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Unit 1•Lesson 3 2 min read

Acids, Bases, and pH

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Learning Objectives

Define acids and bases according to the Arrhenius, Brønsted-Lowry, and Lewis theories.
Explain the concept of pH as a measure of acidity and perform basic pH calculations.
Distinguish between strong and weak acids/bases and the concept of equilibrium.
Understand the function of buffers in resisting pH change.

Theories of Acids and Bases

There are three major definitions of acids and bases that expand in scope:

1.Arrhenius Theory:
An acid is a substance that produces hydrogen ions (H⁺) when dissolved in water.
A base is a substance that produces hydroxide ions (OH⁻) when dissolved in water.
2.Brønsted-Lowry Theory: (Most commonly used)
An acid is a proton (H⁺) donor.
A base is a proton (H⁺) acceptor.

This theory introduces the concept of conjugate acid-base pairs. When a Brønsted-Lowry acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid. (e.g., HCl (acid) + H₂O (base) ⇌ Cl⁻ (conj. base) + H₃O⁺ (conj. acid)).

3.Lewis Theory: (Most general)
A Lewis acid is an electron pair acceptor.
A Lewis base is an electron pair donor.

The pH Scale

The pH scale is a logarithmic scale used to specify the acidity or basicity of an aqueous solution.

pH = -log[H⁺] where [H⁺] is the concentration of hydrogen ions in moles per liter.
Acidic solutions: pH < 7
Neutral solutions: pH = 7
Basic (alkaline) solutions: pH > 7

Because it is a logarithmic scale, a change of 1 pH unit represents a 10-fold change in [H⁺]. A solution with pH 3 is 10 times more acidic than a solution with pH 4.

Strong vs. Weak Acids and Bases

Strong Acids/Bases: Dissociate (ionize) completely in water. The reaction goes to completion. Examples: HCl (hydrochloric acid), NaOH (sodium hydroxide).
Weak Acids/Bases: Only partially dissociate in water. The reaction establishes an equilibrium, and only a fraction of the molecules donate/accept protons. The strength of a weak acid is quantified by its acid dissociation constant, Ka. Examples: CH₃COOH (acetic acid), NH₃ (ammonia).

Buffers

A buffer is a solution that can resist pH change upon the addition of an acidic or basic component. It is able to neutralize small amounts of added acid or base, thus maintaining the pH of the solution relatively stable.

A buffer solution is composed of a weak acid and its conjugate base (e.g., acetic acid and sodium acetate) or a weak base and its conjugate acid (e.g., ammonia and ammonium chloride).
Buffers are crucial in biological systems; for example, the bicarbonate buffer system maintains the pH of human blood close to 7.4.

Key Terms

pH
A logarithmic scale used to specify the acidity or basicity of an aqueous solution. It is the negative of the base-10 logarithm of the hydrogen ion concentration.
Brønsted-Lowry Acid
A chemical species that donates a proton (H⁺).
Brønsted-Lowry Base
A chemical species that accepts a proton (H⁺).
Buffer
A solution containing a weak acid and its conjugate base, or a weak base and its conjugate acid, which resists changes in pH when acid or base is added.
Strong Acid
An acid that ionizes completely in an aqueous solution by losing one proton.

Check Your Understanding

1

What is the pH of a solution with a hydrogen ion concentration [H⁺] of 1 x 10⁻⁴ M? Is this solution acidic or basic?

2

According to the Brønsted-Lowry theory, identify the acid, base, conjugate acid, and conjugate base in the following reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.

3

What are the two components of a buffer solution, and how do they work to resist pH changes?

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