Why is phenolphthalein used

Since one drop of added titrant will cause this large change, even though the change in color of phenolphthalein does not occur right on the equivalence point, it is within approximately one drop. This kind of uncertainty is "acceptable uncertainty" in using titration to volumetrically determine concentrations. To clarify what I mean by "acceptable uncertainty", you should realise that each of your measurements has some kind of uncertainty to them:.

Are you able to state the volume added from the burette to arbitrary precision or is there some uncertainty beyond the two decimal places given by the graduated lines?

Sign up to join this community. The best answers are voted up and rise to the top. Stack Overflow for Teams — Collaborate and share knowledge with a private group. Create a free Team What is Teams? Learn more. Why is phenolphthalein an appropriate indicator for titration of a strong acid with a strong base?

Ask Question. Asked 7 years, 8 months ago. In this experiment, a phenolphthalein color indicator will be used. Phenolphthalein is colorless in acidic solutions and pink in basic solutions.

Phenolphthalein is also used in forensic crime scene analysis to detect the presence of blood, Kastle-Meyer test.

What is the purpose of the phenolphthalein in an acid base titration? Phenolphthalein is often used as an indicator in acid—base titrations. For this application, it turns colorless in acidic solutions and pink in basic solutions.

It belongs to the class of dyes known as phthalein dyes. How is phenolphthalein solution typically used in a titration? A commonly used indicator for strong acid-strong base titrations is phenolphthalein. Solutions in which a few drops of phenolphthalein have been added turn from colorless to brilliant pink as the solution turns from acidic to basic.

Several drops of an indicator are added to the acid and mixed by swirling the flask. The indicator phenolphthalein, whose range spans from pH 8 to 10, therefore makes a good choice for this type of titration.

As we know, phenolphthalein shows pink colour in basic solutions, it will show pink colour in soap solution. Phenolphthalein is an indicator — a chemical which changes colour depending on whether it meets an acid or a base. It turns purple if it meets something basic, such as ammonia; it stays colourless if it meets an acid like vinegar or a neutral substance like water.

It is due to the formation of ions that the solution turns pink. Starch is a viable indicator in the titration process because it turns deep dark blue when iodine is present in a solution. When starch is heated in water, decomposition occurs and beta-amylose is produced.

The intensity of the color is concentration dependent, so if a large amount of phenolphthalein is present the color could be very strong at the end-point. However, the presence of a large amount of phenolphthalein is not usually the reason a titration mixture has an intense color. Phenolphthalein is an indicator of acids colorless and bases pink.

Sodium hydroxide is a base, and it was in the pitcher at the beginning, so when added to the phenolphthalein in beakers 2 and 4, it turned pink top half of the graphic. An aqueous solution of sodium hydroxide, NaOH aq , is a strong base.

Consider thymol blue pH range 8. Phenolphthalein is often used as an indicator in acid-base titrations. It is slightly soluble soluble in water. It turns acidic in acidic solution and turns pink in basic solution. A small amount of the acid solution of known concentration is placed in the burette this solution is called the titrant. A known volume of base with unknown concentration is placed into an Erlenmeyer flask the analyte , and, if pH measurements can be obtained via electrode, a graph of pH vs.

In the case of titrating the acid into the base for a strong acid-weak base titration, the pH of the base will ordinarily start high and drop rapidly with the additions of acid. As the equivalence point is approached, the pH will change more gradually, until finally one drop will cause a rapid pH transition through the equivalence point.

If a chemical indicator is used—methyl orange would be a good choice in this case—it changes from its basic to its acidic color. Titration of a weak base with a strong acid : A depiction of the pH change during a titration of HCl solution into an ammonia solution. The curve depicts the change in pH on the y-axis vs. In strong acid-weak base titrations, the pH at the equivalence point is not 7 but below it.

Polyprotic acids, also known as polybasic acids, are able to donate more than one proton per acid molecule. Monoprotic acids are acids able to donate one proton per molecule during the process of dissociation sometimes called ionization as shown below symbolized by HA :.

Common examples of monoprotic acids in mineral acids include hydrochloric acid HCl and nitric acid HNO 3. On the other hand, for organic acids the term mainly indicates the presence of one carboxylic acid group, and sometimes these acids are known as monocarboxylic acid.

Polyprotic acid are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Certain types of polyprotic acids have more specific names, such as diprotic acid two potential protons to donate and triprotic acid three potential protons to donate.

For example, oxalic acid, also called ethanedioic acid, is diprotic, having two protons to donate. If a dilute solution of oxalic acid were titrated with a sodium hydroxide solution, the protons would react in a stepwise neutralization reaction. Neutralization of a diprotic acid : Oxalic acid undergoes stepwise neutralization by sodium hydroxide solution.

If the pH of this titration were recorded and plotted against the volume of NaOH added, a very clear picture of the stepwise neutralization emerges, with very distinct equivalence points on the titration curves. Titration curve for diprotic acid : The titration of dilute oxalic acid with sodium hydroxide NaOH shows two distinct neutralization points due to the two protons.

Oxalic acid is an example of an acid able to enter into a reaction with two available protons, having different Ka values for the dissociation ionization of each proton. A diprotic acid dissociation : The diprotic acid has two associated values of Ka, one for each proton.

Likewise, a triprotic system can be envisioned. Each reaction proceeds with its unique value of K a. Triprotic acid dissociation : Triprotic acids can make three distinct proton donations, each with a unique Ka.

An example of a triprotic acid is orthophosphoric acid H 3 PO 4 , usually just called phosphoric acid. Another example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. An indicator is a weak acid or a weak base that has different colors in its dissociated and undissociated states. There are many methods to determine the pH of a solution and to determine the point of equivalence when mixing acids and bases.

These methods range from the use of litmus paper, indicator paper, specifically designed electrodes, and the use of colored molecules in solution.

Other than the electrodes, all of the methods are visual and rely on some fundamental changes that occur in a molecule when the pH of its environment changes. In general, a molecule that changes color with the pH of the environment it is in can be used as an indicator. In this reaction, adding acid shifts the indicator equilibrium to the left.

Conversely, adding a base shifts the indicator equilibrium to the right. In the case of the indicator methyl orange, the HIn is colored red and the ionized In — form is yellow.

Methyl orange : The molecule methyl orange is commonly used as an indicator in acid-base equilibrium reactions. In base form, on the left in the figure, the color is yellow. Adding a proton yields the structure on the right, colored red. Note that this color change occurs over the pH range from approximately The eye is sensitive to color changes over a range of concentration ratios of approximately or over two pH units. Below pH 2. Because of the subjective choice determination of color, pH indicators are susceptible to imprecise readings.