Titration Essay

A titration is used to determine the amount of acid in a given solution. This is done by titrating a measured volume of acid (in this instance, acetic acid (CH3COOH)) with a solution of a strong base (usually sodium hydroxide (NaOH)), of a known concentration. The NaOH is added in small aliquots until the acid has been neutralised, and this can be determined with an indicator dye, such as phenolphthalein, or a pH meter (Nelson & Cox, 2008 pg58). In this practical, a pH meter was used and this allows for the acidity or alkalinity of a solution to be measured, and this was more accurate than using an indicator dye. The use of the pH meter in titration is generally preferred more over the visual indicator because the equivalence point can be measured

The pH meter needed to be calibrated before the titration, and this was done by using coloured standards of pH 4.0, 7.0 and 10.0. The NaOH used within this practical was measured out in pellet form, and the amount needed was 0.4g of 0.1M NaOH. The NaOH was then dissolved into 100ml of distilled water by using a magnetic stir bar and a magnetic stirrer, which mixed the solution for around 120 seconds. After the NaOH had been dissolved, 25ml of 0.1M CH3COOH was measured into a measuring cylinder and was then transferred into a 100ml beaker. This was also placed onto the magnetic stirrer and a clean magnetic stirrer bar was then added to avoid any contamination before the NaOH had been added. The calibrated pH meter was then added to the CH3COOH and the initial pH reading was then taken, which was 2.8pH. By using a p200 pipette, 500µl aliquots of NaOH were then added to the CH3COOH solution, and 30 seconds were left between each aliquot to ensure the pH meter registered the changed pH. The pH of the solution was then recorded on a graph of pH vs volume of 0.1M NaOH added. After this, another 500µl (0.5ml) of NaOH was added and recorded until 40ml had been added (Thorne, A.

The shape of any titration curve of a weak acid can be described by using the Henderson-Hasselbach equation. This equation was useful in estimating the pH of a buffer solution (CH3COOH), and allows for an understanding of the acid-base balance within the blood and tissues of vertebrates. This equation fits the titration curve of all weak acids and showed why the pKa of a weak acid was equal to the pH of the solution at the midpoint of the titration (Nelson & Cox, 2008 pg.60-61). When the titration began, the CH3COOH was partly ionised. This means that it was mostly in the form of CH3COOH. When the NaOH was added, the amount of CH3COOH that was dissociated to form acetate ions and H+ ions increased. These ionised protons (H+) reacted with the added OH from the NaOH and formed H2O, which then neutralised them (Joesten, M. 2007). Where the slope of the graph was shallow, the pH changes very slowly with the addition of NaOH and this was where the CH3COOH displayed its maximum buffering ability. This occurred when the pH was resisting change from the addition of NaOH, and it occurred at this point because the amount of CH3COOH was equal to the amount of acetate ions (CH3COO-) (Thorne, A. 2016). This point is also known as half the equivalence point, and this was where half the NaOH added equalled half of the total