As the tables on the previous page show, there was no obvious equivalence point for the latter two reactions. This will be explored in the data processing section.
Phenolphthalein and methyl orange are the two indicators that were used during this investigation. Phenolphthalein was used for the strong acid + strong base experiments and during the weak acid + strong base experiments because the pH jump was expected to be between 2 and 12 for both experiments, and since phenolphthalein’s range is 8.2 to 10 (according to the IB chemistry data booklet), it would be a suitable indicator. Methyl orange was used for weak acid + weak base experiments and strong acid + strong base experiments. The anticipated pH jump for these experiments were estimated to be between 1 and 2.5 for the strong acid + weak base experiment and 2.5 to 4.5 for weak acid +weak base experiments. Since methyl orange has a range of 3.2 to 4.4, it seemed the most appropriate indicator.
* Each time a drop of base was added to the strong acid + strong base and weak acid + strong experiments, there was a hint of purple which would disappear after swirling.
* As the equivalence point for the strong acid + strong base and weak acid + strong base was reached, the solution permanently turned a shade of purple.
* In the weak acid + weak base and strong acid + weak base reactions, the methyl orange did not show any colour change until the equivalence point was reached, when the methyl orange turned from red to yellow.
In each experiment, there is a particular reaction which occurs. Each reaction is unique due to its reactants, and as a result, produces unique products. These products affect the pH, and so cause the pH graph to fluctuate. All applicable reactions are show below. The ions which eventually react to produce water, H+ and OH-, are shown in red. These ions may be in a molecule, but while in water, the molecule dissociates, meaning the ion is free and can affect pH. In all experiments involving a weak acid and/or base, there are at least two reactions. The first reaction shows the dissociation of the molecule through a reversible reaction, and the second shows the appropriate ion reacting with its complementary ion to form water.
Strong Acid + Strong Base:
HCl + NaOH = NaCl + H2O
Strong Acid + Weak Base:
NH3 + H2O – NH4+ + OH-
HCl + OH- = H2O + Cl-
Weak Acid + Strong Base:
CH3COOH + H2O – CH3COO- + H+
H+ + NaOH = H2O + Na-
Weak Acid + Weak Base:
CH3COOH – CH3COO- + H+
NH3 + H2O – NH4 + OH-
H+ + OH- = H2O
Below are graphs which give the pH change a visual representation.
Graph Showing pH Against Volume for Strong Acid + Strong Base
Graph Showing pH Against Volume for Weak Acid and Strong Base
In the first two graphs, a large increase in pH can be seen, meaning this is the equivalence point, and we can determine the buffer regions, equivalence points and 1/2 equivalence points. The buffer region is the region of the graph that plateaus or exhibits slight changes in pH before the sharp increase. This is because the added alkali’s OH ions are reacting with the H+ ions in this region, producing H2O and resulting in a very small amount of pH change. The 1/2 equivalence is shown with the red asterisks. At this point, pKa = pH. This is because the concentration of Cl- or CH3COO-, respectively, is equal to the concentration of HCl and CH3COOH, respectively. This means the two cancel out when divided in the expression for Ka, meaning Ka = [H+]. The x intercept is the initial pH.
Graph Showing pH Against Volume for Strong Acid and Weak Base
Graph Showing pH Against Volume for Weak Acid and Weak Base
On the two graphs in the previous page, there is no definite equivalence point, so it is impossible to find the 1/2 equivalence point and the buffer region. It is also possible that the whole graph is the buffer region, and more alkali was needed to reach the equivalence point. The x intercept on the graph is the initial pH of the solution.
Conclusion and Evaluation
From this investigation, I can conclude that the pH at the equivalence point of titrating a strong acid with a strong base is around 2.14 and that the pH at the equivalence point of titrating a weak acid with a strong base is around 5.56. In the first two experiments, a strong and a weak acid were used. The strong acid, HCl, dissociated into H+ and Cl- ions, whereas the molecules of the CH3COOH do not fully dissociate, so it was considered a weak acid. Fully dissociated molecules such as HCl produce more H+ ions than partially dissociated ones such as CH3COOH, and so these have a higher concentration of hydrogen ions. In a solution of ethanoic acid, an equilibrium is present, meaning both forward and backward reactions take place, and so there will be less H+ ions in this solution, since it will sometimes be in a molecule and sometimes an individual ion. This does not apply to HCl, which permanently dissociates into H+ and Cl-. This concentration of H+ ions immediately affects the pH due to their inverse relationship.
The NaOH was considered as the strong base since the molecules fully dissociate into Na+ and OH- molecules, whereas the molecules of NH3 cannot directly produce OH-, and must react with water to produce ammonium and OH-. The same theory applies to bases as acids; fully dissociated molecules such as NaOH produces more OH-, which means there is a smaller concentration of H+, reducing the concentration and so heightening the pH due to the inverse relationship. Ammonia reacts to form ammonium and OH- in a reversible reaction, so at anytime, there is a small amount of OH- present, so the pH does not increase significantly.
