pH: Measure of $[H^+]$ $pH = -\log[H^+]$

pOH: Measure of $[OH^-]$ $pOH = -\log[OH^-]$

Relationship: $pH + pOH = 14$ (at 25 degrees C)

Strong acids (completely dissociate):

• $HCl$, $HBr$, $HI$, $HNO_3$, $HClO_4$, $H_2SO_4$ (first $H^+$ only)

Strong bases (completely dissociate):

• Group 1 hydroxides ($NaOH$, $KOH$)
• Heavy group 2 hydroxides ($Ca(OH)_2$, $Sr(OH)_2$, $Ba(OH)_2$)

Use ICE tables to calculate equilibrium concentrations

• For weak acid: $HA \rightleftharpoons H^+ + A^-$
• For weak base: $B + H_2O \rightleftharpoons BH^+ + OH^-$

$\% \text{ionization} = \frac{[H^+]_{equilibrium}}{[HA]_{initial}} \times 100\%$

Salt hydrolysis: Ions of salt react with water

• Cation from weak base + Anion from weak acid -> neutral solution
• Cation from strong base + Anion from weak acid -> basic solution
• Cation from weak base + Anion from strong acid -> acidic solution

Equivalence point: Moles acid = moles base

• Strong acid/strong base: pH = 7
• Weak acid/strong base: pH > 7
• Strong acid/weak base: pH < 7

Half-equivalence point: $pH = pK_a$ (maximum buffer capacity)

Binary acids ($HX$): Strength increases with:

• Electronegativity of X (left to right)
• Size of X (down a group for same group)
• Bond strength (weaker bond -> stronger acid)

Oxyacids ($HXO_n$): Strength increases with:

• Number of oxygen atoms (more O -> stronger acid)
• Electronegativity of central atom

Acid dissociation constant ($K_a$): $K_a = \frac{[H^+][A^-]}{[HA]}$

Base dissociation constant ($K_b$): $K_b = \frac{[OH^-][BH^+]}{[B]}$

Relationship: $K_a \times K_b = K_w = 1.0 \times 10^{-14}$

Buffer: Solution that resists pH change

• Consists of weak acid + conjugate base OR weak base + conjugate acid
• Works by consuming added $H^+$ or $OH^-$

$pH = pK_a + \log\frac{[A^-]}{[HA]}$