Topic 8.8: Properties of Buffers

Properties of buffers: why solutions resist a pH change

Two runners are about to race four hundred meters in this hall in Paris, one against one, and they’ve been told to do exactly one thing, win. The four hundred is brutal, because it’s too long to sprint the whole way, and everybody sprints it anyway. Runners even have a name for the last hundred meters. They say the bear jumps on your back.

Keep an eye on her breathing. Look how hard she’s working. Four minutes after the race, a scientist takes a single drop of blood from her fingertip. This really happened, right here in this hall, in a study of twenty-two trained four-hundred-meter runners published in 2024. In one group of eleven of those runners, the average pH of their blood before warming up was seven point three two, and four minutes after the race it was six point nine seven.

pH is a log scale, so that drop means the hydronium in their blood more than doubled. The same samples showed something even stranger, though. A substance in blood called bicarbonate fell from about twenty-four millimoles per liter to under seven, and bicarbonate gets used up when it catches hydronium ions.

Meanwhile the free hydronium went up by less than a ten-thousandth of a millimole per liter. The bicarbonate bar is about three hundred thousand times taller, so we have to zoom way in just to find it.

So their blood caught most of the acid their muscles released, though not all of it, because the pH really did fall, and the chemistry doing the catching is called a buffer. And that little analyzer never measured the bicarbonate at all. It measured the pH and the carbon dioxide, and then it calculated the bicarbonate with a single equation.

So how does a buffer catch acid like that, and what’s that equation? We’ll find out. Then we’ll come back to this hall, and from one pH reading you’ll work out what was happening in the runners’ blood.

Blood is complicated, though. So first, let’s mix a buffer in a beaker. Then we’ll race it against water. On the left is a liter of water with a trace of hydrochloric acid, and on the right a liter of buffer, acetic acid and sodium acetate, a tenth of a mole of each. Both start at exactly the same pH, four point seven four.

Both also have a dye called bromocresol green, which is green right around this pH, yellow when things get more acidic and blue when they get more basic. We’re going to drop one teaspoon of strong acid into each cylinder at the same moment. Before we do, what do you think happens to the pH of the buffer on the right?

The same teaspoon of acid is going into both of them at exactly the same moment. The water turned completely yellow, and its pH crashed from four point seven four to two point three. The buffer flashed yellow for a split second, went straight back to green, and reads four point seven oh, so look closely, because it did move.

Fresh cylinders, and this time a teaspoon of strong base, sodium hydroxide. The water shoots up to eleven point seven, and the buffer barely notices, at four point seven nine. This exact setup is a classic AP question, where they give you two solutions and the same addition and ask which pH changes less, and why.

’Because it’s a buffer’ just gives it a name; it doesn’t explain anything. The explanation AP is looking for names the particle that reacts with the added acid and shows its reaction, which is exactly what we’re about to do. So what’s in that right-hand cylinder that the water doesn’t have? Let’s go inside it and look.

How a buffer works: two partners

This is the buffer up close. The orange ones are acetic acid molecules, the blue ones are acetate ions, about the same number of each, and the little silver ones are sodium ions, with water all around.

Each one stands for an enormous number of real particles. Let’s send in some acid. A hydronium ion is just a water molecule carrying an extra proton, and when it bumps into an acetate, the proton jumps across onto the acetate’s oxygen. Now that acetate is acetic acid, and the hydronium is just water again.

So the acetate caught the proton, and we can write that as one reaction, acetate plus hydronium gives acetic acid plus water. Now there’s a hydroxide ion on its way in. I want you to find the particle that’s going to catch it. Tap whichever one you think the hydroxide is going to react with.

So here’s what a buffer actually is. It’s a solution with a lot of both members of a conjugate acid-base pair, here acetic acid and acetate, which are the same particle with and without one proton. The base partner, acetate, catches any acid you add, and the acid partner, acetic acid, catches any base you add. Those two reactions are the whole reason a buffer resists a change in pH.

Sodium is in the formula. But in the solution it’s a spectator, while acetate does the work. And a lot of both means a lot compared with what you’re adding. Our buffer had a tenth of a mole of each partner, and the teaspoon of acid was only five thousandths of a mole, so there was plenty of acetate to go around.

Two ideas about buffers sound right and aren’t. The first is that a buffer keeps the pH constant, and you just watched one go from four point seven four to four point seven oh. It keeps the change small, which is a different claim. The second is that a buffer neutralizes everything to pH seven. An acetate buffer sits near four point seven and an ammonia buffer near nine point three, and each one holds the pH near its own value.

So how do you spot a buffer? Ask whether the mixture has something to catch an added acid and something to catch an added base. Acetic acid with sodium acetate has both, so it’s a buffer.

Hydrochloric acid with sodium chloride looks similar, but chloride won’t take a proton back from hydronium, because H C L is a strong acid, so nothing catches added acid and it isn’t a buffer. Acetic acid by itself barely ionizes, so there’s hardly any acetate to catch added acid, and one partner isn’t enough.

Ammonia with ammonium chloride is a buffer too. Ammonia catches added acid and becomes ammonium, and ammonium hands a proton to added hydroxide and turns back into ammonia. This base partner has no charge and the acid partner is positive, so you can’t tell the roles from the charges, only from which one has the extra proton. Okay, now you try one. Which of these four mixtures is a buffer?

The buffer in your blood

Now let’s go back to the hall. Her blood is running the same trick, with a different pair. While she was racing, her working muscles were releasing hydrogen ions into her blood much faster than at rest, and each one rides on a water molecule as a hydronium ion.

Her blood is full of bicarbonate ions, the base partner of the main buffer in blood, and its acid partner is carbonic acid, bicarbonate with one more proton. Those red ones are hydronium ions coming in from the muscle, and the blue ones are bicarbonate. Tap the ion in her blood that’s going to catch a hydronium.

Then something happens that never happened in our beaker. Carbonic acid falls apart into carbon dioxide and water, almost as fast as it forms. That carbon dioxide rides in her blood to her lungs, crosses into the air, and she breathes it out. So with every one of those breaths, she’s breathing out carbon dioxide, and some of it came from bicarbonate catching a hydrogen ion. Her blood is a buffer whose acid partner can leave.

Now think about what each catch does to the buffer. It turns one base partner into one acid partner, so a hundred of each becomes ninety-nine and a hundred and one. So what matters is the balance between the two partners, because for any one buffer pair that balance is what sets the pH, and there’s an equation that turns it into a number.