Chemistry · General chemistry II · Concept
Buffers and the Henderson–Hasselbalch equation
A buffer contains a weak acid and its conjugate base in comparable amounts, so it absorbs small additions of strong acid or base with little change in pH. Its pH follows from the Henderson–Hasselbalch equation, pH = pKa + log([A⁻]/[HA]). The same chemistry shapes a weak-acid titration curve, whose half-equivalence point has pH = pKa.
How a buffer works
The weak acid HA neutralizes added OH⁻, and its conjugate base A⁻ neutralizes added H₃O⁺. Each addition converts one form into the other, so the ratio [A⁻]/[HA], and with it the pH, changes only a little.
The Henderson–Hasselbalch equation
Taking the logarithm of the Ka expression gives the pH from the ratio of base to acid. Only the ratio matters, so amounts in moles work as well as concentrations. When the two are equal, pH = pKa.
Adding strong acid or base
React the addition first, in moles: added OH⁻ turns that many moles of HA into A⁻, and added H₃O⁺ turns A⁻ into HA. Then put the new amounts into the Henderson–Hasselbalch equation.
Buffer range and capacity
A buffer works best within about one pH unit of its pKa, where neither form is more than ten times the other. Its capacity, the amount of acid or base it can absorb, grows with the amounts of both forms.
Titration curves
In a titration, a base of known concentration is added to an acid until they have reacted in their stoichiometric ratio: the equivalence point. For a weak acid, the pH starts at the weak-acid value, rises slowly through the buffer region, equals pKa halfway to equivalence and is above 7 at equivalence, because the conjugate base remains. A strong acid titrated with a strong base reaches pH 7.00 at equivalence at 25 °C.
Common mistakes
- Writing the ratio upside down: it is base over acid, [A⁻]/[HA].
- Putting a strong acid or base straight into the Henderson–Hasselbalch equation instead of reacting it first.
- Expecting pH 7 at the equivalence point of a weak-acid titration: the conjugate base makes it basic.
- Using a buffer far from its pKa, where one form is nearly used up.
Key terms
- Buffer
- A solution of a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists pH changes when a little acid or base is added. It stops working once either form is used up.
- Henderson–Hasselbalch equation
- pH = pKa + log([A⁻]/[HA]), the pH of a buffer. It is an approximation that works well when both forms are present in similar, not too dilute, amounts.
- Conjugate acid-base pair
- Two species differing by one proton. An acid loses that proton to form its conjugate base; the base gains it to reform the acid.
- Titration
- Adding a solution of known concentration, the titrant, to a sample until the reaction is complete, to find how much the sample contains. For acids and bases, a titration curve plots pH against volume added.
- Equivalence point
- The point in a titration where just enough titrant has been added to react completely with the sample. Its pH is 7 only for a strong acid with a strong base; halfway there, pH = pKa for a weak acid.
- Indicator
- A dye whose acid and base forms have different colors, so its color change signals the pH, such as phenolphthalein turning pink above about pH 8.2.
- Acid-base neutralization
- An acid reacting with a base to form water and a salt. The resulting solution isn’t always pH 7, especially when a weak acid or base is involved.
Work through an example
Find the pH of a buffer that is 0.100 M acetic acid and 0.150 M sodium acetate. For acetic acid, Ka = 1.8 × 10⁻⁵.
Find the pH of a buffer →Sources and scope
Authored study material. Tool results depend on the stated inputs and model assumptions.
Try in the workspace
Open the example inputs, change a value and keep a useful result on your board.
Find it in Buffers & titration curves Open worked example on a board Chemistry formulas: equilibrium and acidsYour existing work stays on this device. Examples open as editable copies.