An acid-base titration is a laboratory technique used to determine the concentration of an acid or a base by reacting it with a solution of known concentration. By carefully measuring how much of the known solution is needed to neutralize the unknown solution, you can calculate its concentration using a simple chemical relationship.
Titration is widely used in chemistry education, pharmaceutical testing, food analysis, water quality monitoring, and other applications where accurate measurements of acidity or alkalinity matter. Although the technique requires careful handling and precise measurements, the underlying principle is straightforward: acids and bases react in predictable proportions.
What is an acid-base titration?
An acid-base titration is a quantitative chemical analysis in which an acid reacts with a base, or a base reacts with an acid, until the reaction reaches a point that allows the unknown concentration to be calculated.
The solution with a known concentration is called the titrant. It is delivered gradually from a piece of glassware called a burette. The solution being analyzed is called the analyte and is typically placed in an Erlenmeyer flask.
For example, suppose you have a sample of hydrochloric acid (HCl) with an unknown concentration. You can add sodium hydroxide (NaOH) of a known concentration until the hydrochloric acid has reacted completely. By measuring the volume of sodium hydroxide required and applying the balanced chemical equation, you can determine the hydrochloric acid concentration.
The reaction is:
HCl + NaOH → NaCl + H₂O
Hydrochloric acid reacts with sodium hydroxide to form sodium chloride, a salt, and water. In this reaction, one mole of hydrochloric acid reacts with one mole of sodium hydroxide. This 1:1 relationship makes the concentration calculation especially simple.
Not every acid-base titration follows a 1:1 ratio. Some acids can donate more than one proton, so the balanced chemical equation must always guide the calculation.
Understanding the equivalence point and endpoint
Two important concepts determine when a titration is complete: the equivalence point and the endpoint.
The equivalence point is the point at which the acid and base have reacted in the exact proportions required by the balanced chemical equation. For a simple reaction between a strong monoprotic acid and a strong base, this means equal amounts, measured in moles, of acid and base have reacted.
The endpoint is the experimentally observed signal used to identify when to stop adding the titrant. In a typical beginner’s experiment, an indicator changes color near the equivalence point.
These points are related but not identical. An indicator may change color slightly before or after the true equivalence point. Choosing an appropriate indicator helps minimize this difference and improves accuracy.
The pH at the equivalence point is not always 7. For a strong acid reacting with a strong base at approximately room temperature, the equivalence-point pH is close to 7. However, a weak acid reacting with a strong base generally has an equivalence-point pH above 7, while a weak base reacting with a strong acid generally has an equivalence-point pH below 7. These differences arise from the behavior of the ions remaining in solution after the neutralization reaction.
Equipment and materials needed
A standard acid-base titration requires a few basic laboratory tools.
- Burette: A graduated glass tube with a stopcock that controls the delivery of the titrant. It allows small volumes to be added accurately.
- Volumetric pipette: A tool designed to transfer a precise volume of the analyte into the flask.
- Pipette filler: A device used to draw liquid into a pipette safely. Never use your mouth to pipette chemicals.
- Erlenmeyer flask: A conical flask that holds the analyte and allows the solution to be swirled without easily spilling.
- Small funnel: Useful for filling the burette, but it should be removed before the titration begins.
- Suitable indicator: A chemical that changes color over a particular pH range, helping identify the endpoint.
- Wash bottle containing distilled or deionized water: Used to rinse glassware and wash droplets down the inside of the flask.
- White background: A sheet of white paper or a white tile helps make subtle indicator color changes easier to see.
You will also need the acid or base of unknown concentration, a titrant of known concentration, and appropriate personal protective equipment.
Wear chemical splash goggles, a laboratory coat or apron, and suitable gloves according to your laboratory’s safety instructions. Even common laboratory acids and bases can cause injury, especially at higher concentrations.
Choosing the right indicator
An acid-base indicator is a substance that changes color as the pH of a solution changes. Different indicators respond over different pH ranges, so the best choice depends on the expected equivalence-point pH.
