General Chemistry Study Guide: Acids, Bases, Buffers, and Solubility

Chapter 4: Acids and Bases

Acid-Base Equilibria

This section covers the fundamental concepts of acids, bases, and their equilibria in aqueous solutions. Understanding these principles is essential for predicting the behavior of acids and bases in chemical reactions.

• Equilibrium Constant Expression for Weak Acids: The equilibrium constant for the dissociation of a weak acid (Ka) is given by: $K_a = \frac{[H^+][A^-]}{[HA]}$ where [HA] is the concentration of the weak acid, [H+] is the concentration of hydrogen ions, and [A-] is the concentration of the conjugate base.

• Conjugate Acid-Base Pairs: A conjugate acid-base pair consists of two species that differ by a single proton (H+).

• Identifying Acids and Bases: In a reaction, the acid donates a proton, and the base accepts a proton.

• pH and pOH Calculations: The pH of a solution is calculated as $pH = -\log[H^+]$. The pOH is $pOH = -\log[OH^-]$. For aqueous solutions at 25°C, $pH + pOH = 14$.

• Percent Ionization: Percent ionization of a weak acid is calculated as: $\text{Percent Ionization} = \frac{[H^+]_{eq}}{[HA]_{initial}} \times 100\%$

• Ka and Kb Relationship: For a conjugate acid-base pair, $K_a \times K_b = K_w$, where $K_w = 1.0 \times 10^{-14}$ at 25°C.

• Polyprotic Acids: Acids that can donate more than one proton (e.g., H2SO4, H3PO4).

• Relative Strengths: The strength of acids and bases can be compared using their Ka and Kb values. Larger Ka or Kb values indicate stronger acids or bases, respectively.

• Brønsted-Lowry vs. Lewis Definitions: The Brønsted-Lowry definition focuses on proton transfer, while the Lewis definition focuses on electron pair donation/acceptance.

Examples and Applications

• Example: Calculate the pH of a 0.10 M acetic acid solution (Ka = 1.8 × 10-5).

• Example: Identify the conjugate base of HCO3- (answer: CO32-).

Chapter 5: Buffers and Titrations

Buffer Solutions

Buffers are solutions that resist changes in pH upon the addition of small amounts of acid or base. They are typically composed of a weak acid and its conjugate base or a weak base and its conjugate acid.

• Henderson-Hasselbalch Equation: Used to calculate the pH of a buffer solution: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Buffer Capacity: The amount of acid or base a buffer can neutralize before the pH changes significantly.

• Buffer Preparation: Buffers are prepared by mixing a weak acid with its salt (conjugate base) or a weak base with its salt (conjugate acid).

Titrations

• Titration Curves: Graphs showing pH changes as a function of added titrant volume.

• Equivalence Point: The point at which the amount of acid equals the amount of base during a titration.

• Indicator Selection: Choose an indicator whose color change occurs near the equivalence point pH.

• Net Ionic Equations: Show only the species that undergo chemical change during the reaction.

Examples and Applications

• Example: Calculate the pH at the equivalence point for the titration of 0.10 M acetic acid with 0.10 M NaOH.

• Example: Write the net ionic equation for the reaction between HCl and NaOH.

Chapter 6: Solubility and Precipitation

Solubility Product (Ksp)

The solubility product constant (Ksp) describes the equilibrium between a solid and its ions in a saturated solution. It is used to predict whether a precipitate will form when solutions are mixed.

• Writing Ksp Expressions: For a salt $AB_2$ that dissociates as $AB_2 \rightarrow A^{2+} + 2B^-$, the Ksp is: $K_{sp} = [A^{2+}][B^-]^2$

• Calculating Solubility: Use the Ksp value to determine the molar solubility of a salt in water.

• Common Ion Effect: The solubility of a salt decreases in the presence of a common ion.

• Ion Product (Q): The reaction quotient for a solubility equilibrium. If Q > Ksp, a precipitate forms; if Q < Ksp, no precipitate forms; if Q = Ksp, the solution is saturated.

• Selective Precipitation: A technique used to separate ions by adding a reagent that precipitates one ion but not others.

• Qualitative Analysis: The identification of ions in a mixture based on their chemical reactions, often involving precipitation.

Examples and Applications

• Example: Given the Ksp of AgCl is $1.8 \times 10^{-10}$, calculate the solubility of AgCl in water.

• Example: Predict whether a precipitate will form when 0.01 M NaCl is mixed with 0.01 M AgNO3.

Table: Comparison of Acid-Base and Solubility Concepts

Concept	Key Equation	Application
Acid Dissociation (Ka)	$K_a = \frac{[H^+][A^-]}{[HA]}$	Calculating pH of weak acid solutions
Base Dissociation (Kb)	$K_b = \frac{[BH^+][OH^-]}{[B]}$	Calculating pOH of weak base solutions
Buffer pH	$pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$	Determining pH of buffer solutions
Solubility Product (Ksp)	$K_{sp} = [Cation]^m[Anion]^n$	Predicting precipitation and solubility