Problem 34 - General Chemistry: Principles and Modern Applications

Question

Solution (a) is \(100.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{HCl}\) and solution (b) is \(150.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\). A few drops of thymol blue indicator are added to each solution. What is the color of each solution? What is the color of the solution obtained when these two solutions are mixed?

Step-by-Step Solution

Answer

Solution (a) will turn the thymol blue red. Solution (b) will turn the thymol blue blue. When the two solutions are mixed, the thymol blue will turn yellow/green.

Step 1: Identify the nature of solutions

First identify the nature of solutions. Here, solution (a) is hydrochloric acid (HCl), therefore it's acidic. Solution (b) is sodium acetate (NaCH3COO), a salt of a strong base and weak acid, so it is slightly basic.

Step 2: Determine the color of the indicator in each solution

Thymol blue changes color based on the pH of the solution it is in. In acidic solutions (pH < 2.8), it is red. In neutral solutions (pH = 7), it is yellow/green and in basic solutions (pH > 8.2) it is blue. So, solution (a) will turn the thymol blue red because it's acidic. Solution (b), being slightly basic, will turn the thymol blue blue.

Step 3: Determine the color of indicator when solutions are mixed

When these two solutions are mixed, they will neutralize each other as an acid reacts with a base to make a salt and water. The pH of the resulting solution will be about 7, which is neutral. So, upon mixing, the thymol blue will become yellow/green.

Key Concepts

Thymol Blue Indicator

Thymol blue is a pH indicator used in chemistry to determine the acidity or basicity of a solution. It changes color at different pH levels, allowing scientists to estimate the pH of the solution visually.

Thymol blue is particularly useful across certain pH ranges with specific color transitions:

• In strongly acidic solutions with a pH below 2.8, thymol blue appears red.
• In neutral solutions, roughly around a pH of 7, it shifts to a yellow/green color.
• In basic environments with a pH above 8.2, thymol blue takes on a blue hue.

This visual guide can be extremely helpful when conducting experiments, as it provides a quick and simple way to gauge the pH without using electronic equipment.

It's important to note that thymol blue is one of many pH indicators available, each with its own range and color changes. Selecting the right indicator depends on the expected pH change in the experiment.

pH Scale

The pH scale is a numerical scale that ranges from 0 to 14 and is used to specify how acidic or basic a water-based solution is. "pH" stands for "potential of hydrogen" and reflects the concentration of hydrogen ions in the solution.

Here's a simple breakdown of the scale:

• Solutions with a pH less than 7 are considered acidic. The lower the pH, the stronger the acid.
• A pH of exactly 7 is neutral, meaning the solution is neither acidic nor basic. Pure water is a common example of a neutral solution.
• Solutions with a pH greater than 7 are basic (also known as alkaline). The higher the pH, the stronger the base.

The pH scale is logarithmic, meaning each whole number increase on the scale represents a tenfold increase in basicity, and conversely, each decrease represents a tenfold increase in acidity.

The pH value is a crucial element in chemistry because it influences the behavior of molecules, the rate of chemical reactions, and biological processes.

Neutralization Reaction

A neutralization reaction is a chemical reaction that occurs when an acid and a base interact to form a salt and water. This type of reaction is critical in maintaining pH balance in various chemical processes.

During a neutralization reaction, the hydrogen ions (\( ext{H}^+ \)) from the acid react with hydroxide ions (\( ext{OH}^- \)) from the base, resulting in the formation of water (\( ext{H}_2 ext{O} \)). The general equation for a neutralization reaction can be written as:
\[ ext{Acid} + ext{Base} ightarrow ext{Salt} + ext{Water} \]
In the specific case of the original exercise with hydrochloric acid (\( ext{HCl} \)) and sodium acetate (\( ext{NaCH}_3 ext{COO} \)), they react to produce sodium chloride (\( ext{NaCl} \)) and acetic acid (\( ext{CH}_3 ext{COOH} \)). However, what is notable in solution is the dominant reaction between the hydrogen and hydroxide ions forming water, leading to a neutral pH level of around 7.

Neutralization reactions are not just laboratory exercises but are fundamental to many real-world applications, such as antacid tablets neutralizing stomach acid or wastewater treatment processes adjusting industrial effluent pH levels.