In the first graph (Strong acid and strong base), the hydroxide ion dissociated from the NaOH and the hydrogen ion from the HCl combine to form water at the buffer region. Immediately before the equivalence point, many of the H+ ions have been used by the OH to form water, keeping the solution at a relatively low pH. When the next drop of NaOH is added, there are not enough H+ ions remaining in the solution to react with the newly introduces OH ions, and so the solution becomes basic very suddenly. This applies to the second graph (Weak acid and strong base) as well, but, theoretically, less NaOH would have needed to be added due to the lesser concentration of H+ ions compared to the first experiment.
In the third graph (strong acid and weak base) no apparent equivalence point can be observed. This is likely due to the fact that not enough alkali was added to fully titrate the solution. Since we were titrating a strong acid with a weak base, there were few OH ions present, but many H+ ions. This means a high volume of NH3 would have needed to be added in order to produce enough OH- and form water from the H+ ions. Judging by the line in the graph, which starts to rise at around 22ml, the equivalence point was being approached. In the final graph (weak base and weak acid), there is a relatively general trend of slowly increasing pH. The reason for this very gradual increase is due to there not being
many H+ nor OH- ions, thus resulting in there being very few water molecules formed. There is a slight increase in pH as more alkali is added, but there is no equivalence point because the NH3 does not dissociate enough to be noticeable.
The literature value for titrating 25ml of a strong acid with a strong base is that 25ml of NaOH needed to be added, with a relatively flat buffer region and a long equivalence point. and an equivalence pH of 7. The literature value for titrating 25ml of a weak base with a strong acid is that 25ml is needed, with a steeper buffer region and shorter equivalence point and an equivalence pH of 8.73. These results are similar to the acquired results since the amount needed to change the solution in experiments one and two was 8, and there was 10ml of solution in the beaker. These numbers are very close, similar to the 1:1 ratio of acid volume and amount of base that was needed volume. The graphs also exhibit similar patterns of a steeper buffer region and shorter equivalence point for the weak acid and strong base. The acquired equivalence point for strong acid and strong base is 7.75 and for weak acid and strong base it is 9.5, both values are relatively close to the literature values.
The literature value for titrating 25ml of HCl with NH3 of the same concentration is that 25ml would be needed. It is likely that more NH3 would have been needed to be added in this investigation in order to determine the equivalence point and the other regions of the graph. The literature value for the weak acid and weak base titration was that there was a pH equivalence at of 7.1, but that there was a general increase in the graph, and the equivalence point occurred when 25ml of CH3COOH was added. The final graph reflects this, and shows similar behaviour to the literature values.
Source for literature values:
Owen, Steve. Chemistry for the IB Diploma. Cambridge: Cambridge UP, 2011. Print.
There are several aspects of the investigation which could be improved upon. The first point is the fact that the set of acquired data has largely been limited by its numbers. Due to there being only one trial per titration, we assume that this titration applied to all titrations using these chemicals, where, in fact, it does not. We also assume, by assigning each chemical as a strong or weak acid and a strong or weak base, that these titration curves apply to all strong acids or weak bases etc. Titrating H2SO4 with KOH may have lead to different results and a different titration curve, but we immediately assume this applies to all appropriate chemicals. Another area for improvement would be the amount of alkali added in the strong acid and weak base titration. By adding more alkali, the equivalence point could be determined and, by doing so, the buffer region and 1/2 equivalence point could also be determined. Since too small an amount of NH3 was added, the equivalence point was never reached, and so the other regions of the graph could not be determined. Uncertainties were not an issue in this investigation; all equipment used had negligible uncertainties, and so had minimal effects on the investigation. A second assumption made was that all chemicals used in the investigation were pure and reacted only with their
respective acid or base. Had the chemicals not been pure and had reacted with something other than their respective chemical, the balance between the OH- and H+ ions might have been upset, causing incorrect results to be acquired.
Steps can be taken in order to remove or minimize the above weaknesses. An ideal amount of repetitions for each titration would be 5. This would ensure that the chemicals being tested show the same trends in the graphs and that the graph is indeed the correct one for the titration in question. The experiment could be carried out with a number of different weak and strong acids and bases, not only the ones used in this investigation. This would provide more substantial results and eliminate the assumption that the titration curves in this experiment apply to all strong and weak acids and bases titrations. More NH3 should be added in the weak base and strong acid titration so the equivalence point can be reached and the data can be properly processed. As mentioned before, uncertainties were not a problem in this investigation, but they could be further reduces by using more accurate equipment which can measure the volume of a liquid to a number of decimal places. The last area for improvement, assumptions of purity and other reactions, can be monitored by measuring the actual yield of products and comparing it to the theoretical yield of products, the difference would be assumed to be due to purity and other reactions. This is calculated with the formula below:
Percentage Difference = (Actual Percentage – Theoretical percentage) / Actual Percentage.
Error Uncertainty or Assumption
Data set limitations
Equivalence point not reached for one graph
By adding more alkali, the equivalence point could be determined and, by doing so, the buffer region and 1/2 equivalence point could also be determined.