chevron_left End-point the point in a titration at which the indicator changes colour, signalling that the reaction is complete chevron_right

The End-Point: The Colorful Finale of a Titration

The Basics of Titration

Imagine you are a detective trying to find out the exact concentration of a mysterious solution. Titration is your go-to investigative tool. It is a common laboratory method of quantitative chemical analysis used to determine the concentration of an identified analyte. The process involves slowly adding a solution of known concentration, called the titrant, from a burette to a known volume of another solution of unknown concentration, called the analyte, until the reaction between the two is complete.

The key apparatus used in a titration includes a burette, which is a long, graduated tube with a stopcock at its bottom for precise liquid dispensing, and a conical flask (or Erlenmeyer flask) that holds the analyte. The magic happens when an indicator is added to the analyte. This substance changes color in response to the chemical conditions in the flask, signaling the end-point.

The Indicator: The Messenger of Change

The indicator is the star of the show when it comes to detecting the end-point. Most indicators are weak acids or bases that have different colors in their acidic and basic forms. They are chosen because their color change occurs over a specific pH range that closely matches the pH expected at the equivalence point of the titration.

For instance, in a strong acid-strong base titration, the equivalence point is at pH 7. An indicator like bromothymol blue, which changes color between pH 6.0 and 7.6, would be a good choice. The end-point is recorded as the volume of titrant added when the color change persists, for example, when the solution in the flask turns from yellow to a stable blue.

End-Point vs. Equivalence Point: A Critical Distinction

It is crucial to understand that the end-point and the equivalence point are not the same thing, though they are often very close.

• Equivalence Point (Theoretical): This is the ideal point in the titration where the amount of titrant added is exactly stoichiometrically equivalent to the amount of analyte in the solution. For the reaction $H^+ + OH^- \rightarrow H_2O$, this is the point where moles of acid equal moles of base. It is a theoretical point calculated from the reaction's stoichiometry.
• End-Point (Experimental): This is the observed point where the indicator changes color. It is our practical signal to stop adding titrant.

The goal of a well-designed titration is to choose an indicator whose end-point is as close as possible to the true equivalence point. A slight mismatch between the two points results in a small, inherent error in the experiment known as the indicator error.

A Practical Application: Titrating Vinegar

Let's follow a real-world example to see the end-point in action: determining the concentration of acetic acid in vinegar.

1. Setup: A known volume of vinegar (the analyte, a weak acid) is placed in a conical flask. A few drops of phenolphthalein indicator are added, making the solution colorless.
2. Titration: A sodium hydroxide (NaOH) solution of known concentration (the titrant, a strong base) is slowly added from a burette to the vinegar.
3. The Reaction: The hydroxide ions from the NaOH react with the acetic acid ($CH_3COOH$) molecules: 
   $CH_3COOH + OH^- \rightarrow CH_3COO^- + H_2O$
4. The End-Point: Initially, each drop of NaOH is quickly consumed by the abundant acetic acid. As the reaction nears completion, the solution is mostly water and acetate ions. The moment the last of the acetic acid is neutralized, the next single drop of NaOH introduces free $OH^-$ ions into the solution. This slight excess of base causes the pH to rise sharply into the range where phenolphthalein changes color. The solution in the flask turns a permanent pale pink. This is the end-point.
5. Calculation: The volume of NaOH used is recorded. Using the concentration of the NaOH and the balanced chemical equation, the concentration of acetic acid in the vinegar can be accurately calculated.

Beyond Acid-Base: Other Types of Titrations

While acid-base titrations with color-changing indicators are the most common introduction to the end-point, the concept applies to other types of titrations as well.

Redox Titrations: These involve reduction-oxidation reactions. The end-point can be detected by a color change from the titrant itself or a special redox indicator. A classic example is the titration of iron(II) with potassium permanganate. The deep purple $MnO_4^-$ ion is reduced to nearly colorless $Mn^{2+}$. The end-point is the first permanent pink color that appears, indicating that all the iron(II) has been oxidized and there is now an excess of permanganate.

Precipitation Titrations: These involve the formation of a precipitate. A common method is the Mohr method for chloride determination, using silver nitrate as the titrant. Potassium chromate is used as an indicator. Silver chloride precipitates first as a white solid. Once all chloride ions have been precipitated, the next drop of silver nitrate reacts with chromate to form a brick-red silver chromate precipitate, signaling the end-point.

Important Questions

Conclusion

Footnote

¹ Analyte: The substance whose concentration is being determined in a titration. 
² Titrant: The solution of known concentration that is added from the burette during a titration. 
³ Stoichiometry: The calculation of relative quantities of reactants and products in chemical reactions. 
⁴ Burette: A graduated glass tube with a tap at one end, used for delivering known volumes of a liquid, especially in titrations. 
⁵ pH: A scale used to specify the acidity or basicity of an aqueous solution.