Book cover for Chemistry The Science in Context

Chemistry The Science in Context

Thomas R. Gilbert

ISBN #9780393615142

5th Edition

2,675 Questions

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191,124 Students Helped

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Summary
Learning Objectives
Key Concepts
Example Problems
Explanations
Common Mistakes

Summary

This chapter elucidates the complexity of aqueous equilibria, emphasizing the regulation of pH in natural and laboratory settings through buffer systems and the common-ion effect. By integrating the Henderson–Hasselbalch equation, acid–base titration principles, and advanced concepts such as Lewis acid–base theory and complex ion formation, students gain a comprehensive understanding of processes integral to ocean acidification and environmental chemistry. The interplay of these equilibria demonstrates how small changes in ion concentration or external stresses can significantly influence system behavior.

Learning Objectives

1

Explain the principles of additional aqueous equilibria and how they affect ocean chemistry and acidification.

2

Describe the common-ion effect and its impact on solubility equilibria and buffer systems.

3

Apply the Henderson–Hasselbalch equation to design and analyze pH buffers in both environmental and laboratory settings.

4

Understand the mechanics of acid–base titrations, including indicators, titration curves, and the roles of Lewis acids and bases.

5

Analyze the formation of complex ions and the behavior of hydrated metal ions in aqueous solutions.

Key Concepts

CONCEPT

DEFINITION

Additional Aqueous Equilibria

Equilibrium processes involving multiple species in solution, including buffer systems, titrations, and complex ion formation.

Ocean Acidification

The decrease in pH of ocean water primarily due to the absorption of atmospheric CO2, affecting marine life and chemical equilibria.

Common-Ion Effect

The shift in equilibrium that occurs when a compound containing an ion in the equilibrium is added to the solution, reducing solubility or altering pH.

pH Buffers

Solutions that resist changes in pH upon the addition of small amounts of acids or bases, typically analyzed using the Henderson–Hasselbalch equation.

Henderson–Hasselbalch Equation

An equation that relates the pH of a buffer solution to the pKa of the acid and the ratio of the concentrations of the conjugate base and the acid.

Acid–Base Titrations

A quantitative analytical procedure that determines the concentration of an acid or base in solution using a titrant and indicators.

Indicators

Substances that change color at a specific pH range, used to signal the endpoint in titration procedures.

Lewis Acids and Bases

Chemical species that accept (Lewis acids) or donate (Lewis bases) electron pairs, extending acid–base concepts beyond traditional proton transfer.

Complex Ion Formation

The process of formation of a complex ion where central metal ions bind with surrounding ligands, affecting solubility and reactivity in solution.

Solubility Equilibria

The equilibrium between a solid and its ions in a saturated solution, often impacted by the common-ion effect.

Hydrated Metal Ions

Metal ions that are coordinated by water molecules in a solution, which can act as acids due to their ability to donate protons.

Example Problems

Example 1

The graph in Figure P16.1 shows the titration curves of a 1 $M$ solution of a weak acid with a strong base and a $1 M$ solution of a strong acid with the same base. Which curve is which? (FIGURE CAN'T COPY)

Example 2

Estimate (to two significant figures) the $\mathrm{p} K_{\mathrm{a}}$ of the weak acid in Problem 16.1

Example 3

Suppose you have four color indicators to choose from to detect the equivalence point of the titration reaction represented by the red curve in Figure P16.1. The p $K_{\mathrm{a}}$ values of the four indicators are $3.3,5.0,7.0,$ and 9.0 Which indicator would be the best one to choose? (FIGURE CAN'T COPY)

Example 4

Explain why the slope of the red titration curve in Figure $P 16.1$ is nearly flat in the region extending about halfway from the start of the titration to its equivalence point. (FIGURE CAN'T COPY)

Example 5

One of the titration curves in Figure $\mathrm{P} 16.5$ represents the titration of an aqueous sample of $\mathrm{Na}_{2} \mathrm{CO}_{3}$ with strong acid; the other represents the titration of an aqueous sample of $\mathrm{NaHCO}_{3}$ with the same acid. Which curve is which?(FIGURE CAN'T COPY)

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Step-by-Step Explanations

QUESTION

How do you calculate the pH of a buffer solution using the Henderson–Hasselbalch equation?

STEP-BY-STEP ANSWER:

Step 1: Identify the acid (HA) and its conjugate base (A-) present in the buffer solution, and determine their molar concentrations.
Step 2: Obtain or calculate the pKa value for the weak acid from literature or experimental data.
Step 3: Plug the values into the Henderson–Hasselbalch equation: pH = pKa + log([A-]/[HA]).
Step 4: Calculate the log ratio to determine the contribution of the conjugate base to the overall pH.
Step 5: Sum the pKa and the log ratio to arrive at the final pH value.
Final Answer: The pH of the buffer solution is given by pH = pKa + log([A-]/[HA]).

Henderson–Hasselbalch Equation

QUESTION

How does the addition of a common ion affect the solubility of a sparingly soluble salt?

STEP-BY-STEP ANSWER:

Step 1: Write the solubility equilibrium equation for the sparingly soluble salt (e.g., AB ⇌ A+ + B-).
Step 2: Identify the ion that is already present in the solution (the common ion).
Step 3: Apply Le Chatelier's principle: the addition of the common ion will shift the equilibrium to the left, reducing ionization.
Step 4: Conclude that the solubility of the salt decreases due to the increased concentration of the common ion.
Final Answer: The presence of a common ion decreases the solubility of a sparingly soluble salt by shifting the equilibrium to favor the undissolved form.

Common-Ion Effect on Solubility

QUESTION

How can you determine the equivalence point in an acid–base titration using a titration curve?

STEP-BY-STEP ANSWER:

Step 1: Record the pH values as titrant is added to the analyte solution.
Step 2: Plot the titration curve with pH on the y-axis and volume of titrant on the x-axis.
Step 3: Identify the point on the curve where the pH changes most rapidly; this is the steepest part of the curve.
Step 4: The midpoint of this steep change corresponds to the equivalence point of the titration.
Final Answer: The equivalence point is determined by the steepest part of the titration curve, where the pH undergoes a sudden change.

Acid–Base Titration Curve

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Common Mistakes

  • Assuming that buffers can maintain a constant pH regardless of the amount of acid or base added.
  • Neglecting the impact of the common-ion effect on solubility when additional ions are introduced into the system.
  • Incorrectly applying the Henderson–Hasselbalch equation by mixing up the roles of the acid and its conjugate base.
  • Failing to recognize that titration curves must be interpreted carefully, especially near the equivalence point where changes are most pronounced.
  • Overlooking the broader acid–base interactions beyond Brønsted–Lowry definitions by not considering Lewis acid–base behavior.