Introduction to Acids and Bases

Learning Objectives

In the topic “Introduction to Acids and Bases” for the AP Chemistry exam, you should learn to define and distinguish between acids and bases using the Arrhenius, Brønsted-Lowry, and Lewis theories. Understand the properties and behaviors of strong and weak acids and bases, and calculate the pH and pOH of solutions. Master the concept of acid-base titrations, including identifying the equivalence point and using indicators. Recognize the composition and function of buffer solutions, and apply your knowledge to solve practice problems involving acid and base concentrations, dissociation constants, and pH calculations. These objectives will prepare you to analyze and interpret acid-base reactions effectively.

Introduction

Acids and bases are fundamental concepts in chemistry, playing a crucial role in various chemical reactions and processes. Acids, characterized by their sour taste and ability to turn blue litmus paper red, release hydrogen ions (H⁺) in solution, while bases, known for their bitter taste and slippery feel, release hydroxide ions (OH⁻). The pH scale, ranging from 0 to 14, measures the acidity or basicity of a solution, with values below 7 indicating acidity and values above 7 indicating basicity. Understanding acids and bases is essential for exploring topics like neutralization reactions, acid-base titrations, and the behavior of substances in different environments.

What are Acids and Bases?

Acids: Acids are substances that can donate a proton (H⁺ ion) to another substance. They have a pH less than 7 in aqueous solutions and can turn blue litmus paper red.

Bases: Bases are substances that can accept a proton (H⁺ ion) or donate a hydroxide ion (OH⁻) to another substance. They have a pH greater than 7 in aqueous solutions and can turn red litmus paper blue.

Properties of Acids and Bases

Properties of Acids

• Sour taste
• Conduct electricity in aqueous solutions
• React with metals to produce hydrogen gas
• Turn blue litmus paper red
• pH less than 7

Properties of Bases

• Bitter taste
• Slippery texture
• Conduct electricity in aqueous solutions
• Neutralize acids to form water and a salt
• Turn red litmus paper blue
• pH greater than 7

Theories of Acids and Bases

Arrhenius Theory

• Acid: A substance that increases the concentration of hydrogen ions (H⁺) in an aqueous solution.
  • Example: Hydrochloric acid (HCl) dissociates in water to form H⁺ and Cl⁻ ions.
• Base: A substance that increases the concentration of hydroxide ions (OH⁻) in an aqueous solution.
  • Example: Sodium hydroxide (NaOH) dissociates in water to form Na⁺ and OH⁻ ions.

Brønsted-Lowry Theory

• Acid: A proton (H⁺ ion) donor.
  • Example: Ammonium ion (NH₄⁺) donates a proton to form ammonia (NH₃).
• Base: A proton (H⁺ ion) acceptor.
  • Example: Ammonia (NH₃) accepts a proton to form ammonium ion (NH₄⁺).

Lewis Theory

• Acid: An electron pair acceptor.
  • Example: Boron trifluoride (BF₃) accepts an electron pair from ammonia (NH₃) to form a coordinate covalent bond.
• Base: An electron pair donor.
  • Example: Ammonia (NH₃) donates an electron pair to boron trifluoride (BF₃) to form a coordinate covalent bond.

Strength of Acids and Bases

Strong Acids

Definition: Strong acids completely dissociate into their ions in water.

Examples:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)

Weak Acids

Definition: Weak acids partially dissociate in water, establishing an equilibrium between the undissociated acid and the ions.

Examples:

• Acetic acid (CH₃COOH)
• Formic acid (HCOOH)

Strong Bases

Definition: Strong bases completely dissociate into their ions in water.

Examples:

• Sodium hydroxide (NaOH)
• Potassium hydroxide (KOH)

Weak Bases

Definition: Weak bases partially dissociate in water, establishing an equilibrium between the undissociated base and the ions.

Examples:

• Ammonia (NH₃)
• Methylamine (CH₃NH₂)

pH and pOH

pH

Definition: pH is a measure of the hydrogen ion concentration [H⁺] in a solution. It indicates how acidic or basic a solution is.

Formula: pH = −log⁡[H⁺]

Range:

• Acidic: pH < 7
• Neutral: pH = 7
• Basic (Alkaline): pH > 7

pOH

Definition: pOH is a measure of the hydroxide ion concentration [OH⁻] in a solution. It is related to the basicity of a solution.

