Question

Phenol red indicator changes from yellow to red in the pH range from 6.6 to \(8.0 .\) Without making detailed calculations, state what color the indicator will assume in each of the following solutions: (a) \(0.10 \mathrm{M} \mathrm{KOH}\) (b) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH} ;\) (c) \(0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{NO}_{3} ;\) (d) \(0.10 \mathrm{M}\) HBr; (e) \(0.10 \mathrm{M} \mathrm{NaCN} ;\) (f) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}-0.10 \mathrm{M}\) \(\mathrm{NaCH}_{3} \mathrm{COO}\).

Short answer

The colors phenol red would change to in each solution are: (a) red, (b) yellow, (c) orange, (d) yellow, (e) red, and (f) yellow.

Step-by-step solution

Significance of KOH

KOH is a strong base. Therefore, its solution will have a pH greater than 7, and specifically for this exercise, it will bring the pH over 8.0. Therefore, in a \(0.10 M KOH\) solution, the phenol red indicator will turn red.

Significance of CH3COOH

CH3COOH is a weak acid, which means that it will not fully ionize in solution, thereby resulting in a solution with a pH of less than 7, and specifically for this exercise, a pH below 6.6. Therefore, in a \(0.10 M CH3COOH\) solution, the phenol red indicator will turn yellow.

Significance of NH4NO3

NH4NO3, a salt of a weak base (NH4) and a strong acid (NO3), results in a neutral solution. The phenol red indicator will neither be yellow nor red but will be orange, which is the indicator color at neutral pH.

Significance of HBr

HBr is a strong acid. When in solution, it will cause the pH to fall below 6.6. Thus, in a \(0.10 M HBr\) solution, the phenol red indicator will turn yellow.

Significance of NaCN

NaCN, being a salt of a weak acid (HCN) and a strong base (NaOH), will create a weakly basic solution. This means the pH will be over 8.0. Therefore, adding the phenol red indicator to a \(0.10 M NaCN\) solution will result in a red color.

Significance of CH3COOH-NaCH3COO

The solution of \(0.10 M CH3COOH-0.10 M NaCH3COO\) is a buffer solution with a pH less than 7 and, for this exercise, a pH below 6.6. This will result in the phenol red indicator turning yellow.

Key concepts

Acid-Base Chemistry

Acid-base chemistry is a fundamental concept in chemistry that helps us understand the properties and reactions of acids and bases in aqueous solutions. Acids are substances that release hydrogen ions (\( \text{H}^+ \)) in solution, while bases release hydroxide ions (\( \text{OH}^- \)). The pH scale, which ranges from 0 to 14, measures the acidity or basicity of a solution. A pH below 7 indicates an acidic solution, a pH of 7 is neutral, and a pH above 7 is basic or alkaline.
When an acid reacts with a base, they can neutralize each other, forming water and a salt. This is called neutralization. Different indicators, such as phenol red, change color depending on the acidity or basicity of the solution. Phenol red turns from yellow to red as the pH increases from 6.6 to 8.0.
Understanding acid-base chemistry is crucial for predicting the behavior of substances in various solutions, such as those mentioned in the original exercise.

Buffer Solutions

Buffer solutions are special systems in acid-base chemistry that resist changes in their pH when small amounts of acid or base are added. Often consisting of a weak acid and its conjugate base, or a weak base and its conjugate acid, buffer solutions maintain a steady pH level.
An example of this is the solution of acetic acid (\( \text{CH}_3\text{COOH} \)) and sodium acetate (\( \text{NaCH}_3\text{COO} \)) mentioned in the exercise. This combination acts as a buffer, stabilizing the pH around a certain value due to the presence of both the weak acid and its conjugate base. When an external acid or base is added to the buffer, the weak acid/base pair reacts to neutralize the additional ions, preventing a significant shift in pH.
Buffer solutions are essential in many biological and chemical processes, allowing for control and stability in chemical reactions and maintaining conditions necessary for systems to function properly.

Weak Acids and Bases

Weak acids and bases only partially ionize in solution, which affects their behavior in acid-base reactions. Weak acids, like acetic acid (\( \text{CH}_3\text{COOH} \)), do not fully dissociate into hydrogen ions and their conjugate base. This means they produce fewer hydrogen ions, resulting in a higher pH compared to strong acids of the same concentration.
Similarly, weak bases do not completely dissociate to absorb hydrogen ions or produce hydroxide ions (\( \text{OH}^- \)). An example is ammonia (\( \text{NH}_3 \)), which is a weak base that reacts partially with water to form ammonium ions (\( \text{NH}_4^+ \)) and hydroxide ions.
Understanding the behavior of weak acids and bases is crucial, as their incomplete ionization influences the pH of solutions and impacts their interaction with indicators like phenol red. In the original exercise, weak acids and bases contribute to the color changes noted with different solutions, guiding the prediction of their properties.

Strong Acids and Bases

Strong acids and bases are characterized by their complete ionization in aqueous solutions. Unlike their weak counterparts, strong acids dissociate fully, releasing a high concentration of hydrogen ions (\( \text{H}^+ \)) into the solution. An example is hydrochloric acid (\( \text{HCl} \)), which when dissolved in water, completely ionizes to form hydrogen and chloride ions.
In contrast, strong bases like potassium hydroxide (\( \text{KOH} \)), fully dissociate to provide hydroxide ions (\( \text{OH}^- \)), resulting in a significantly high pH. These substances can greatly affect the pH of a solution and consequently, the color change of pH indicators like phenol red.
In the context of the exercise, the solution of strong acids and bases resulted in an immediate and predictable pH change. Understanding the dissociation behavior of strong acids and bases allows for accurate predictions and manipulations of pH in various chemical contexts.