The Core of Acid-Base Chemistry: Definitions, pH, and Conjugate Pairs

Acid-base chemistry is a cornerstone of the physical sciences, appearing in everything from biological enzyme reactions to industrial manufacturing. For students, however, it often presents a hurdle: shifting between multiple definitions (Arrhenius vs. Brønsted-Lowry), mastering logarithmic math (pH), and predicting how molecules behave in water.

Whether you are calculating the pH of a strong acid or identifying conjugate pairs in a complex equilibrium, success comes down to understanding the movement of protons ($H^+$) and electrons. This guide breaks down the core theories, the mathematics of the pH scale, and the essential rules for identifying strong acids and conjugate pairs.

Arrhenius, Brønsted-Lowry, and Lewis Theories

To master acid-base questions, you must first determine which definition the problem requires. Chemistry has evolved three major theories to describe acidity.

1. Arrhenius Theory (The Classical Definition)

Svante Arrhenius provided the earliest definition, focusing on what ions are produced when a substance dissolves in water.

• Arrhenius Acid: A substance that increases the concentration of hydrogen ions, $H^+$, in water (e.g., $HCl \rightarrow H^+ + Cl^-$).
• Arrhenius Base: A substance that increases the concentration of hydroxide ions, $OH^-$, in water (e.g., $NaOH \rightarrow Na^+ + OH^-$).

2. Brønsted-Lowry Theory (The Proton Transfer Definition)

This is the most common definition used in general chemistry exams because it explains reactions that occur without hydroxide ions.

• Acid: A proton ($H^+$) donor.
• Base: A proton ($H^+$) acceptor.

Illustrative Example: The Cyanide Ion
Consider an aqueous solution of cyanide ($CN^-$). Will the pH be greater or less than 7? To answer this, we apply Brønsted-Lowry logic:

1. $CN^-$ has a negative charge and lone pair electrons. It has no protons to donate, so it cannot be an acid.
2. It can, however, accept a proton from water.
3. The reaction is: $$CN^-(aq) + H_2O(l) \rightleftharpoons HCN(aq) + OH^-(aq)$$

Because the cyanide ion accepts a proton from water, generating hydroxide ions ($OH^-$) in the process, it acts as a base. Consequently, the solution will be basic with a pH > 7.

3. Lewis Theory (The Electron Definition)

The broadest definition, often used in organic chemistry.

• Lewis Acid: An electron-pair acceptor.
• Lewis Base: An electron-pair donor.

Understanding the pH and pOH Scales

The concentration of hydrogen ions in a solution can range from very high (10 M) to incredibly low ($10^{-15}$ M). To make these numbers manageable, we use the logarithmic pH scale.

The Core Formulas

The "p" in pH stands for the negative base-10 logarithm. Therefore:

Because water naturally dissociates into protons and hydroxide ions ($H_2O \rightleftharpoons H^+ + OH^-$), there is a constant relationship between the two. At standard temperature (25°C), the product of their concentrations is always $1.0 \times 10^{-14}$ (the dissociation constant, $K_w$).

This leads to the most useful equation for conversion:

Interpreting the Scale

• pH < 7: Acidic (High $[H^+]$).
• pH = 7: Neutral ($[H^+] = [OH^-] = 1.0 \times 10^{-7} M$).
• pH > 7: Basic (High $[OH^-]$).

Strong vs. Weak Acids and Bases

Differentiation between "strong" and "weak" refers to the extent of dissociation, not the concentration of the chemical.

Strong Acids: Complete Dissociation

A strong acid dissociates essentially 100% in water. If you dissolve 1.0 mole of $HCl$ in water, you get 1.0 mole of $H^+$ and 1.0 mole of $Cl^-$. There are virtually no $HCl$ molecules left intact.

Memorize these 7 Common Strong Acids:

1. Hydrochloric acid ($HCl$)
2. Hydrobromic acid ($HBr$)
3. Hydroiodic acid ($HI$)
4. Nitric acid ($HNO_3$)
5. Chloric acid ($HClO_3$)
6. Perchloric acid ($HClO_4$)
7. Sulfuric acid ($H_2SO_4$) — (Note: Only the first proton dissociates completely).

Nomenclature Note: Naming Oxyacids

Identifying acids often requires knowing the rules for polyatomic ions. The name of the acid is derived from the anion:

• If the anion ends in -ate, the acid is -ic acid. (e.g., Chlorate $ClO_3^- \rightarrow$ Chloric acid $HClO_3$).
• If the anion ends in -ite, the acid is -ous acid. (e.g., Phosphite $PO_3^{3-} \rightarrow$ Phosphorous acid $H_3PO_3$).
• Prefixes: "Hypo-" indicates one less oxygen than "-ite"; "Per-" indicates one more oxygen than "-ate".

Weak Acids: Partial Dissociation

Weak acids, such as Hydrofluoric acid ($HF$) or Acetic acid ($CH_3COOH$), establish an equilibrium where most of the acid remains bound (undissociated). Calculating the pH of these solutions requires an ICE table and the acid dissociation constant ($K_a$).

Identifying Conjugate Acid-Base Pairs

In Brønsted-Lowry theory, acids and bases always come in pairs. When an acid donates a proton, it becomes a conjugate base. When a base accepts a proton, it becomes a conjugate acid.

The Golden Rule: A conjugate pair differs by exactly one proton ($H^+$).

Step-by-Step Case Study: Phosphate Buffer

Let's identify the pairs in the following equilibrium reaction:

1. Analyze Species 1 ($H_2PO_4^-$): It turns into $HPO_4^{2-}$. To do this, it lost an $H^+$. Therefore, $H_2PO_4^-$ is the Acid, and $HPO_4^{2-}$ is its Conjugate Base.
2. Analyze Species 2 ($H_2O$): It turns into $H_3O^+$. To do this, it gained an $H^+$. Therefore, $H_2O$ is the Base, and $H_3O^+$ is its Conjugate Acid.

Practice Drills

Can you identify the conjugate acid of the following bases?

• Base: $Br^-$
  Action: Add $H^+$
  Conjugate Acid: $HBr$ (Hydrobromic acid)
• Base: $NH_3$
  Action: Add $H^+$
  Conjugate Acid: $NH_4^+$ (Ammonium ion)

Properties and Indicators: Testing Acids and Bases

Beyond the formulas, acids and bases have distinct physical properties and reactions with indicators.

General Properties

• Conductivity: Both acids and bases conduct electricity in solution because they dissociate into mobile ions (electrolytes).
• Tactile Feel: Bases often feel slippery. This is because they react with the fatty acids in your skin oils in a process called saponification (essentially making soap on your skin).
• Neutralization: An acid reacts with a base to produce water and a salt.

Indicators

Indicators are weak acids or bases that change color depending on the pH of the environment.

• Litmus Paper:
  • Turns Red in Acid (Think "Red" for "Risk" or "Acid").
  • Turns Blue in Base (Think "Blue" for "Base").
• Phenolphthalein: Colorless in acidic/neutral solutions; turns bright pink/magenta in basic solutions.