Investigation of the Suitability of Indicators and Instrumental Detection of Equivalence Points in Acid-Base Titrations

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

pH was recorded every time 1.00 mL of NaOH was added to beaker. When the amount of NaOH added to the beaker was about 5.00 mL away from the expected end point, NaOH was added very slowly. Approximately 0.20 mL of NaOH was added until the pH made a jump. The pH was recorded until it reached ~12. This was repeated two more times. The pKa of each trial are determined using the graphs made on excel.

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.

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

To begin, three sets ofabout 0.3000g of KHP are weighed out on an analytical balance. Put the three sets of KHP into three separate, labeled flasks. All three sets of the KHP is then dissolved with approximately 50mL of deionized water. Next, a buret is used to start the actual titration. Buret is initially filled to 0.00mL mark with the NaOH solution, this is recorded as initial volume. Next, add 2-3 drops of phenolphthalein indicator into each of the three flasks containing KHP. A magnetic stir bar is then added to the first flask, and placed above a stir plate. Everything is positioned under the buret. Stirrer is put on medium speed and the titration can start. Slowly release the NaOH into the KHP flask. As the end point is reached, a pink color will be seen in the flask. When the lightest pink possible remains in the solution for more than 30 seconds titration is complete. The final volume is recorded, and the same steps are taken for the other two sets of KHP solution. Finally, blank titration is completed to determine deviation.

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

The first titration (with use of the PASCO program) is the most accurate titration that the lab group can perform. The titration curve gives many different variables, with accuracy that passes what humans can do without a machine. But the problem with this titration is that the curve or change to neutrality can often be sudden and out of nowhere. This means that it is very hard to just titrate a solution to become neutral, and often will change into the opposite property, essentially an acid to base and vice versa. The titration with use of an indicator is less accurate, but shows a more visual representation about how a titration takes place. The change in color is a less accurate indication of the volume it takes a solution to change to be neutral, but the visual change is helpful in a classroom laboratory setting. The preference is the first titration method, because it is more accurate and can give more reliable data, since the PASCO machine and drips added are not affected by human

Within an acid-base titration the titration curve resembles the strengths of the corresponding acids and bases. A strong acid will correspond with a weak conjugate base, and a weak conjugate acid will correspond with a strong base. This is based on the Bronsted-Lowry model. The weak acid will donate protons to the hydroxide ion. Weak acids will have a low Ka value, the Ka value is the tendency of the acid to dissociate:

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

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

The purpose of this experiment was to determine how much KMnO4 was needed to titrate approximately 1 mL of an Unknow X101 concentrated solution of Oxalic Acid. A standardized KMnO4 solution was used on a known solution of Oxalic acid to help determine the unknown percent oxalic acid in unknown X101. The unknown sample for this experiment was sample x101 which theoretically was a % Oxalic Acid dehydrate sample but, the average of all three trials determined it to be a 6.7% percent Oxalic acid.

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.