Titration is a common laboratory method of quantitative chemical analysis used to determine the unknown concentration of a substance (the analyte) in a solution. This technique involves the controlled addition of a precisely measured reagent (the titrant) into the analyte until the reaction reaches completion. By monitoring the reaction, the exact amount of titrant required to fully react with the analyte can be measured, allowing for a precise calculation of the unknown concentration.
Preparation and Essential Equipment
Before beginning the experiment, all glassware must be meticulously cleaned and rinsed to ensure that residual contaminants do not dilute or interfere with the solutions. The apparatus centers around the burette, a long, graduated glass tube designed to dispense the titrant with high accuracy. This burette must be mounted securely on a stand, filled completely with the titrant, and the initial volume reading carefully recorded after ensuring no air bubbles are trapped.
The analyte is accurately measured using a volumetric pipette and transferred into an Erlenmeyer flask. The pipette is calibrated to deliver a fixed, precise volume of the analyte, making its accuracy paramount for the final calculation. The conical shape of the Erlenmeyer flask is specifically chosen because it allows the mixture to be swirled vigorously without the risk of splashing or losing any of the reacting solution.
The titrant is prepared in a volumetric flask to ensure its concentration is known with a high degree of certainty. The last preparatory step involves adding a few drops of a chemical indicator to the analyte in the Erlenmeyer flask. This indicator changes color abruptly when the chemical reaction between the titrant and analyte is complete, providing a visual signal for the end of the experiment.
The Step-by-Step Titration Procedure
The physical process begins by positioning the Erlenmeyer flask containing the analyte and indicator directly beneath the burette tip. The titrant is then slowly added from the burette while continuously swirling the flask to ensure the two solutions mix thoroughly and the reaction proceeds uniformly. Continuous mixing prevents localized concentration differences that could cause a premature or inaccurate color change.
To save time and estimate the required volume, a preliminary “rough” titration is often performed first, where the titrant is added relatively quickly. The goal of this initial run is simply to determine the approximate volume needed to cause the indicator to change color, which establishes a range for the precise trials. This rough estimate is discarded for calculation purposes but informs the technique for subsequent, more accurate runs.
The precise titration is performed by adding the titrant rapidly until the volume is within a few milliliters of the rough estimate, then slowing the addition significantly. The final milliliters are added drop-wise, allowing a moment for the color to fade with each drop while the flask is swirled. The endpoint is reached on the addition of a single drop that causes a persistent, permanent color change, signifying that the reaction is complete.
The volume of titrant used is determined by subtracting the initial burette reading from the final reading at the endpoint. The endpoint is the point where the indicator changes color, which is chosen to be as close as possible to the theoretical equivalence point. The equivalence point is the exact moment when the moles of titrant added are stoichiometrically equal to the moles of analyte present in the flask.
Calculating and Analyzing Results
The mathematical analysis begins by accurately calculating the net volume of titrant delivered from the burette, which is the difference between the final and initial readings. This precise volume, along with the titrant’s known concentration, allows for the determination of the moles of titrant used. The fundamental relationship used here is that Moles equals Molarity multiplied by Volume in liters.
For example, in a common acid-base titration like hydrochloric acid (\(\text{HCl}\)) and sodium hydroxide (\(\text{NaOH}\)), the balanced chemical equation shows a 1:1 mole ratio. This means that at the equivalence point, the moles of \(\text{NaOH}\) titrant added are exactly equal to the moles of \(\text{HCl}\) analyte initially in the flask. The moles of titrant calculated in the previous step are therefore directly equivalent to the moles of the unknown analyte.
If the reaction had a different stoichiometry, such as a 1:2 ratio, the moles of analyte would be calculated by applying the appropriate conversion factor from the balanced equation. After the moles of the analyte are determined, the final step is to calculate its concentration, or molarity. This is achieved by dividing the calculated moles of analyte by the original, fixed volume of the analyte that was measured into the Erlenmeyer flask.
Since molarity is defined as moles per liter, the initial volume of the analyte must be converted from milliliters to liters before performing the division. This sequential process moves from measured volume to moles of known substance, then to moles of unknown substance using the mole ratio. Finally, the concentration of the unknown is determined, providing a highly accurate quantitative result for the analyte.