D38.1 Buffer Solutions

A mixture containing a weak acid and its conjugate base, such as acetic acid and sodium acetate [CH3COOH(aq) + CH3COONa(aq)], or a mixture containing a weak base and its conjugate acid, such as ammonia and ammonium chloride [NH3(aq) + NH4Cl(aq)], is a buffer solution. A buffer solution resists changes in pH when small amounts of a strong acid or a strong base are added.

Two images are shown. Image a on the left shows two beakers that each contain yellow solutions. The beaker on the left is labeled “Unbuffered” and the beaker on the right is labeled “p H equals 8.0 buffer.” Image b similarly shows 2 beakers. The beaker on the left contains a bright orange solution and is labeled “Unbuffered.” The beaker on the right is labeled “p H equals 8.0 buffer.”
Figure: Buffered vs. unbuffered. (a) The unbuffered solution on the left and the buffered solution on the right have the same pH (pH 8), showing the yellow color of the indicator methyl orange. (b) After the addition of 1 mL of a 0.01-M HCl solution, the buffered solution pH has not changed detectably. The unbuffered solution has become acidic, as indicated by the red color of the methyl orange, which turns red at a pH of about 4. (credit: modification of work by Mark Ott)

A solution made from equal concentrations of CH3COOH and CH3COO– is slightly acidic because Ka,acetic acid > Kb,acetate anion. When a strong base, such as NaOH, is added to this solution, the OH– anions can react with the H3O+ cations present in the solution in a very product-favored reaction, decreasing concentrations of H3O+:

1)     NaOH(aq) + H3O+(aq) ⇌ Na+(aq) + 2H2O(ℓ)          K1 = 1.0 × 1014

In response to this decrease in [H3O+] upon the addition of NaOH, the Ka,acetic acid equilibrium reaction:

2)     CH3COOH(aq) + H2O(l) ⇌ CH3COO–(aq) + H3O+(aq)          K2 = Ka,acetic acid = 1.8 × 10-5

shifts towards producing more products, thereby restoring [H3O+] to almost the value it was before NaOH was added. Therefore, the net effect of the added NaOH is to convert some of the CH3COOH to CH3COO−. This is evident if you add the two reaction equations to give:

3)     CH3COOH(aq) + NaOH(aq) ⇌ CH3COONa(aq) + H2O(ℓ)          K3 = K1 × K2 = 1.8 × 109

Overall, there is minimal change in the solution’s H3O+ concentration. (Note that the added NaOH may react directly with H3O+ or CH3COOH, because both species are present in the buffer solution when NaOH is added. Regardless of which acidic species NaOH is reacting with directly, the final result is the same, and there is minimal change in the solution’s pH.)

Similarly, when a strong acid, such as HCl, is added, the net effect of the added H3O+ is to convert some of the CH3COO− to CH3COOH:

CH3COO–(aq) + HCl(aq) ⇌ CH3COOH(aq) + Cl–(aq)          K = 5.6 × 104

And again, there is only a minimal change in the solution’s pH.

These concepts are illustrated in the figure below. When a small amount of a strong acid or strong base is added, a buffer solution can moderate changes to pH because it consists of a weak acid that can react with added strong base as well as a weak base that can react with added strong acid.

This figure begins with a chemical reaction at the top: C H subscript 3 C O O H ( a q ) plus H subscript 2 O ( l ) equilibrium arrow H subscript 3 O superscript positive sign ( a q ) plus C H subscript 3 C O O superscript negative sign ( a q ). Below this equation are two arrows: one pointing left and other pointing right. The arrow pointing left has this phrase written above it, “H subscript 3 O superscript positive sign added, equilibrium position shifts to the left.” Below the arrow is the reaction: C H subscript 3 C O O H ( a q ) left-facing arrow C H subscript 3 C O O superscript negative sign ( a q ) plus H subscript 3 O superscript positive sign. The arrow pointing right has this phrase written above it, “O H subscript negative sign added, equilibrium position shifts to the right.” Below the arrow is the reaction: O H superscript negative sign plus C H subscript 3 C O O H ( a q ) right-facing arrow H subscript 2 O ( l ) plus C H subscript 3 C O O superscript negative sign ( a q ). Below all the text is a figure that resembles a bar graph. In the middle are two bars of equal height. One is labeled, “C H subscript 3 C O O H,” and the other is labeled, “C H subscript 3 C O O superscript negative sign.” There is a dotted line at the same height of the bars which extends to the left and right. Above these two bars is the phrase, “Buffer solution equimolar in acid and base.” There is an arrow pointing to the right which is labeled, “Add O H superscript negative sign.” The arrow points to two bars again, but this time the C H subscript 3 C O O H bar is shorter than that C H subscript 3 C O O superscript negative sign bar. Above these two bars is the phrase, “Buffer solution after addition of strong base.” From the middle bars again, there is an arrow that points left. The arrow is labeled, “Add H subscript 3 O superscript positive sign.” This arrow points to two bars again, but this time the C H subscript 3 C O O H bar is taller than the C H subscript 3 C O O superscript negative sign bar. These two bars are labeled, “Buffer solution after addition of strong acid.”
Figure: Buffer Action. When a strong acid (producing H3O+) is added to a buffer solution, the weak base is converted to its weak conjugate acid. When a strong base (producing OH–) is added to a buffer solution, the weak acid is converted to its weak conjugate base.

The weak base and weak acid in a buffer solution are typically a conjugate acid-base pair, both serving to maintain a single dynamic equilibrium that responds to additions of other acids and bases. If they are not a conjugate acid-base pair, then there would be two dynamic equilibria at play, which significantly complicates the buffering actions.

Activity: pH of a Buffer Solution

Exercise: Characteristics of Buffer Solutions

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