A weak acid–strong base titration determines the concentration of a weak acid by gradually adding a strong base of known concentration until the acid has reacted completely. As the base is added, the solution’s pH changes in a predictable pattern, allowing a chemist to identify the amount of acid present and calculate its concentration.
This type of titration is different from a strong acid–strong base titration because a weak acid only partially ionizes in water. Its reaction with a strong base, however, proceeds essentially to completion. This difference affects the solution’s pH throughout the titration, particularly at the equivalence point, where the amount of base added is chemically equivalent to the original amount of acid.
What is a weak acid–strong base titration?
A titration is a laboratory technique used to determine the concentration of a substance by reacting it with a solution of known concentration. The solution being added is called the titrant, and the solution being analyzed is called the analyte.
In a weak acid–strong base titration, the analyte is a weak acid, such as acetic acid, while the titrant is a strong base, such as sodium hydroxide. A weak acid does not release all its available hydrogen ions into water. Instead, it establishes an equilibrium between its ionized and unionized forms. A strong base, by contrast, dissociates almost completely in water, providing hydroxide ions that react with the acid.
For example, acetic acid, the acid responsible for vinegar’s characteristic acidity, reacts with sodium hydroxide according to this net ionic equation:
CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O
Here, CH₃COOH represents acetic acid, OH⁻ represents the hydroxide ion, CH₃COO⁻ represents the acetate ion, and H₂O represents water.
The hydroxide ions remove acidic hydrogen from the acetic acid, producing water and acetate ions. As long as acid remains available, the added hydroxide is consumed by the reaction.
By measuring how much sodium hydroxide is required to reach the equivalence point, a chemist can determine the original amount of acetic acid and, if the sample volume is known, its concentration.
How the titration proceeds
The titration begins with a measured volume of weak acid in a flask. A few drops of a suitable acid–base indicator may be added, or a pH meter may be used to monitor the solution continuously.
The strong base is delivered gradually from a burette, a graduated laboratory instrument that allows precise measurement of the volume added. Each addition introduces hydroxide ions, which react with the weak acid. The pH is monitored as the reaction proceeds.
At the beginning, the solution contains mostly the weak acid, along with a smaller amount of its ions. Its initial pH depends on the acid’s concentration and its tendency to ionize in water.
As the base is added, some of the weak acid is converted into its conjugate base. A conjugate base is the species that remains after an acid donates a proton, or hydrogen ion. In the acetic acid example, acetate is the conjugate base.
Before the equivalence point, both the weak acid and its conjugate base are present in substantial amounts. Together, they form a buffer solution, which resists large changes in pH when small quantities of acid or base are added.
As more base is introduced, the amount of weak acid decreases and the amount of conjugate base increases. The pH rises gradually through much of this region because the buffer moderates the effect of additional hydroxide ions.
Near the equivalence point, a relatively small addition of base produces a much larger pH change. Once the equivalence point has been passed, the base is present in excess, and the pH is increasingly determined by the remaining hydroxide ions.
Why the pH changes in a distinctive pattern
The pH curve for a weak acid–strong base titration typically begins at a value above that of a comparable strong acid at the same concentration. This is because the weak acid only partially ionizes in water, producing fewer free hydrogen ions.
The curve then rises gradually as the acid is converted into its conjugate base. This buffer region is an important feature of the titration because it shows how the two forms of the acid work together to limit changes in pH.
At the half-equivalence point, exactly half of the original weak acid has been neutralized. Under the usual assumptions for a simple monoprotic weak acid, the concentrations of the acid and its conjugate base are equal. The Henderson–Hasselbalch equation then gives a particularly useful relationship:
pH = pKₐ
The pKₐ is a measure of an acid’s tendency to donate a proton. A lower pKₐ generally indicates a stronger acid, while a higher pKₐ indicates a weaker acid. The half-equivalence point therefore provides a practical way to estimate the pKₐ of a weak acid from its titration curve.
As the titration approaches the equivalence point, the buffer becomes less effective because relatively little unreacted weak acid remains. The pH rises sharply over a comparatively narrow volume range.
The most important difference from a strong acid–strong base titration appears at the equivalence point itself: the pH is generally greater than 7 at 25°C. This happens because the solution contains the conjugate base of the weak acid, which reacts with water to produce hydroxide ions.
After the equivalence point, additional strong base remains in the solution. The pH is then governed mainly by the excess hydroxide ions rather than by the weak acid–conjugate base equilibrium.
Why the equivalence point has a pH above 7
The equivalence point is reached when the amount of strong base added is chemically equivalent to the original amount of acid. For a monoprotic acid, which can donate one proton per molecule, this means that the number of moles of hydroxide added equals the initial number of moles of acid.
Although the acid has been neutralized in the stoichiometric sense, the solution is not necessarily neutral in the pH sense. The conjugate base produced during the reaction can interact with water, a process called base hydrolysis.