Phenolphthalein is commonly used in titrations involving a weak acid and a strong base. It is colorless in acidic solution and becomes pink as the solution becomes sufficiently basic. In a titration of a strong acid with a strong base, it can also be suitable because the pH changes sharply near the equivalence point.
Methyl orange changes from red in acidic solution through an intermediate orange range to yellow in less acidic or basic conditions. It is often suitable for titrations involving a strong acid and a weak base, where the equivalence point is acidic.
For a strong acid and a strong base, several indicators may work because the pH changes rapidly near the equivalence point. For weak-acid or weak-base titrations, the indicator must be selected more carefully.
A pH meter is another option. It measures the solution’s pH directly and can be used to follow the pH change throughout the titration. The equivalence point can then be estimated from the resulting titration curve, which plots pH against the volume of titrant added.
How to perform an acid-base titration step by step
A successful titration depends on careful preparation, controlled addition of the titrant, and accurate volume measurements. The following procedure describes a typical titration in which a known sodium hydroxide solution is used to determine the concentration of hydrochloric acid.
1. Prepare the glassware
Clean the burette, pipette, and Erlenmeyer flask before beginning. Contamination can change the effective concentration of a solution or interfere with the reaction.
Rinse the burette with a small amount of the titrant you will use. This conditions the inside surface with the correct solution and prevents residual rinse water from unnecessarily diluting the titrant.
Rinse the pipette with a small amount of the analyte if the laboratory procedure calls for it. The Erlenmeyer flask can generally be rinsed with distilled or deionized water because the amount of analyte transferred into it is determined by the pipette, and additional pure water does not change the number of moles of analyte present.
2. Fill the burette with the titrant
Close the burette stopcock and use a funnel to add the titrant. Fill it above the intended starting level, then open the stopcock briefly to fill the tip and remove any trapped air bubbles.
Remove the funnel before taking the initial reading. If liquid drips from the funnel into the burette during the experiment, it can change the volume measurement.
Check that the burette tip is full and that no liquid is leaking from the stopcock. Record the initial burette reading at eye level to avoid parallax error, which occurs when a reading appears different because it is viewed from an angle.
For a typical transparent aqueous solution, read the bottom of the curved liquid surface, called the meniscus. Use the reading convention appropriate to your burette and record all digits justified by its graduations.
3. Measure the analyte into the flask
Use a volumetric pipette to transfer a known volume of the acid sample into the Erlenmeyer flask. For example, you might use 25.00 mL if that is the volume specified in your laboratory procedure.
Add two or three drops of a suitable indicator, following the indicator’s instructions. Swirl the flask gently so the indicator is distributed evenly.
Place the flask on a white background beneath the burette. This makes it easier to recognize the color change at the endpoint.
4. Begin adding the titrant
Open the stopcock and allow the titrant to flow into the flask while swirling the flask gently. Swirling helps the acid and base mix thoroughly so that each addition reacts with the analyte.
At the beginning, the titrant can usually be added in relatively large increments because the analyte is present in excess. As the titration progresses and the indicator’s color change becomes more noticeable, reduce the flow rate.
Rinse any droplets adhering to the inner walls of the flask down into the main solution with distilled or deionized water. This ensures that any splashed analyte or titrant participates in the reaction. Adding pure water to the flask does not change the number of moles of acid or base already present.
5. Approach the endpoint carefully
As the titration nears completion, add the titrant one drop at a time while continuously swirling the flask. Near the endpoint, a drop may produce a temporary color change that disappears as the solution mixes.
Continue cautiously until the indicator reaches the specified endpoint color. With phenolphthalein, a common endpoint is the first very pale pink color that persists for approximately 30 seconds, provided that this matches the laboratory’s procedure.
Do not continue adding titrant after the endpoint has been reached. Excess titrant can move the solution beyond the desired point and cause a systematic error in the calculated concentration.
6. Record the final burette reading
Read the burette at eye level using the same method as for the initial reading. Record the final volume reading and subtract the initial reading to find the volume of titrant delivered.
