A titration curve is a graph with the volume of titrant added (in mL) on the x-axis and the pH of the solution on the y-axis. Reading one comes down to identifying four key features: the starting pH, the buffer region, the equivalence point, and the final pH. Each of these tells you something specific about the acid or base being analyzed.
What the Axes Tell You
The x-axis tracks how much titrant (the solution being dripped in from a burette) has been added, measured in milliliters. The y-axis tracks the pH of the solution in the flask as it changes in response. The resulting S-shaped curve captures the entire neutralization reaction from start to finish. Every point on the curve represents a snapshot of the solution’s acidity at a specific volume of added titrant.
The Starting pH
Before any titrant is added, the curve’s y-intercept (the leftmost point) tells you the initial pH of the solution in the flask. This first reading already gives you useful information. If you’re titrating an acid with a base, a very low starting pH (around 1 to 2) suggests a strong acid, while a higher starting pH (around 3 to 5) suggests a weak acid. The reverse logic applies when titrating bases: a starting pH near 13 to 14 points to a strong base, and something closer to 9 to 11 suggests a weak one.
The Buffer Region
After the titration begins, you’ll notice a long, relatively flat stretch of the curve where adding titrant barely changes the pH. This is the buffer region. It occurs because the solution contains a mixture of the original acid (or base) and its partially neutralized form, which together resist pH changes. The nearly flat portion extends from roughly one pH unit below the pKa to one pH unit above it. Recognizing this zone matters because it tells you the pH range where the acid-base pair acts as an effective buffer.
The exact center of the buffer region is the half-equivalence point. At this specific volume, exactly half of the original acid has been neutralized, meaning the concentration of the acid equals the concentration of its conjugate base. When those two concentrations are equal, the pH of the solution equals the pKa of the acid. So if you need to find the pKa from a titration curve, locate the equivalence point first, then look at the pH value at exactly half that volume. That pH is your pKa.
The Equivalence Point
The equivalence point is the most dramatic feature on the curve: a near-vertical spike where the pH changes rapidly over a very small addition of titrant. At this point, the moles of titrant added exactly equal the moles of the substance in the flask. The acid has been completely neutralized. On the graph, you find it at the midpoint of the steep vertical section, not at the top or bottom of the rise.
The pH at the equivalence point is one of the most informative readings on the entire curve:
- Strong acid + strong base: pH at the equivalence point equals 7 (neutral).
- Weak acid + strong base: pH at the equivalence point is greater than 7 (slightly basic), because the conjugate base of the weak acid makes the solution basic.
- Strong acid + weak base: pH at the equivalence point is less than 7 (slightly acidic), because the conjugate acid of the weak base makes the solution acidic.
If you need a more precise location for the equivalence point than eyeballing the steep section, you can use derivatives. The first derivative of the curve (plotting the rate of pH change against volume) shows the equivalence point as a sharp peak. The second derivative crosses zero at the equivalence point, giving an even more exact location.
Strong vs. Weak: How Curve Shapes Differ
The overall shape of the curve tells you what type of acid or base you’re dealing with. A strong acid titrated with a strong base starts at a very low pH, stays low with a gradual rise, then shoots up steeply through pH 7, and levels off at a high pH. The curve is roughly symmetrical around the equivalence point.
A weak acid titrated with a strong base looks different in two key ways. First, the starting pH is higher because weak acids don’t fully dissociate. Second, there’s a visible, extended buffer region before the steep rise, and the equivalence point lands above pH 7 rather than at it. These differences make it straightforward to tell strong and weak acid curves apart at a glance: look for the starting pH and the position of the equivalence point relative to 7.
When a weak acid is titrated with a weak base, the curve loses its dramatic steep section almost entirely. Instead of a sharp vertical rise, you get a gentle S-shape with only an inflection point (a subtle change in curvature) marking the equivalence point. These curves are harder to read and are one reason weak acid/weak base titrations are avoided in practice when possible.
Polyprotic Acids: Multiple Equivalence Points
Some acids can donate more than one proton. These polyprotic acids produce titration curves with multiple steep sections, each representing the neutralization of one proton. A diprotic acid (two protons to donate) shows two steep rises and two equivalence points. A triprotic acid like citric acid shows three.
Each segment of the curve behaves like its own mini-titration. Between each pair of equivalence points, you’ll find another buffer region with its own half-equivalence point, giving you a separate pKa for each proton. For a triprotic acid, the three equivalence points are spaced so that the first occurs at one-third of the total base volume needed, the second at two-thirds, and the third at the full volume. Each halfway point within those segments gives you one of the three Ka values.
Choosing an Indicator From the Curve
Chemical indicators change color over a specific pH range, typically spanning about two pH units centered on the indicator’s own pKa. To pick the right indicator for a titration, you match the indicator’s color-change range to the pH at the equivalence point. For a strong acid/strong base titration with an equivalence point at pH 7, an indicator that changes color near pH 7 works well. For a weak acid/strong base titration where the equivalence point sits around pH 8 or 9, you’d pick an indicator with a pKa near that value, like phenolphthalein. Using an indicator whose range falls outside the steep portion of the curve will give you an inaccurate endpoint.
Amino Acid Titration Curves
In biochemistry, titration curves for amino acids follow the same principles but carry additional information. Simple amino acids have two ionizable groups: a carboxyl group (pKa typically between 1.8 and 2.4) and an amino group (pKa between 8.8 and 9.7). Their titration curves show two inflection points and two buffer regions, one for each group.
The key feature on an amino acid curve is the isoelectric point (pI), the pH where the molecule carries no net charge. For amino acids with neutral side chains, the pI is simply the average of the two pKa values. Alanine, for example, has pKa values of 2.34 and 9.69, giving a pI of 6.02. On the curve, this pH falls between the two buffer regions. For amino acids with acidic side chains, the pI is the average of the two lowest pKa values. For those with basic side chains, it’s the average of the two highest. Knowing how to pull pI from a titration curve is essential for techniques like electrophoresis, where proteins are separated based on their charge at a given pH.