Analysis of Acid by Titration with Sodium Hydroxide

Purpose/Hypothesis: The purpose of this experiment is to use both cabbage juice and pH paper to determine the pH of household items. This way, we can tell which products are basic and which one are acidic. If we use cabbage juice as an universal pH indicator by comparing it to pH paper then pH determined by the cabbage juice will be unstable because by using cabbage juice, it can be different depending on how diluted it is.

The hypothesis is; as we know the concentration of sodium carbonate (Na2CO3) we can obtain the concentration of hydrochloric acid using the titration of a standard solution.

We only added a small amount of HCl to the water and sodium chloride. We did not continue to add more HCl after a significant drop in pH was recorded. We added a total of 2 mL of HCl to both H20 and NaCl before the pH changed. The 1 gram solution of sodium acetate and acetic acid changed after a 8 mL, and the other two never dropped before we reached our total of 10 mL HCl.

Table 2: Consists of color extract taken from a red cabbage for a natural indicator. The pH reading that was measured by using the pH meter and the result of the pH reading to determine whether the solution was acidic or basic.

Chemistry 102 is the study of kinetics – equilibrium constant. When it comes to the study of acid-base, equilibrium constant plays an important role that tells how much of the H+ ion will be released into the solution. In this lab, the method of titrimetry was performed to determine the equivalent mass and dissociation constant of an unknown weak monoprotic acid. For a monoprotic acid, it is known that pH = pKa + log (Base/Acid). When a solution has the same amount of conjugate base and bronsted lowry acid, log (Base/Acid) = 0 and pH = pKa. By recording the pH value throughout the titration process and determining the pH at half- equivalence point, the value of Ka can be easily calculated. In this experiment, the standardized NaOH solution has a concentration of 0.09834 M. The satisfactory sample size of known B was 0.2117 g. The average equivalent mass of the unknown sample was found to be 85.01 g, pKa was found to be 4.69, which was also its pH at half-equivalence point and Ka was found to be 2.0439×〖10〗^(-5). The error was 1.255% for equivalent mass and 0.11% for Ka. In other word, the experiment was very precise and accurate; the identity of the unknown sample was determined to be trans-crotonic by the method of titrimetry.

Methods: First, a burette, ring stand, clamp, and an empty flask were obtained. The burette, with the valve closed, was attached to the ring stand with a clamp, and the empty flask was placed below the burette. Next, 50mL of the NaOH solution were poured into the burette, and a small bit was drained into the empty flask to ensure that the tip of the burette was also full of NaOH solution. The volume of the NaOH in the burette was recorded. Next, approximately 0.6 grams of KHP were massed poured into an empty 125mL flask. Two drops of an indicator solution were added to the KHP

In order for the media to show the change in acidity the solutions are modified and include an indicator chemical. This indicator will change color depending on the ph level of the media it is in. For all the media used in this experiment, the indicator changes to a yellow color when in the presence of an acid and turns magenta/pink when in the presence of a base or alkali.

When handling the Acid or Base to avoid getting any substance in eyes: Use protective eyewear

The purpose of the experiment was to determine how a buffer works and how to use an acid-base indicator. The way a buffer works was determined by observing the changes in pH of solutions of different concentrations weak acids and their conjugate bases to determine how a buffer affects the pH change. The solution of 10 mL of 0.20 M CH3COOH and 10 mL of 0.20 M CH3COONa had slighter changes in pH than the solution of 10 mL of 0.0020 M CH3COOH and 10 mL of 0.0020 M CH3COONa. Both of these solutions were buffers, shown because they had slighter changes in pH than the solutions with only the weak acid or conjugate base and water. The determination of how buffers work was also tested with observing that the solution of NaC4H3O4 and Na2C4H2O4 had smaller

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

By using acid-base titration, we determined the suitability of phenolphthalein and methyl red as acid base indicators. We found that the equivalence point of the titration of hydrochloric acid with sodium hydroxide was not within the ph range of phenolphthalein's color range. The titration of acetic acid with sodium hydroxide resulted in an equivalence point out of the range of methyl red. And the titration of ammonia with hydrochloric acid had an equivalence point that was also out of the range of phenolphthalein.. The methyl red indicator and the phenolphthalein indicator were unsuitable because their pH ranges for their color changes did not cover the equivalence points of the trials in which they were used. However, the

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

In this lab a acid-base indicator phenolphthalein was used to determine endpoint of a reaction HCl(aq) and KOH(aq). At the end point all of the HCl(aq) would have reacted with KOH(aq), and the pH becomes 7. The phenolphthalein would changed colours from colourless to pink indication when enough KOH(aq) was added. The purpose of numerous trials was to use the average volume of the 3 trials with similar measurements.

An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid

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.