The volume delivered is calculated as:
Volume of titrant = final burette reading − initial burette reading
For example, if the initial reading is 0.35 mL and the final reading is 24.80 mL, the volume delivered is:
24.80 mL − 0.35 mL = 24.45 mL
Use the delivered volume, not simply the final burette reading, in the concentration calculation.
7. Repeat the titration
The first titration is often treated as a rough trial to estimate the endpoint volume. Repeat the procedure with fresh portions of the analyte, taking greater care near the endpoint.
Obtain several results that agree closely with one another, according to the precision expected in your laboratory. Such results are called concordant titrations.
Calculate the concentration using the reliable trials rather than automatically averaging every result. A rough trial or a result in which the endpoint was clearly overshot should generally be excluded from the final mean, with the reason documented.
How to calculate the unknown concentration
The central calculation in a titration uses the relationship between the amount of each reactant and the coefficients in the balanced chemical equation.
The amount of a substance, measured in moles, is calculated using:
n = C × V
Here, n is the amount in moles, C is the molar concentration in moles per liter (mol/L), and V is the solution volume in liters (L).
For a reaction between hydrochloric acid and sodium hydroxide:
HCl + NaOH → NaCl + H₂O
The balanced equation shows that one mole of HCl reacts with one mole of NaOH. At the equivalence point:
moles of HCl = moles of NaOH
Therefore:
Cₐ × Vₐ = Cᵦ × Vᵦ
In this equation, Cₐ and Vₐ represent the concentration and volume of the acid, while Cᵦ and Vᵦ represent the concentration and volume of the base. The volumes must use the same units.
Rearranging gives:
Cₐ = (Cᵦ × Vᵦ) ÷ Vₐ
This equation works for the 1:1 reaction between HCl and NaOH. It should not be applied blindly to reactions with different mole ratios.
Worked example
Suppose you titrate 25.00 mL of hydrochloric acid with sodium hydroxide that has a known concentration of 0.1000 mol/L. The average volume of sodium hydroxide required to reach the endpoint is 24.45 mL.
The balanced equation is:
HCl + NaOH → NaCl + H₂O
Because the reaction ratio is 1:1, the acid concentration can be calculated directly:
Cₐ = (0.1000 mol/L × 24.45 mL) ÷ 25.00 mL
The milliliter units cancel because both volumes are expressed in the same unit.
Cₐ = 0.09780 mol/L
The hydrochloric acid concentration is therefore 0.09780 mol/L, based on the stated measurements and the assumption that the endpoint closely represents the equivalence point.
What if the acid and base do not react in a 1:1 ratio?
Consider sulfuric acid reacting with sodium hydroxide:
H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
One mole of sulfuric acid reacts with two moles of sodium hydroxide under the complete-neutralization conditions represented by this equation.
At the equivalence point:
n(NaOH) = 2 × n(H₂SO₄)
The corresponding concentration relationship is:
Cᵦ × Vᵦ = 2 × Cₐ × Vₐ
Therefore:
Cₐ = (Cᵦ × Vᵦ) ÷ (2 × Vₐ)
The factor of 2 comes from the balanced equation. This illustrates why identifying the correct reaction stoichiometry—the numerical relationship between reactants and products—is essential for accurate titration calculations.
For polyprotic acids, which can donate more than one proton, the appropriate relationship also depends on the reaction endpoint and the extent of neutralization.
Common titration errors and how to avoid them
Even when the chemistry is correct, measurement errors can produce inaccurate results. Recognizing the most common problems helps improve both precision and reliability.
Overshooting the endpoint: Adding too much titrant is one of the most common errors. The excess titrant increases the measured volume and can cause the calculated analyte concentration to be too high or too low, depending on which solution is unknown. Approach the endpoint slowly and add titrant dropwise near the color change.
Reading the burette incorrectly: Viewing the meniscus from above or below eye level introduces parallax error. Keep your eyes level with the liquid surface and use the same reading method at the beginning and end.
Leaving air in the burette tip: An air bubble can fill with titrant during the experiment, causing the apparent delivered volume to exceed the volume that actually entered the flask. Check the tip before recording the initial reading.
