pH of Acids and Bases Study Guide

Introduction

The applications of acid-base chemistry extend beyond what was learned in Lesson 6, "Aqueous Reactions." The acidity of an acid solution can be measured by its pH, and equilibria can be established between an acid and a base to create a buffer.

Acid and Base Definitions

Acids are proton donors (according to Brönsted) or electron acceptors (according to Lewis, this is a more general concept). Strong acids are completely dissociated in water, releasing protons (H ⁺) and anionic conjugate bases. Acids have a sour taste.

Bases are proton acceptors (Brönsted) or electron donors (Lewis).When dissolved in water, strong bases such as NaOH dissociate to release hydroxide ions and sodium cation. Bases have a bitter taste and feel slippery like soap. (See also Lesson 6, Table 6.2, for more on acids and bases.)

Reactions of Acids

When acids react with another substance, the products can be predicted based on solubility rules and the four principal acid reactions (see Table 15.1)

Example:

Show the reaction when HCl reacts with Mg(OH)₂ (the neutralization of stomach acid with milk of magnesia).

2HCl + Mg(OH)₂ MgCl₂ + H₂O

Show the reaction when magnesium carbonate reacts with nitric acid.

Autoionization of Water

In pure water, H₂O dissociates to H⁺ ions (protons) and OH^– ions (hydroxide):

H₂O H⁺ + OH^–

The molar concentration of H⁺ equals the molar concentration of OH^–:

[H⁺] = [OH^–] = 1 * 10^–7 M In turn, the ion product of water is K_w = [H⁺] [OH^–] = 1 * 10^–14

pH Scale

The pH measures the negative logarithm (for the presentation of a very small number in a large scale) of the hydrogen ion concentration (in mol/L):

pH = –log [H⁺]

If pure water has a hydrogen ion concentration of 1 * 10^–7 M, the pH of neutral water is 7. A pH of 7 defines a neutral solution. The pH scale ranges from 0 to 14 with acids in the lower end of the scale (smaller than pH 7), whereas bases are at the higher end (greater than pH 7). Some strong acids will result in a negative pH value.

pH of Strong Acids and Bases

Strong acids and bases completely dissociate (ionize) in solution. Therefore, the concentration of the solute is the same as the concentration of the [H⁺] for acid or [OH^–] for base. A 6.0 M HCl solution produces 6.0 M H⁺ ions, and a 3.5 M solution of sodium hydroxide produces 3.5 M of OH^– ions.

Example 1:

Calculate the pH of a 0.050 M HCl solution.

Solution 1:

Because HCl is a strong acid, the [H⁺] = 0.05 M; pH = – log[H⁺] = – log(0.050) =

Example 2:

Calculate the pH of a 0.0030 M NaOH solution.

Solution 2:

Because NaOH is a strong base, [OH–] = 0.0030 M;

1.0 * 10^–14 = [OH^–] [H⁺] [H⁺] = pH = –log[H⁺] = –log (3.33333 * 10^–12) = 11.477 = 11

pH of Weak Acids

Weak acids are weak electrolytes and do not dissociate completely. An equilibrium exists between the reactants and the products, and the equilibrium constant must be taken into account to solve for the pH value. When a weak acid (HA) is dissolved in water, the conjugate base (A^–) and conjugate acid (H⁺) are formed. The equilibrium constant for an acid is called the acid dissociation constant.

HA A^– + H⁺

[Alternatively written HA + H₂O A^– + H₃O⁺]

and pK_a = –log [K_a]