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Chapter 8: Problem 25

The indicator used for titration of weak base and strong acid is (a) thymol blue (b) methyl orange (c) phenolpthalein (d) fluorescein

Short Answer

Expert verified
The indicator used is (b) methyl orange.

Step by step solution

01

Understand the Problem

We need to pick an indicator suitable for the titration of a weak base with a strong acid. An indicator changes color over a specific pH range, and we want one that changes around the equivalence point of the titration.
02

Identify the pH Range for Weak Base-Strong Acid Titration

In a titration between a weak base and a strong acid, the equivalence point occurs in an acidic pH range, typically around pH 4 to 6.
03

Evaluate the Indicators

- Thymol blue has a pH range of 8.0 to 9.6 for its transition. - Methyl orange changes color between pH 3.1 to 4.4. - Phenolphthalein changes color between pH 8.2 to 10. - Fluorescein changes around pH 4.5 to 6.5. Analyzing these, methyl orange changes at the desired pH range of around 3.1 to 4.4, which is within the equivalence point for this titration.
04

Select the Indicator

The best indicator for the titration of a weak base with a strong acid is methyl orange since its color change occurs in the desired pH range around the equivalence point of the titration.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Weak Base
A weak base is a substance that partially dissociates in water to produce hydroxide ions (OH鈦�). Unlike strong bases, which dissociate completely, weak bases only partially break apart into their respective ions. This means that at equilibrium, both the base and its ions are present in solution.
Common examples of weak bases include ammonia (NH鈧�) and certain organic compounds like amines. Because they do not entirely disassociate, weak bases create a solution with mildly higher pH levels compared to strong bases.
  • Understanding the concept of weak bases is crucial in titration because the concentration of hydroxide ions varies and affects the overall pH of the solution.
  • In a titration setup, knowing that one is dealing with a weak base helps predict that the reaction will not severely spike the pH level immediately.
Strong Acid
A strong acid is characterized by its complete dissociation in water, producing hydrogen ions (H鈦�). When a strong acid dissolves, it separates entirely into its ions, leading to a high concentration of hydrogen ions in solution. This results in a low pH value, typically below 3.
Some of the most common strong acids include hydrochloric acid (HCl), sulfuric acid (H鈧係O鈧�), and nitric acid (HNO鈧�). These acids exhibit high conductivity due to the abundance of free ions.
  • In a titration involving a strong acid, the changes in pH tend to be more abrupt than with weaker acids, especially near the equivalence point.
  • It鈥檚 important to understand the properties of strong acids, as this knowledge helps in predicting the behavior of a titration and choosing the appropriate indicator.
Indicators
Indicators are substances used in titrations to signal the end point of the reaction by changing color. They work by exhibiting a noticeable color change at a particular pH level and are essential tools in determining the equivalence point of a titration.
The choice of an indicator depends greatly on the expected pH range of the equivalence point. For weak base-strong acid titrations, the equivalence point falls within an acidic pH range. This specificity is why certain indicators are more suitable than others.
  • Methyl orange is an ideal indicator for these titrations because its pH range (3.1 to 4.4) aligns well with the expected equivalence point range (4 to 6).
  • On the other hand, indicators like phenolphthalein and thymol blue are not suitable, as their color change intervals do not match the optimal pH range for a weak base-strong acid titration.
Equivalence Point
The equivalence point in a titration is the moment when the amount of titrant added equals the amount of substance present in the sample. For weak base-strong acid titrations, the equivalence point occurs when all of the weak base has reacted with the strong acid.
Since weak bases do not fully dissociate and strong acids completely dissociate, the resulting solution at the equivalence point is typically acidic. The pH at the equivalence point for these reactions generally falls between 4 and 6.
  • Identifying the equivalence point is crucial, as it allows for accurate determination of the concentration of the unknown solution.
  • The target is to choose an indicator that will change color exactly or very close to this pH, ensuring precise titration results.

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Most popular questions from this chapter

When \(\mathrm{H}_{2} \mathrm{~S}\) is passed through an aqueous solution of an equilimolar mixture of \(\mathrm{Zn}^{2+}\) and \(\mathrm{Pb}^{2+}\) acidified with dilute acetic acid, \(\mathrm{ZnS}\) is not precipitated, because (a) \(\mathrm{K}_{\mathrm{s}}(\mathrm{ZnS})<\mathrm{K}_{\mathrm{si}}(\mathrm{PbS})\) (b) \(\mathrm{K}_{\mathrm{ss}}(\mathrm{ZnS})>\mathrm{K}_{s p}(\mathrm{PbS})\) (c) \(\mathrm{H}_{2} \mathrm{~S}\) decreases the \(\mathrm{K}_{\text {sp }}\) of \(\mathrm{ZnS}\) (d) \(\mathrm{H}_{2} \mathrm{~S}\) increases the \(\mathrm{K}_{\text {p }}\) of \(\mathrm{PbS}\)

Which of the following are the correct statements (a) The \(\mathrm{pH}\) of blood is same in summer and winter (b) \(\mathrm{pH}\) of an acidic buffer increases if more salt is added (c) \(\mathrm{pH}\) of a basic buffer decreases if more salt is added (d) The term solubility product is only for sparingly soluble salts

If \(\mathrm{Ksp}\) of \(\mathrm{Al}(\mathrm{OH})_{3}\) is \(1.0 \times 10^{-15} \cdot \mathrm{M}\). Find at what \(\mathrm{pH}\) does \(1.0 \times 10^{-3} \cdot \mathrm{M} \mathrm{Al}^{3+}\) precipitate on the addition of buffer of \(\mathrm{NH}_{4} \mathrm{Cl}\) and \(\mathrm{NH}_{4} \mathrm{OH}\) solution. (a) 10 (b) \(10.5\) (c) 11 (d) 12

A buffer solution is prepared by mixing \(10 \mathrm{~mL}\) of \(1.0\) M acetic and \(20 \mathrm{~mL}\) of \(0.5 \mathrm{M}\) sodium acetate and then diluted to \(100 \mathrm{~mL}\) with distilled water. If the pKa of \(\mathrm{CH}_{3} \mathrm{COOH}\) is \(4.76\), what is the \(\mathrm{pH}\) of the buffer solution prepared? (a) \(5.21\) (b) \(4.76\) (c) \(4.34\) (d) \(5.35\)

At what concentration of \(\mathrm{CH}_{3} \mathrm{COOH}\) will the \(\left[\mathrm{H}^{+}\right]\) obtained will be same as that obtained from \(10^{-2} \mathrm{M}\) \(\mathrm{HCOOH},\left(\mathrm{Ka}\left(\mathrm{CH}_{3} \mathrm{COOH}\right)=10^{-5}, \mathrm{Ka}(\mathrm{HCOOH})=10^{-4}\right)\) (a) \(10 \mathrm{M}\) (b) \(5 \mathrm{M}\) (c) \(10^{-1} \mathrm{M}\) (d) \(6 \mathrm{M}\)
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