Can a buffer run out?

Yes. This is called 'buffer capacity.' A buffer only has a certain amount of weak acid and weak base molecules. If you add enough strong acid to react with all of the weak base present, the buffer is 'broken,' and the pH will start dropping rapidly.

Do all buffers keep the pH at exactly 7?

No. A buffer just maintains the pH at a specific, steady value. That value depends entirely on the $pK_a$ of the weak acid used. For instance, an acetic acid buffer keeps the pH around 4.74, while an ammonia buffer keeps the pH around 9.25.

How do I choose the right chemicals to make a buffer?

You should pick a weak acid that has a $pK_a$ very close to the target pH you want to maintain. Usually, a buffer is effective within one pH unit above or below its $pK_a$ (so a $pK_a$ of 5 is good for a pH range of 4 to 6).

What is a buffer and how does it work?

A buffer is a special type of chemical solution designed to keep its pH relatively constant, even when you add a strong acid or a strong base to it. Think of a buffer like a shock absorber on a car. Just as a shock absorber smooths out a bumpy road so you don't feel every pothole, a buffer absorbs extra acid or base so the overall pH of the solution doesn't experience a massive spike or drop.

To be precise, a buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. When you add a bit of strong acid or base, these two components work together to neutralize the intruders, keeping the free hydrogen ion concentration—and therefore the pH—stable.

What makes up a buffer?

A buffer must have two working parts: a weak acid to neutralize added bases, and a weak base to neutralize added acids. Crucially, these two parts cannot neutralize each other. This is why we use a 'conjugate pair.' A common example is acetic acid ( $CH_3COOH$ ) and its conjugate base, the acetate ion ( $CH_3COO^-$ ), usually added in the form of a salt like sodium acetate. Because they are a conjugate pair, they exist in equilibrium without destroying one another.

How does it neutralize intruders?

When you add a strong acid (which releases $H^+$ ions) to the buffer, the weak base component reacts with the $H^+$ to form more of the weak acid. Because the strong acid is converted into a weak acid, the pH drops only slightly. Conversely, if you add a strong base (which releases $OH^-$ ions), the weak acid component reacts with it to form water and more of the weak base. The buffer essentially trades a strong, pH-swinging chemical for a weak, gentle one.

Where students slip: Strong vs. Weak

A very common mistake is trying to make a buffer out of a strong acid and its conjugate base, like hydrochloric acid ( $HCl$ ) and sodium chloride ( $NaCl$ ). This does not work. Strong acids dissociate completely in water, meaning the chloride ion has virtually no ability to act as a base and absorb added $H^+$ ions. A buffer absolutely requires a weak acid or weak base so that an equilibrium can be established.

Worked through

You have a buffer solution containing 0.50 M acetic acid ( $CH_3COOH$ ) and 0.50 M sodium acetate ( $CH_3COONa$ ). The $K_a$ for acetic acid is $1.8 imes 10^{-5}$ . What is the pH of this buffer?

To find the pH of a buffer, we use the Henderson-Hasselbalch equation: $pH = pK_a + \log\left(\frac{[Base]}{[Acid]}\right)$

First, calculate the $pK_a$ from the given $K_a$ : $pK_a = -\log(1.8 imes 10^{-5}) = 4.74$

Next, identify the concentrations of your acid and conjugate base: [Acid] = 0.50 M ( $CH_3COOH$ ) [Base] = 0.50 M ( $CH_3COO^-$ )

Now, plug these into the equation: $pH = 4.74 + \log\left(\frac{0.50}{0.50}\right)$ $pH = 4.74 + \log(1)$

Since the log of 1 is 0, the equation simplifies to: $pH = 4.74 + 0 = 4.74$

The pH of the buffer is 4.74. Notice that when the concentrations of the acid and base are equal, the pH exactly equals the $pK_a$ !

Questions students ask

Ask about this topic

Explain at

Where this comes from: OpenStax Chemistry 2e, Chapter 14: Acid-Base Equilibria · Khan Academy, Unit: Buffers, titrations, and solubility equilibria