In this lab exercise we will evaluate the effectiveness of several indicators for the determination of the point of completion of a specific acid-base neutralization reaction. We will also determine the unknown concentration of the strong base NaOH by its reaction with a known amount of the weak acid, potassium acid phtalate (HKC8H4O4, abbreviated KHP). This will be accomplished using the titration method. The KHP solution will be created and its volume and concentration recorded. The KHP solution will be poured in a flask along with a few drops of one of three indicators we will be evaluating. The NaOH solution will be poured into a buret (with volume markers) and will be used as the titrant. The strong base will be added slowly to the acidic solution, gradually neutralizing the acid. The volume of base added can be determined by the difference in the initial and final volume marks on the buret.
At a certain volume of added NaOH, all the KHP acid will be neutralized due to the large equilibrium dissociation constant (Kb) of the base. This point of titration is referred to as the equivalence point. Considering the 1:1 stoichiometry of this acid-base reaction
NaOH(aq) + C6H4(COOH)(COOK) (aq) C6H4(COONa)(COOK)(aq) + H2O(l)
the point of equivalence is the point of titration when the number of moles of NaOH (Na) added is equal to the number of moles of KHP (Nb) in the solution. The number of moles of KHP in the solution can be calculated very simply by dividing the known mass of the sample in the solution by its molecular mass. The unknown concentration of the NaOH can then be calculated in the following manner:
At the point of equivalence of a reaction of 1:1 stoichiometric ratio, Na = Nb.
The number of moles of a solute is the concentration times the volume (N = Vc ).
Thus Vaca = Vbcb.
Knowing all other variables we can solve for cb by restructuring the previous equation as cb = caVb/Va.
However, in order to determine the equivalence point the dissociation of the indicators being used must coincide with the pH at the equivalence point. The indicator, a weak organic acid, will dissociate at a certain pH. The dissociation of an indicator is concurrent with a color change or some other physical change which informs the observer of the solution’s approximate pH. A decreased amount of H3O+ (a product of acid dissociation) makes it more probable for the dissociation reaction of the indicator to occur since equilibrium must be maintained. Depending on the specific dissociation constant of each indicator a different H3O+ concentration (and thus pH) will trigger the dissociation of each indicator.
Since we do not know the dissociation equilibrium of each indicator, we cannot calculate the exact range of pH at which a color change will appear. Thus we will must repeat the titration experiment with a pH-meter and record the pH of the acid-base solution per milliliter of NaOH added. The results of this part of the experiment will be used as the correct reference in order to determine which indicators change color at a pH range that coincides with the approximate pH at the equivalence point of the given titration. The calculations of the concentration of NaOH must thus exclude the unsuitable indicator(s).
Method and Explanations
·Acid Base Titration with Different Indicators
We first created a solution of NaOH by adding 10ml of 6M NaOH to 500 ml of distilled water. This solution was poured into a plastic bottle with a lid and was shook vigorously for a few minutes. It is essential that the solution be homogeneous for the titration experiment to be successful for in order to investigate and calculate the NaOH concentration it must be constant throughout the solution. We then rinsed and dried four clean beakers and labeled them from one to four. We weighed precisely 0.50 g of KHP in each beaker, with an accuracy of + .001g. The mass of KHP added to each beaker was recorded. 50ml of distilled water was then added to each beaker.
The solution was then swirled carefully in order to dissolve the solute. Since our beakers were large we were able to stir the solution contents with the magnetic stirrer without the fear of spilling any solution. This method was more effective and less time consuming than swirling the beaker. This solution was then poured into a 250ml Erlenmeyer’s flask and a magnetic stir bar and several drops of phenolphthalein indicator were added. The flask was then placed on the magnetic stirrer with a white paper under the flask to allow for more contrast and facilitate the detection of a color change.
Once the experimental setup was complete, a 50ml buret was rinsed twice with 10ml of the NaOH solution from the plastic bottle. The buret was then filled with the NaOH solution and the initial NaOH volume mark was recorded. With the magnetic stirrer still on, we then placed the buret directly above the opening of the flask and slowly add NaOH to the acidic solution in the flask by slightly turning the stopcock. (It is essential that the magnetic stirrer be mixing the solution continuously so that there is no delay due to the time it take for the hydronium ions to collide with the hydroxide ions.) The instant the color of the solution changes permanently from clear to pink the stopcock must be closed and the final NaOH volume mark must be recorded. The resultant solution was then poured into the designated waste beaker, eventually to be discarded in the waste container.
The color of the solution should fluctuate for a few seconds from clear to pink and back again but this is simply because the equilibrium of the solution was temporarily thrown out of balance with the presence of more hydroxide ions (OH-). The additional hydroxide ions neutralized the hydronium ions (H3O+) in the solution. This causes a temporary lack of hydronium and thus shifts the dissociation equilibrium of the indicator. The indicator dissociates temporarily revealing an instance of pink coloration. However, as there was still some KHP present, the acid, having a stronger dissociation constant, dissociates and H3O+ is produced. The indicator’s dissociation reaction is then forced backward and the solution once again appears clear.
This experiment was then repeated using two other indicators: bromothymol blue and methyl orange. For the latter experiment, twice the amount of indicator drops was added. This makes it easier to detect the instant the color change occurs since the methyl orange indicator continually and gradually changes its color. It is thus difficult to determine exactly when the first permanent color change occurs. The gradual color changes of this indicator may be due to multiple steps of dissociation which may occur for example if the acid can release more than one H+ ion (this is simply a speculation, do you know why the color change is so gradual?).
The experiment was then repeated once more. This time no indicator was used. Rather, a pH meter was used to devise a correct reference with which to determine which indicators are appropriate for the determination of the equivalence point for this specific reaction. The experimental setup is the same as previously however, we also inserted the probe of a pH meter in the flask.
The probe was rinsed with distilled water and dried. We then proceeded with the experiment recording the pH of the solution after every ml. of added NaOH. With these results we then constructed a titration curve and determined which indicator(s) was inappropriate for this experiment so that we may calculate the experimental concentration of NaOH excluding the unsuitable indicator(s).
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