Lesson 1: Acid-Base Concepts

Video Lesson

Competencies

At the end of the lesson, you will be able to:

• Explain the Arrhenius Theory of Acids and Base
• Understand the Basics of Acids and Bases
• Identify Strong Acids and Bases
• Discuss the Limitations of Arrhenius’s Theory
• Understand the Bronsted-Lowry Concept
• Explore Conjugate Acid-Base Pairs
• Examine the Autoionization of Water
• Identify Amphiprotic Species

Brainstorming Questions

• What are acids and bases ?
• Mention some examples of acid and base which can be found at your home?

Key Terms

1.1. Arrhenius Basic Theory

The Swedish chemist Svante Arrhenius was the first to propose a successful concept of acids and bases in the late 19th century. His theory, which defines acids and bases based on their behavior in water, remains foundational in the study of chemistry despite some limitations. This note elaborates on Arrhenius’s theory, its examples, and its limitations.

1.1.1. Definition of Acids and Bases

According to Arrhenius, acids are substances that increase the concentration of H⁺ ions (protons) when dissolved in water, while bases are substances that increase the concentration of OH⁻ ions (hydroxide ions) in aqueous solution.

Acids:

• When an acid dissolves in water, it releases H⁺ ions.

For example, hydrochloric acid (HCl) dissociates in water to form H⁺ and Cl⁻ ions: HCl → H⁺ + Cl⁻

Bases:

• When a base dissolves in water, it releases OH⁻ ions.

For example, sodium hydroxide (NaOH) dissociates in water to form Na⁺ and OH⁻ ions: NaOH → Na⁺+OH⁻

Strong Acids and Bases

In Arrhenius’s theory, strong acids and bases are those that completely ionize in aqueous solution.

Strong Acids:

• A strong acid fully dissociates in water to produce H_3​O⁺ (hydronium ion) and an anion.

For example, perchloric acid (HClO₄​) dissociates as follows: HClO₄(aq) + H₂O(l) → H₃O⁺(aq) + ClO₄⁻⁽aq)

• Other examples of strong acids include sulfuric acid (H₂SO_4​), hydroiodic acid (HI), hydrobromic acid (HBr), hydrochloric acid (HCl), and nitric acid (HNO₃​).

Strong Bases:

• A strong base fully dissociates in water to produce OH⁻ ions and a cation.

For example, sodium hydroxide (NaOH) dissociates as follows: NaOH(s)→Na⁺(aq)+OH⁻(aq)

• The principal strong bases are the hydroxides of Group IA elements (such as NaOH and KOH) and Group IIA elements (such as Ca(OH)₂and Ba(OH)₂, except for beryllium hydroxide Be(OH)₂.

1.1.2. Limitations of Arrhenius’s Theory

Despite its significant contributions and continued usefulness, Arrhenius’s theory has limitations:

• Non-Aqueous Solutions: The theory does not account for acids and bases in non-aqueous solutions. It is limited to substances that produce H⁺ and OH⁻ ions in water.
• Broader Definitions: The theory does not include substances that do not produce H⁺ or OH⁻ ions but still exhibit acidic or basic properties.

For example, ammonia (NH₃) acts as a base by accepting a proton, but it does not produce OH⁻ ions directly.

• Incomplete Ionization: It does not address weak acids and bases that only partially ionize in water.

1.1.3. Bronsted-Lowry Basic Theory

• In 1923, J. N. Bronsted from Denmark and T. M. Lowry from Great Britain independently proposed a new theory for acids and bases. Their theory, known as the Bronsted-Lowry concept, describes acid-base reactions as proton-transfer reactions, defining acids as proton donors and bases as proton acceptors. This concept extends the definition of acids and bases beyond aqueous solutions and addresses some of the limitations of the Arrhenius theory.

Definition:

Proton Transfer: Acid-base reactions are seen as proton-transfer reactions.

Acid Definition: An acid is a proton donor.

Base Definition: A base is a proton acceptor.

Example: In the ionization of ammonia in water: NH₃+H₂O→NH4⁺+OH ⁻ Water acts as an acid (proton donor), and ammonia acts as a base (proton acceptor).

1.1.4. Bronsted-Lowry and Arrhenius Theories

Scope:

• Arrhenius: Limited to aqueous solutions, defining acids and bases in terms of H⁺ and OH⁻ ions.
• Bronsted-Lowry: Not limited to aqueous solutions, defining acids and bases based on proton transfer.

Conjugate Acid-Base Pairs

• Definition: Conjugate acid-base pairs consist of two species that differ by a single proton.
• Formation:
  • A conjugate base is formed when an acid donates a proton.
  • A conjugate acid is formed when a base accepts a proton.
• Examples:
  • In the reaction of acetic acid CH₃​COOH with water: CH₃COOH+H₂O⇌CH₃COO⁻+H₃O⁺
    • CH₃COOH is the acid.
    • H_2​O is the base.
    • CH₃COO⁻is the conjugate base.
    • H₃​O⁺ is the conjugate acid.

• HCl (acid) and Cl⁻ (conjugate base).
• NH₃​ (base) and NH₄⁺ (conjugate acid).

