Chapter 17: Acids and Bases – General Chemistry II Study Notes

Acids and Bases

Definitions of Acids and Bases

Understanding acids and bases is fundamental in chemistry. There are three main definitions: Arrhenius, Brønsted-Lowry, and Lewis. Each definition broadens the scope of what can be considered an acid or a base.

• Arrhenius Acids and Bases: An Arrhenius acid produces H+ ions in water, while an Arrhenius base produces OH− ions in water.

• Brønsted-Lowry Acids and Bases: A Brønsted-Lowry acid is a proton (H+) donor, and a Brønsted-Lowry base is a proton acceptor. Acid-base reactions involve the transfer of a proton from the acid to the base.

• Lewis Acids and Bases: A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor. This definition includes more substances, such as metal ions, as acids.

Conjugate Acid-Base Pairs

In Brønsted-Lowry theory, acids and bases always occur in pairs. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

• Every acid has a conjugate base, and every base has a conjugate acid.

• Example: NH3 + H2O ⇌ NH4+ + OH−

Water as a Solvent

Auto-Ionization of Water

Water is an amphoteric substance, meaning it can act as both an acid and a base. In pure water, two water molecules react to form hydronium and hydroxide ions:

• H2O(l) + H2O(l) ⇌ H3O+(aq) + OH−(aq)

• Or simply: H2O(l) ⇌ H+(aq) + OH−(aq)

Ion-Product Constant of Water, Kw

The equilibrium constant for the auto-ionization of water is called the ion-product constant, Kw:

• $K_w = [H_3O^+][OH^-] = [H^+][OH^-]$

• At 25°C, $K_w = 1.0 \times 10^{-14}$

• In neutral water, $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}$ M$

pKw, pH, and pOH

The pH and pOH scales are logarithmic measures of hydronium and hydroxide ion concentrations, respectively:

• $pH = -\log[H_3O^+]$

• $pOH = -\log[OH^-]$

• $pK_w = -\log K_w = 14.00$ at 25°C

• $pH + pOH = 14.00$ at 25°C

pH Scale and Measurements

The pH Scale

The pH scale typically ranges from 0 (very acidic) to 14 (very basic), but values outside this range are possible. Each unit change in pH represents a tenfold change in [H+].

• pH < 7: Acidic solution ($[H^+] > [OH^-]$)

• pH = 7: Neutral solution ($[H^+] = [OH^-]$)

• pH > 7: Basic solution ($[H^+] < [OH^-]$)

Calculating pH and pOH

• To find [H3O+] from pH: $[H_3O^+] = 10^{-pH}$

• To find [OH−] from pOH: $[OH^-] = 10^{-pOH}$

• To find pH from [OH−]: Calculate pOH, then use $pH = 14.00 - pOH$

pH Measurement

The pH of a solution is most accurately measured using a pH meter, which uses a glass electrode sensitive to [H3O+]. Less accurate methods include pH paper and indicators.

Weak Acids: Acid Dissociation Constant (Ka), % Ionization, and pH Calculations

Strong vs. Weak Acids

Strong acids ionize completely in water, while weak acids only partially ionize. The concentration of H+ in a strong acid solution equals the acid's molarity. For weak acids, equilibrium calculations are required.

• Strong acids to memorize: HCl, HBr, HI, HNO3, HClO4, HClO3, H2SO4 (first proton only)

Acid Dissociation Constant (Ka)

The strength of a weak acid is quantified by its acid dissociation constant, Ka:

• For HA(aq) + H2O(l) ⇌ H3O+(aq) + A−(aq):

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• The larger the Ka, the stronger the acid.

• pKa is defined as $pK_a = -\log K_a$; smaller pKa means a stronger acid.

% Ionization of a Weak Acid

The percent ionization of a weak acid is the fraction of acid molecules that ionize in solution, expressed as a percentage:

• $\%\text{ionization} = \frac{[H_3O^+]_{eq}}{[HA]_i} \times 100$

• As the acid is diluted, % ionization increases.

Calculating pH of a Weak Acid Solution

To calculate the pH of a weak acid solution, set up an ICE (Initial, Change, Equilibrium) table and solve for [H3O+] using Ka. If % ionization is less than 5%, the change in concentration (x) can be neglected in the denominator for simplification.

• Example: For 0.15 M acetic acid (Ka = 1.8 × 10−5):

• Set up: $K_a = \frac{x^2}{0.15}$, solve for x, then $pH = -\log x$

• If % ionization ≤ 5%, the approximation is valid.

Concentration	[HC2H3O2], M	[H3O+], M	[C2H3O2−], M
Initial	0.15	0	0
Change	−x	+x	+x
Equilibrium	0.15−x	x	x

Additional info: The ICE table method is a standard approach for equilibrium calculations in weak acid and base chemistry.