Using contaminated glassware: Residual chemicals can react with the analyte or titrant and alter the results. Clean the apparatus thoroughly and rinse the burette and pipette with the appropriate solutions when required.
Choosing an unsuitable indicator: If the indicator changes color far from the equivalence point, the measured endpoint volume may not accurately represent the required reaction volume. Select an indicator whose transition range matches the expected steep pH change.
Failing to mix the solution: Without adequate swirling, the titrant may create a localized region of high or low pH. The indicator can change color temporarily before the solution becomes uniform. Swirl continuously and allow each addition to mix before deciding whether the endpoint has been reached.
Adding excessive indicator: Indicators are usually used in small quantities. Adding too much can alter the solution slightly or make the color change more difficult to interpret. Follow the prescribed number of drops.
Using an incorrect concentration for the titrant: The calculation is only as reliable as the known concentration. If the titrant has absorbed moisture, reacted with carbon dioxide from the air, or otherwise changed during storage, its actual concentration may differ from its label. Some titrants therefore require standardization, a process in which their concentration is determined using a suitable reference substance.
Recording too many or too few digits: Report results according to the precision of the measuring equipment and the reliability of the experimental procedure. Extra decimal places do not make a result more accurate.
Understanding accuracy and precision
Accuracy describes how close a measured value is to the true or accepted value. Precision describes how closely repeated measurements agree with one another.
A set of titration results may be precise but inaccurate if the same systematic error affects every trial. For example, an unsuitable indicator could produce similar endpoint volumes each time while consistently shifting the result away from the true equivalence point.
Random errors, such as small variations in drop size or subjective judgments of color, can cause repeated results to differ. Careful technique and repeated measurements help reduce the influence of these errors.
Systematic errors require a different response. They may result from incorrectly calibrated equipment, an inaccurate titrant concentration, contamination, or an unsuitable endpoint method. Repeating the experiment without correcting the underlying problem will not necessarily improve accuracy.
For important measurements, a laboratory may use a primary standard—a highly pure, stable substance suitable for preparing or verifying a solution of known concentration—to standardize the titrant. This helps establish a more reliable concentration for subsequent analyses.
Safety precautions during titration
Acid-base titrations are common teaching experiments, but the chemicals must still be handled responsibly. Strong acids and bases can cause serious skin and eye injuries, and even relatively dilute solutions may irritate tissue.
Wear splash goggles, suitable protective clothing, and gloves appropriate for the chemicals being used. Follow the laboratory’s instructions for handling, labeling, and disposing of solutions. Never pipette by mouth, and do not taste or directly smell laboratory chemicals.
If a chemical contacts your skin or eyes, immediately use the laboratory’s designated eyewash or safety shower as appropriate and follow the established emergency procedure. Notify the instructor or laboratory supervisor promptly. Do not attempt to neutralize a chemical on the skin with another chemical, because the reaction may generate heat or cause further injury.
Keep burettes securely mounted, handle glassware carefully, and clean up spills according to the laboratory’s safety procedures. Dispose of the resulting solutions as directed rather than assuming that every neutralized mixture is safe to pour down a drain.
Why acid-base titration is useful
Titration converts an observable laboratory measurement—the volume of a reacting solution—into a quantitative estimate of chemical concentration. This makes it useful whenever a substance’s acidity, alkalinity, or concentration needs to be measured reliably.
In water analysis, acid-base titrations can help determine alkalinity, which reflects water’s ability to neutralize added acid. In food science, titration can measure quantities such as titratable acidity in products including juices and other foods. In pharmaceutical and chemical manufacturing, suitable titration methods can help verify the composition or concentration of raw materials and finished products.
Titration also demonstrates an important principle of quantitative chemistry: chemical reactions follow predictable relationships between the amounts of substances consumed and produced. Once the reaction equation is known, a carefully measured volume can reveal information that cannot be determined reliably from appearance alone.
The most important habits for a beginner are to use clean equipment, measure volumes carefully, choose an appropriate indicator, slow down near the endpoint, repeat the experiment, and apply the balanced chemical equation correctly. With these practices, acid-base titration becomes a systematic and understandable way to determine an unknown concentration.