Strengths of Conjugate Acid-Base Pairs

• How to Determine the Strengths of Conjugate Acid-Base Pairs

The strength of conjugate acid-base pairs can be determined based on their ability to donate or accept protons. The concept relies on the relationship between the strengths of acids and their conjugate bases and the strengths of bases and their conjugate acids.

• Relative Strengths:
  • Stronger Acid, Weaker Conjugate Base: The stronger an acid, the weaker its conjugate base. This means that if an acid readily donates a proton, its conjugate base will have a low tendency to accept a proton.
  • Stronger Base, Weaker Conjugate Acid: Conversely, the stronger a base, the weaker its conjugate acid. A strong base readily accepts a proton, making its conjugate acid less likely to donate a proton.
• Equilibrium Position:
  • The net direction of an acid-base reaction depends on the relative strengths of the acids and bases involved. A reaction will proceed towards the formation of a weaker acid and a weaker base from a stronger acid and a stronger base. This is because the equilibrium favors the side with weaker acids and bases.
• Examples:
  • Hydrochloric Acid (HCl):
    • HCl is a strong acid because it completely dissociates in water to form H⁺ and Cl⁻.
    • The conjugate base Cl⁻ is very weak because it has a very low tendency to accept a proton and reform HCl.
  • Acetic Acid (CH₃COOH):
    • CH₃COOH is a weak acid because it partially dissociates in water.
    • The conjugate base CH₃COO⁻ is relatively stronger compared to Cl⁻ because it has a greater tendency to accept a proton and reform acetic acid.

1.1.5. Molecular Auto Ionization (Self-Ionization)

• Definition: A reaction between two identical neutral molecules, especially in a solution, to produce an anion and a cation.
  • Example: In the case of water, this can be represented as: 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

Ions Present in Water and their Formation

• Ions in Water: The ions present in water are hydrogen ions (H⁺) and hydroxide ions (OH⁻).
• Formation of Ions: These ions are formed through the process of molecular auto ionization (or self-ionization).

Unique Properties of Water

• Dual Role: Water can act either as an acid or as a base.
• Reactions with Acids and Bases:
  • With Acids: Water acts as a base (e.g., HCl and CH₃COOH).
  • With Bases: Water acts as an acid (e.g., NH₃).

Weak Electrolyte

• Conductivity: Water is a very weak electrolyte and a poor conductor of electricity.
• Ionization: Despite this, water does undergo ionization to a small extent.

Autoionization of Water

• Reaction: The autoionization of water involves water molecules acting as both acids and bases: H₂O(l)+H₂O(l)⇌H₃O⁺(aq)+OH⁻(aq)
• Water Molecule as Acid: Donates a proton.
• Water Molecule as Base: Accepts a proton.

Amphiprotic Species

Definition

• Amphiprotic Species: Molecules or ions that can either donate or accept a proton, depending on the other reactant.

Examples

• Bicarbonate Ion (HCO₃^–): HCO₃ + OH⁻ ⇌ CO₃^2- + H₂O
  • Acts as an Acid: In the presence of OH⁻. HCO₃ + OH⁻ ⇌ CO₃^2- + H₂O
  • Acts as a Base: In the presence of HF. HCO₃ + HF ⇌ H₂CO₃ + F^–
• Water: The most important amphiprotic species.
  • When Acting as a Base: Water accepts a proton from an acid. H₂O +HF ⇌ H₃O⁺ + F^–
  • When Acting as an Acid: Water donates a proton to a base. H₂O +HCO₃ ⇌ OH⁻ + H₂CO₃

Reactions Involving Water

• With Ammonia (NH₃): NH₃(aq)+H₂O(l)⇌NH₄⁺(aq)+OH⁻(aq)
  • Water acts as an acid.
• With Acetic Acid (CH3COOH): CH₃COOH(aq)+H₂O(l)⇌CH₃COO⁻(aq)+H₃O⁺(aq)
  • Water acts as a base.

1.1.6. Lewis Concepts

• G. N. Lewis’s proposed the concept of acids and bases is expanded beyond proton-transfer reactions to include reactions of acidic and basic oxides and many other reactions.

Lewis Acid-Base Definition

• Lewis Base: A species that donates an electron pair to form a bond.
• Lewis Acid: A species that accepts an electron pair to form a bond.

Example: Reaction of Boron Trifluoride (BF₃) with Ammonia (NH₃)

• Reaction Illustration: BF₃+NH₃→F₃B⁻NH₃
  • Lewis Acid: BF₃ (boron trifluoride) accepts an electron pair.
  • Lewis Base: NH₃ (ammonia) donates an electron pair.
• Explanation:
  • The boron atom in BF₃ has only six electrons in its valence shell and needs two more electrons to satisfy the octet rule.
  • BF₃ (Lewis acid) accepts a pair of electrons from NH₃ (Lewis base).

Key Points in a Lewis Acid-Base Reaction

1. Electron Acceptor: Look for a species with an empty orbital to accommodate an electron pair (e.g., the B atom in BF₃).
2. Electron Donor: Look for a species with lone-pair electrons (e.g., the N atom in NH₃).

Lewis Bases and Lewis Acids

• Lewis Bases: Species with electron pairs available to donate, such as:
  • OH⁻
  • NH₃
  • H₂O
• Characteristics of Lewis Bases: Any molecule or negatively charged species having an excess of electrons.
• Lewis Acids: Electron-deficient molecules or positively charged species