Buffer Summary

Key Concepts and Summary

Solutions that contain appreciable amounts of a weak conjugate acid-base pair are called buffers. A buffered solution will experience only slight changes in pH when small amounts of acid or base are added. Addition of large amounts of acid or base can exceed the buffer capacity, consuming most of one conjugate partner and preventing further buffering action.

Key Equations

pKa = −log Ka
pKb = −log Kb
$\text{pH}=\text{p}K_a+log\frac{[A^-]}{[HA]}$

Practice Questions

Explain why a buffer can be prepared from a mixture of NH4Cl and NaOH but not from NH3 and NaOH.

Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the acid H3PO4 and a salt of its conjugate base NaH2PO4.

Solution

Excess H3O+ is removed primarily by the reaction:

$$H_3O^+(aq)+H_2PO_4^-(aq)⟶H_3PO_4(aq)+H_2O(l)$$



Excess base is removed by the reaction:

$$OH^-(aq)+H_3PO_4(aq)⟶H_2PO_4^-(aq)+H_2O(l)$$

 

Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the base NH3 and a salt of its conjugate acid NH4Cl.

What is [H3O+] in a solution of 0.25 M CH3CO2H and 0.030 M NaCH3CO2?

$$CH_3CO_2H(aq)+H_2O(l)⇌H_3O^+(aq)+CH_3CO_2^-(aq)\qquad K_a=1.8×10^{-5}$$

 

Solution

[H3O+] = $1.5×10^{-4}\;M$

What is [H3O+] in a solution of 0.075 M HNO2 and 0.030 M NaNO2?

$$HNO_2(aq)+H_2O(l)⇌H_3O^+(aq)+NO_2^-(aq)\qquad K_a=4.5×10^{-5}$$

 

What is [OH−] in a solution of 0.125 M CH3NH2 and 0.130 M CH3NH3Cl?

$$CH_3NH_2(aq)+H_2O(l)⇌CH_3NH_3^+(aq)+OH^-(aq)\qquad K_b=4.4×10^{-4}$$

 

Solution

[OH−] = $4.2×10^{-4}\;M$

What is [OH−] in a solution of 1.25 M NH3 and 0.78 M NH4NO3?

$$NH_3(aq)+H_2O(l)⇌NH_4^+(aq)+OH^-(aq)\qquad K_b=1.8×10^{-5}$$

 

What is the effect on the concentration of acetic acid, hydronium ion, and acetate ion when the following are added to an acidic buffer solution of equal concentrations of acetic acid and sodium acetate:

(a) HCl

(b) KCH3CO2

(c) NaCl

(d) KOH

(e) CH3CO2H

Solution

(a) The added HCl will increase the concentration of H3O+ slightly, which will react with $CH_3CO_2^-$ and produce CH3CO2H in the process. Thus,$[CH_3CO_2^-]$ decreases and [CH3CO2H] increases.
(b) The added KCH3CO2 will increase the concentration of $[CH_3CO_2^-]$ which will react with H3O+ and produce CH3CO2 H in the process. Thus, [H3O+] decreases slightly and [CH3CO2H] increases.
(c) The added NaCl will have no effect on the concentration of the ions.
(d) The added KOH will produce OH− ions, which will react with the H3O+, thus reducing [H3O+]. Some additional CH3CO2H will dissociate, producing $[CH_3CO_2^-]$ ions in the process. Thus, [CH3CO2H] decreases slightly and $[CH_3CO_2^-]$ increases.
(e) The added CH3CO2H will increase its concentration, causing more of it to dissociate and producing more $[CH_3CO_2^-]$ and H3O+ in the process. Thus, [H3O+] increases slightly and $[CH_3CO_2^-]$ increases.

What is the effect on the concentration of ammonia, hydroxide ion, and ammonium ion when the following are added to a basic buffer solution of equal concentrations of ammonia and ammonium nitrate:

(a) KI

(b) NH3

(c) HI

(d) NaOH

(e) NH4Cl

What will be the pH of a buffer solution prepared from 0.20 mol NH3, 0.40 mol NH4NO3, and just enough water to give 1.00 L of solution?

Solution

pH = 8.95

Calculate the pH of a buffer solution prepared from 0.155 mol of phosphoric acid, 0.250 mole of KH2PO4, and enough water to make 0.500 L of solution.

How much solid NaCH3CO2•3H2O must be added to 0.300 L of a 0.50-M acetic acid solution to give a buffer with a pH of 5.00? (Hint: Assume a negligible change in volume as the solid is added.)

Solution

37 g (0.27 mol)

What mass of NH4Cl must be added to 0.750 L of a 0.100-M solution of NH3 to give a buffer solution with a pH of 9.26? (Hint: Assume a negligible change in volume as the solid is added.)

A buffer solution is prepared from equal volumes of 0.200 M acetic acid and 0.600 M sodium acetate. Use $1.80×10^{-5}$ as Ka for acetic acid.

(a) What is the pH of the solution?

(b) Is the solution acidic or basic?

(c) What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to 0.200 L of the original buffer?

Solution

(a) pH = 5.222;
(b) The solution is acidic.
(c) pH = 5.221

A 5.36–g sample of NH4Cl was added to 25.0 mL of 1.00 M NaOH and the resulting solution

diluted to 0.100 L.

(a) What is the pH of this buffer solution?

(b) Is the solution acidic or basic?

(c) What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to the solution?