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Chapter 16: Problem 40

Which of the following solutions has the highest pH? (a) \(0.40 \mathrm{M} \mathrm{HCOOH}\) (b) \(0.40 \mathrm{M} \mathrm{HClO}_{4}\) (c) \(0.40 M\) \(\mathrm{CH}_{3} \mathrm{COOH}\)

Short Answer

Expert verified
(c) 0.40 M CH鈧僀OOH has the highest pH.

Step by step solution

01

Identify the Nature of Each Acid

We need to determine the type of acid for each option. - HCOOH (formic acid) and CH鈧僀OOH (acetic acid) are weak acids. - HClO鈧� (perchloric acid) is a strong acid.
02

Calculate or Determine the pH of Each Acid Solution

Generally, strong acids completely dissociate in solution, yielding a lower pH than weak acids with the same concentration, which only partially dissociate. - The pH of a strong acid like HClO鈧� with a concentration of 0.40 M is very low (less than 1). - Weak acids such as HCOOH and CH鈧僀OOH partially dissociate, resulting in a higher pH relative to their concentration. These will not be as low as those of strong acids.
03

Compare pHs of Weak Acids

Weak acids have varying strengths, usually expressed by their dissociation constants (Ka). - Formic acid (HCOOH) has a higher Ka than acetic acid (CH鈧僀OOH), leading to a more significant dissociation and, therefore, a lower pH compared to acetic acid of the same concentration. - Since CH鈧僀OOH dissociates less, it will have the highest pH among the weak acids.
04

Determine Which Solution Has the Highest pH

Combine results from Steps 2 and 3. HClO鈧� has the lowest pH because it's a strong acid. Among weak acids, CH鈧僀OOH has a higher pH than HCOOH. Therefore, 0.40 M CH鈧僀OOH solution has the highest pH.

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Key Concepts

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

Strong Acids
Strong acids are a category of acids that fully dissociate in water. This means that when you dissolve them in solution, they completely break apart into their ions. This makes strong acids very good at producing hydrogen ions (H鈦�), which leads to a lower pH. For example, perchloric acid (HClO鈧�) is considered a strong acid.
Here are some key points about strong acids:
  • They completely dissociate in water, resulting in a high concentration of hydrogen ions.
  • They typically have a pH less than 1 for concentrated solutions.
  • The complete dissociation contributes to their corrosive nature.
If you encounter a solution of a strong acid, you can expect it to significantly affect the pH level.
Weak Acids
Unlike strong acids, weak acids only partially dissociate in water. This means that not all of the acid molecules release their hydrogen ions, leading to fewer free hydrogen ions in solution compared to strong acids. Formic acid (HCOOH) and acetic acid (CH鈧僀OOH) are perfect examples of weak acids, especially when comparing the same molarity solutions.
Some important things to know about weak acids:
  • They dissociate only partially, so they have a higher pH than strong acids of the same concentration.
  • The extent of dissociation is determined by their acid dissociation constant (Ka).
  • Since they dissociate less, they're less corrosive compared to strong acids.
Understanding weak acids will help you comprehend why their solutions are less acidic compared to those of strong acids with equivalent concentrations.
Acid Dissociation Constant (Ka)
The acid dissociation constant (Ka) helps to quantify the strength of a weak acid. It gives us an idea about the extent to which an acid can dissociate in solution.
A larger Ka value means the acid more readily donates its hydrogen ions, indicating a stronger acid among weak acids.
Here is more about the significance of Ka:
  • Ka is derived from the equilibrium constant for the dissociation of an acid.
  • The formula used is: \[ ext{Ka} = \frac{[ ext{H}^+][ ext{A}^-]}{[ ext{HA}]}\]where \([ ext{H}^+]\) is the concentration of hydrogen ions, \([ ext{A}^-]\) is the concentration of the conjugate base, and \([ ext{HA}]\) is the concentration of the un-dissociated acid.
  • A higher Ka value implies a lower pH for the acid solution, as more hydrogen ions are present in the solution.
Using Ka, you can compare the relative strengths of weak acids, allowing you to predict their behavior in solutions.
Acidic Solutions
Acidic solutions are characterized by the presence of excess hydrogen ions. They have a pH value less than 7, with the pH decreasing as the concentration of hydrogen ions increases. The pH scale is logarithmic, so even small changes in the number of hydrogen ions can greatly affect the pH.
Understanding acidic solutions involves a few important points:
  • Strong acids in solution have lower pHs due to complete dissociation, while weak acids exhibit higher pH due to partial dissociation.
  • The strength of an acid in solution is determined by its dissociation behavior, expressed by Ka for weak acids.
  • Solutions with strong acids can be highly reactive and corrosive, demanding proper handling and safety precautions.
By analyzing acidic solutions, you can predict their reactivity and effect on pH, which is crucial for many chemical reactions.

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Most popular questions from this chapter

All Bronsted acids are Lewis acids, but the reverse is not true. Give two examples of Lewis acids that are not Bronsted acids.

(a) Calculate the percent ionization of a \(0.20 \mathrm{M}\) solution of the monoprotic acetylsalicylic acid (aspirin). \(\left(K_{\mathrm{a}}=3.0 \times 10^{-4} .\right)\). (b) The \(\mathrm{pH}\) of gastric juice in the stomach of a certain individual is \(1.00 .\) After a few aspirin tablets have been swallowed, the concentration of acetylsalicylic acid in the stomach is \(0.20 \mathrm{M}\). Calculate the percent ionization of the acid under these conditions.

You are given two beakers containing separately an aqueous solution of strong acid (HA) and an aqueous solution of weak acid (HB) of the same concentration. Describe how you would compare the strengths of these two acids by (a) measuring the \(\mathrm{pH},\) (b) measuring electrical conductance, (c) studying the rate of hydrogen gas evolution when these solutions are reacted with an active metal such as \(\mathrm{Mg}\) or \(\mathrm{Zn}\).

Like water, ammonia undergoes autoionization in liquid ammonia: $$ \mathrm{NH}_{3}+ \mathrm{NH}_{3} \rightleftharpoons \mathrm{NH}_{4}^{+}+ \mathrm{NH}_{2}^{-} $$ (a) Identify the Br?nsted acids and Br贸nsted bases in this reaction. (b) What species correspond to \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-},\) and what is the condition for a neutral solution?

Classify each of these following species as a Lewis acid or a Lewis base: (a) \(\mathrm{CO}_{2},\) (b) \(\mathrm{H}_{2} \mathrm{O},\) (c) \(\mathrm{I}^{-},\) (d) \(\mathrm{SO}_{2}\), (e) \(\mathrm{NH}_{3},\) (f) \(\mathrm{OH}^{-},\) (g) \(\mathrm{H}^{+}\), (h) \(\mathrm{BCl}_{3}\)
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