other hand, bases have a bitter taste, feel slippery and also cause color changes in plant dyes.

In 1884, Svante Arrhenius proposed the first theoretical model for acids and bases. According to the Arrhenius theory, pure water dissociates to some extent to produce hydrogen ions, H+ and hydroxide ions, OH-. When this occurs, equal amounts of H+ and OH- ions are produced: H2O ↔ H++ OH-

The list of strong acids are as follows:
* HCl- HNO3
* HBr- H2SO4
* HI- HClO4
The list of strong bases include:
* LiOH- RbOH
* NaOH- CsOH
Strong bases are all composed of a Group IA metal and a hydroxide. Based on the solubility rules, all compounds containing a Group IA metal are soluble. That’s what these strong bases have in common. In addition to that, all of these acids and bases listed above dissociate 100% in water.

pH is used to describe the acidic or basic nature of a solution. The H in pH means the concentration of hydrogen ion, H+ in a solution. The p in pH is just a mathematical trick to make a decimal number a whole number and it is in terms of –log (logarithms of base ten). The pH is the logarithm of the concentration of hydrogen ions, which can be express as: pH=-log[H+]. Thus, when the pH has low values, the concentration of hydrogen ion is high.

Titration is a convenient quantitative method for accurately determining unknown concentrations of solutions. A necessary requirement for...
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