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Chapter 4: Problem 85

Describe the basic steps involved in an acid-base titration. Why is this technique of great practical value?

Short Answer

Expert verified
An acid-base titration involves the setup of apparatus, initial measurements, gradual addition of titrant to the analyte with measurements, endpoint determination and calculation of the unknown concentration. It is of great importance in industries like chemical, pharmaceutical, and water treatment for determining the concentration of substances, and in preparing solutions with desired pH.

Step by step solution

01

Setup

Acid-base titrations typically include the use of a burette containing the titrant (a solution with known concentration), and a flask containing the analyte (solution of unknown concentration). An appropriate indicator is also part of the initial setup, it helps to determine the endpoint of the titration.
02

Initial Measurements

Record the initial volume of titrant in the burette. Also, if applicable, measure the initial pH of the analyte solution.
03

Titrant Addition and Measurements

Add the titrant to the analyte slowly, while constantly stirring to ensure the reaction proceeds properly. After each addition, measure and record the volume of the titrant delivered and the pH or observed indicator change of the analyte solution.
04

Endpoint Determination

Continue adding titrant until the physical or chemical indicator signals that the endpoint of the titration has been reached. This is usually a sudden change in pH, or a color change in the analyte solution.
05

Calculations

Use the balanced chemical equation for the reaction, the volume of titrant delivered, and the molarity of the titrant to calculate the concentration of the analyte.
06

Importance of Acid-Base Titration

Acid-base titration is a crucial process in chemistry. It is used in a wide range of practical applications, such as in the chemical, pharmaceutical, food, and water treatment industries to determine the concentration of different substances. It also helps in the preparation of solutions with desired pH and in the analysis of chemical reactions and their mechanisms.

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Most popular questions from this chapter

Nitric acid is a strong oxidizing agent. State which of the following species is least likely to be produced when nitric acid reacts with a strong reducing agent such as zinc metal, and explain why: \(\mathrm{N}_{2} \mathrm{O}\) \(\mathrm{NO}, \mathrm{NO}_{2}, \mathrm{~N}_{2} \mathrm{O}_{4}, \mathrm{~N}_{2} \mathrm{O}_{5}, \mathrm{NH}_{4}^{+}\).

You have \(505 \mathrm{~mL}\) of a \(0.125 \mathrm{M} \mathrm{HCl}\) solution and you want to dilute it to exactly \(0.100 \mathrm{M}\). How much water should you add? Assume volumes are additive.

A 3.664 -g sample of a monoprotic acid was dissolved in water. It took \(20.27 \mathrm{~mL}\) of a \(0.1578 \mathrm{M}\) \(\mathrm{NaOH}\) solution to neutralize the acid. Calculate the molar mass of the acid.

Ammonium nitrate \(\left(\mathrm{NH}_{4} \mathrm{NO}_{3}\right)\) is one of the most important nitrogen-containing fertilizers. Its purity can be analyzed by titrating a solution of \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) with a standard \(\mathrm{NaOH}\) solution. In one experiment a \(0.2041-\mathrm{g}\) sample of industrially prepared \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) required \(24.42 \mathrm{~mL}\) of \(0.1023 \mathrm{M} \mathrm{NaOH}\) for neutralization. (a) Write a net ionic equation for the reaction. (b) What is the percent purity of the sample?

Calculate the concentration (in molarity) of a \(\mathrm{NaOH}\) solution if \(25.0 \mathrm{~mL}\) of the solution are needed to neutralize \(17.4 \mathrm{~mL}\) of a \(0.312 \mathrm{M} \mathrm{HCl}\) solution.
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