Formula: pOH = −log⁡[OH⁻]

Relationship to pH:

pH + pOH = 14 (at 25°C)

Acid-Base Titrations

Acid-base titration is a quantitative analytical method used to determine the concentration of an acid or a base by reacting it with a standard solution of known concentration.

Key Concepts

1. Titrant: The solution of known concentration, usually a strong acid or base, added from a burette.
2. Analyte: The solution of unknown concentration that reacts with the titrant.
3. Indicator: A substance that changes color at the endpoint of the titration.
4. Endpoint: The point in the titration where the indicator changes color, signaling that the reaction is complete.
5. Equivalence Point: The point at which the amount of titrant added is stoichiometrically equivalent to the amount of analyte in the solution. Ideally, this is the same as the endpoint.

Procedure

1. Preparation: Fill a burette with the titrant of known concentration.
2. Indicator Addition: Add a few drops of an appropriate indicator to the analyte solution.
3. Titration: Slowly add the titrant to the analyte while continuously stirring until the indicator changes color.
4. Calculation: Use the volume of titrant added to calculate the concentration of the analyte.

Indicators

Common indicators and their pH range for color change:

• Phenolphthalein: Colorless in acidic solutions, pink in basic solutions (pH 8.2 – 10.0).
• Methyl Orange: Red in acidic solutions, yellow in basic solutions (pH 3.1 – 4.4).
• Bromothymol Blue: Yellow in acidic solutions, blue in basic solutions (pH 6.0 – 7.6).

Types of Titrations

1. Strong Acid-Strong Base: Sharp equivalence point; any suitable indicator can be used.
2. Weak Acid-Strong Base: Equivalence point above pH 7; phenolphthalein is suitable.
3. Strong Acid-Weak Base: Equivalence point below pH 7; methyl orange is suitable.
4. Weak Acid-Weak Base: Less sharp equivalence point; requires more careful selection of indicators.

Buffer Solutions

A buffer solution is a solution that resists significant changes in pH when small amounts of an acid or a base are added. Buffers are essential in maintaining stable pH levels in various chemical and biological systems.

Components of a Buffer

1. Weak Acid and Its Conjugate Base
   • Example: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa)
2. Weak Base and Its Conjugate Acid
   • Example: Ammonia (NH₃) and ammonium chloride (NH₄Cl)

How Buffers Work

Buffers work through the principle of Le Chatelier’s Principle. When an acid (H⁺) or base (OH⁻) is added to the buffer, the equilibrium shifts to minimize the change in pH.

Acid Addition

• Weak Acid Buffer: Added H⁺ ions react with the conjugate base (A⁻) to form the weak acid (HA), thus reducing the impact on pH. H⁺+A⁻→HA\text{H⁺} + \text{A⁻} \rightarrow \text{HA}H⁺+A⁻→HA

Base Addition

• Weak Acid Buffer: Added OH⁻ ions react with the weak acid (HA) to form water and the conjugate base (A⁻), thus reducing the impact on pH. OH⁻ + HA → A⁻ + H₂O

Buffer Capacity

Buffer capacity is the measure of a buffer’s ability to resist changes in pH. It depends on the concentrations of the weak acid and its conjugate base (or weak base and its conjugate acid).

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentration of the acid and its conjugate base:

\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)

Where:

• pKₐ​ is the negative logarithm of the acid dissociation constant Kₐ.
• [A⁻] is the concentration of the conjugate base.
• [HA]is the concentration of the weak acid.

Common Buffer Systems

1. Acetic Acid-Acetate Buffer:
   • Components: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa)
   • pH Range: 3.7 – 5.7
2. Ammonium Buffer:
   • Components: Ammonia (NH₃) and ammonium chloride (NH₄Cl)
   • pH Range: 9.2 – 11.2
3. Phosphate Buffer:
   • Components: Dihydrogen phosphate (H₂PO₄⁻) and hydrogen phosphate (HPO₄²⁻)
   • pH Range: 5.8 – 8.0

Examples of Acids and Bases

Examples of Acids

1. Phosphoric Acid (H₃PO₄)
2. Formic Acid (HCOOH)
3. Lactic Acid (C₃H₆O₃)
4. Carbonic Acid (H₂CO₃)
5. Hydrofluoric Acid (HF)

Examples of Bases

1. Barium Hydroxide (Ba(OH)₂)
2. Lithium Hydroxide (LiOH)
3. Sodium Bicarbonate (NaHCO₃)
4. Zinc Hydroxide (Zn(OH)₂)
5. Triethanolamine (C₆H₁₅NO₃)