There are four significant regions in titration curves: the initial point, pre-equivalence point, equivalence point, and post-equivalence point. Each region's characteristics are related to the pH results obtained, and the points are named according to their position relative to the equivalence point. The equivalence point is where the solution becomes neutral, which is represented on the graph as a sudden increase and decrease in pH. At the initial point, the solution's pH equates to the number of H+ ions initially present in the analyze. At the pre-equivalence point, the iterant is added, and the solution resists drastic pH changes. For this experiment, a strong base (Noah) was added to a weak acid (KIP), resulting in a buffer-like phenomenon. At the post-equivalence point, excess Noah causes further pH increases. Continuous stirring is important in potentiometers titration becaus
...e it allows for incremental hotshots' formation and a higher concentration of iterant in a localized area compared to the fluid around it.
The usage of smaller increments of addition towards the equivalence point aids in achieving quicker endpoints. As the titration nears the equivalence point, pH changes become larger, especially at the equivalence point where the most significant change occurs. Continuously adding increments of 1 ml will not accurately determine the volume at which the equivalence point occurred. The determination of weak acid peak through potentiometers titration relates the potential between electrodes and iterant volume. This method ensures that similar amounts of iterant are used for weak and strong acids because it is the number of protons or H+ ions measured by pH that determines the equivalence point. Potentiometers titration is advantageous because indicators are not used
reducing errors due to human recognition failure of color change. Possible discrepancies in experimental and theoretical peak values can occur due to various reasons.The results may have been affected by human errors, including miscalculations, overvaluations, and failure to determine the correct pH value due to the fluctuation of the pH meter. The electrodes of the pH meter apparatus could also have been a source of error due to imperfect classroom conditions or failure to rinse the electrodes. Additionally, inaccuracies in the analytical balance resulting from its prior use could have led to errors. One member of the group accidentally contaminated the setup by touching the magnetic stirrer, but its effect on acidity cannot be determined. Standardization error due to incorrect measurement of final and initial values could also have caused significant changes. Lastly, preparing the electrodes inaccurately, such as failure to recalibrate them after being put in a new test solution, would have resulted in inaccurate readings due to previous pH values still affecting the pH meter.
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