Introduction to Acids and Bases

Table 1: Common Acids

Table 2: Common Bases

Defining Acids and Bases

To understand how these substances react, chemists use different definitions. The Arrhenius definition states that an acid produces $H^+$ in solution (which associates with water to form hydronium, $H\_3O^+$), while a base produces $OH^-$.

However, the Brønsted–Lowry definition is more widely applicable. Under this theory:

• An acid is an $H^+$ (proton) donor.
• A base is an $H^+$ acceptor.

When an acid and base react, they typically form a salt and water:

$$\begin{aligned} \text{acid} + \text{base} \rightarrow \text{salt} + \text{water} \\ HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l) \end{aligned}$$

Conjugate Pairs and Amphoterism

In a Brønsted–Lowry reaction, the relationship between reactants and products is described by conjugate pairs. When an acid donates a proton, it becomes a conjugate base. Conversely, when a base accepts a proton, it becomes a conjugate acid.

Some substances are amphoteric, meaning they can act as either an acid or a base depending on the reaction partner. Water is the most famous example: it can act as a base by accepting $H^+$ from $HCl$, or as an acid by donating $H^+$to$NH_3$.

Measuring Strength: $K_a$, $K_b$, and pH

The "strength" of an acid or base refers to how completely it ionizes in water.

Strong vs. Weak

• Strong acids and bases are strong electrolytes; they ionize almost 100% in water. Examples include $HCl$, $HNO_3$, and $NaOH$.
• Weak acids and bases are weak electrolytes; only a small fraction of the molecules ionize. Examples include acetic acid ($HC_2H_3O_2$) and ammonia ($NH_3$).

Ionization Constants

The strength of a weak acid is quantified by its acid ionization constant ($K_a$):

$$\begin{aligned} K_a = \frac{[H_3O^+][A^-]}{[HA]} \end{aligned}$$

A larger $K_a$ indicates a stronger acid. Similarly, the base ionization constant ($K_b$) measures base strength.

Exam Tip: Because $K_a$ and $K_b$ are logarithmic scales, it is often easier to use $pK_a$ and $pK_b$. The relationship is $pK_a = -\log(K_a)$. Note that a smaller $pK_a$ value means a stronger acid.

The pH Scale

The acidity of a solution is expressed as pH, defined as:

$$\begin{aligned} \text{pH} = -\log[H_3O^+] \end{aligned}$$

At $25^\circ\text{C}$, the product of hydronium and hydroxide concentrations is always constant: $[H_3O^+][OH^-] = K_w = 1.00 \times 10^{-14}$. This leads to the relationship:

$$\begin{aligned} \text{pH} + \text{pOH} = 14.0 \end{aligned}$$

• pH < 7: Acidic solution
• pH = 7: Neutral solution
• pH > 7: Basic solution

Polyprotic Acids

Some acids, known as polyprotic acids, can donate more than one proton per molecule. These acids ionize in steps, and each step has its own ionization constant ($K_{a1}, K_{a2}$, etc.). Generally, the first ionization is much stronger than the subsequent ones ($K_{a1} > K_{a2} > K_{a3}$).

For example, in sulfuric acid ($H_2SO_4$), the first ionization is considered strong, while the second ionization is a weak acid reaction with $K_{a2} = 1.2 \times 10^{-2}$.