Buffer Solutions

Definition and Components

Buffer solutions consist of a weak acid and its conjugate weak base.
Examples:
- Hydrofluoric acid (HF): Weak acid with conjugate base fluoride (F-), paired with sodium to form a buffer.
- Acetic acid (CH3COOH): Weak acid with conjugate base acetate (CH3COO-), paired with sodium acetate.
- Hydrocyanic acid (HCN): Weak acid with conjugate base sodium cyanide.

Maintains constant pH: Resists changes in pH by neutralizing added acids or bases.
Mechanism:
- Acids added to the buffer react with the weak base component.
- Bases added react with the weak acid component.

pH and pOH relationship: pH = 14 - pOH.
Addition of strong acids/bases changes pH and pOH but is moderated by buffer reactions.

Used to calculate the pH of a buffer solution:

[ \text{pH} = \text{pKa} + \log \left( \frac{[\text{Base}]}{[\text{Acid}]} \right) ]
pKa is derived from the acid dissociation constant ( K_a ):
- ( \text{pKa} = -\log(K_a) )

Equal concentrations: When [HA] = [A-], pH = pKa.
Concentration ratios:
- 10:1 ratio of base to acid increases pH by 1 unit above pKa.
- 1:10 ratio decreases pH by 1 unit below pKa.

Acetic Acid and Sodium Acetate
- Given: 0.75 M acetic acid, 0.5 M sodium acetate, ( K_a = 1.8 \times 10^{-5} )
- Calculate pH using Henderson-Hasselbalch equation.
Ammonium Chloride and Ammonia
- Given: 0.15 moles NH4Cl, 1.5 moles NH3, ( K_b = 1.8 \times 10^{-5} )
- Calculate pH, noting the ratio of base to acid affects the pH relative to pKa.
Hydrofluoric Acid and Sodium Fluoride
- Given: 15g HF, 21g NaF, 750 mL solution, ( K_a = 7.2 \times 10^{-4} )
- Convert grams to moles and use the equation to find pH.
Unknown Acid
- Given: Solution pH of 5.62, [HA] = 0.45 M, [A-] = 0.85 M
- Use the equation to determine pKa.

Buffer solutions are crucial for maintaining stable pH levels in various chemical and biological systems.
Understanding the relationship between pH, pKa, and concentration ratios is key to using buffer solutions effectively.