For acetic acid, the acetate ion reacts with water as follows:
CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻
This reaction produces a small amount of hydroxide ions, making the solution basic. The equilibrium lies mainly toward acetate and water because acetate is a weak base, but the hydroxide produced is sufficient to raise the pH above 7 under typical conditions.
The strength of the conjugate base depends on the strength of its corresponding acid. A weaker acid generally has a stronger conjugate base, so its conjugate base produces a more noticeable basic effect in water.
The equivalence-point pH depends on the acid’s properties, the solution’s concentration, and the temperature. It is not always the same for every weak acid–strong base titration.
How to calculate the unknown acid concentration
The main quantitative goal of the titration is to determine how much acid was originally present. For a monoprotic weak acid reacting with a strong base that supplies one hydroxide ion per formula unit, the reaction follows a 1:1 mole ratio.
At the equivalence point:
moles of acid = moles of base
The number of moles of a dissolved substance is calculated by multiplying its molar concentration by its volume in liters. Therefore:
MₐVₐ = MᵦVᵦ
In this equation, Mₐ is the molarity of the acid, Vₐ is the volume of the acid, Mᵦ is the molarity of the base, and Vᵦ is the volume of base required to reach the equivalence point. Both volumes must use the same units.
Rearranging the equation gives the unknown acid concentration:
Mₐ = (MᵦVᵦ) ÷ Vₐ
Consider a sample containing 25.0 mL of an unknown acetic acid solution. Suppose it requires 20.0 mL of 0.100 M sodium hydroxide to reach the equivalence point.
Substituting these values into the equation gives:
Mₐ = (0.100 mol/L × 20.0 mL) ÷ 25.0 mL
Mₐ = 0.0800 M
The original acid concentration is therefore 0.0800 mol/L, or 0.0800 M.
This calculation relies on identifying the equivalence point accurately and knowing the base concentration. For acids that can donate more than one proton, the stoichiometric relationship may differ, so the balanced chemical equation must be used rather than assuming a 1:1 ratio.
How to choose an appropriate indicator
An acid–base indicator is a substance that changes color over a particular pH range. It can help identify the endpoint of a titration, the experimentally observed point used to signal that the reaction has reached its intended completion.
The endpoint and equivalence point are related but not identical. The equivalence point is defined by the reaction’s stoichiometry, while the endpoint is the point at which an observable signal, such as an indicator color change, occurs. A good indicator changes color close enough to the equivalence point that the resulting measurement error is small.
Because the equivalence point in a weak acid–strong base titration generally lies above pH 7, an indicator that changes color in the mildly basic region is often appropriate. Phenolphthalein is commonly used because its transition from colorless to pink occurs within a pH range that often overlaps the steep section of the titration curve.
An indicator whose color changes well before or after the equivalence point may produce a less accurate result. The best choice depends on the shape of the titration curve, which in turn depends on the acid’s strength and concentration.
A pH meter offers another approach. By recording pH after successive additions of base, the chemist can construct a titration curve and locate the region of rapid pH change. This approach can be especially useful when the indicator’s color change is difficult to judge or when a more detailed analysis of the acid’s properties is needed.
What affects the accuracy of the titration?
Accurate results depend on careful preparation, measurement, and recognition of the endpoint. The concentration of the strong base must be known reliably, because any error in that concentration carries into the calculated acid concentration.
Volume measurements also matter. The burette should be read carefully, and the volume delivered should be calculated from the difference between the initial and final readings. Adding the base slowly near the endpoint helps prevent overshooting the equivalence point.
Mixing the solution thoroughly after each addition helps ensure that the reaction proceeds throughout the sample and that the measured pH represents the mixture as a whole. When using an indicator, only a small amount is generally needed, since the indicator itself is a weak acid or base and can slightly affect the solution.
Temperature can influence acid–base equilibria and pH measurements. Carbon dioxide absorbed from the air can also affect some solutions, particularly by introducing additional acidity. Such effects may be important when high precision is required.
The assumptions behind the calculation should also be checked. The simple concentration equation works directly for a monoprotic acid reacting with a strong base in a 1:1 ratio. Polyprotic acids, which can donate more than one proton, may have multiple equivalence points if their successive ionization steps are sufficiently distinct. Their calculations require the appropriate reaction stoichiometry.
What a weak acid–strong base titration reveals
Beyond determining an unknown concentration, a weak acid–strong base titration provides information about how an acid behaves in water. The initial pH reflects its partial ionization, the buffer region demonstrates the relationship between the acid and its conjugate base, and the half-equivalence point can reveal its pKₐ.
The equivalence point illustrates an important principle of acid–base chemistry: neutralization does not necessarily produce a solution with a pH of 7. The pH depends not only on whether the original acid and base have reacted in the correct proportions, but also on the chemical behavior of the substances left in the solution.
Together, the titration curve, the neutralization reaction, and the stoichiometric calculation explain why this method is useful. By adding a strong base in measured amounts and observing how the solution responds, a chemist can determine an unknown acid concentration while gaining a clearer understanding of acid strength, chemical equilibrium, and the behavior of solutions.