Introduction:

The goal of this experiment is to learn how to properly standardize an acidic solution of which we do not know the exact concentration. A secondary goal is to properly learn how to titrate a solution.

I hypothesize that by titrating the unknown solution with a standardized titrant, one can then use the data gathered by this to learn exactly how much titrant was needed to reach equilibrium, and many moles of the titrant were required to completely react with the solution of unknown concentration. By taking the stoichiometric ratios of these compounds into account, we can then calculate the concentration of the unknown solution by using the formula:

Equipment:

• Data Collection System
• pH Sensor
• Magnetic Stirrer
• Ring Stand
• Beaker (2), 100-mL
• Beaker (2), 10-mL
• Volumetric Flask, 1000-mL
• Erlenmeyer Flask, 250-mL
• Hot Plate/Magnetic Plate
• Automatic Titrator
• Funnel
• Jar with custom-made hockey puck-derived lid, 180-mL
• Hydrochloric Acid, unknown concentration (approximately 6M), 70 mL
• Sodium Hydroxide (NaOH) solution, standardized in Lab 6a, 100 mL
• Buffers, pH 4 and pH 10, 10 mL
• Water, deionized, 100 mL
• Plastic Bottle, 1-L
• Wash Bottle
• Pipet
• Aluminum Foil
• Tissue

Method:

• Gathered all equipment and materials. Inspected all equipment for damage or contamination. Hotplate/Magnetic Plate and Stand were already assembled.
• Obtained 20 mL of 6 M HCL stock solution. Pipetted 16.5 of the stock solution into a 1000 mL volumetric flask, and diluted solution to 1 L with deionized water.
• Sealed flask, inverted flask, shook gently to stir. Vented flask once solution had become homogenous. Transfered solution to 1 L plastic bottle. Labeled bottle.
• Started a new experiment with the data collection system. Set up graph.
• Set up the Titrator, flushing it once with deionized water.
• Calibrated the pH sensor, by rinsing it, measuring the pH of the buffer solutions, taking care to rinse the sensor between measurements, and then testing it with the buffers again.
• Added the magnetic stirrer to the end of the pH sensor, and passed it through the hockey-puck lid, and into the jar, ensuring that the entire apparatus was secured to the stand.
• Transferred, via pipet, 20 mL of the HCl solution to the 180 mL jar, and attaching it to the hockey puck lid. Ensured that the pH sensor was submerged within the solution 
• Loaded the Titrator with the NaOH titrant that was standardized in Lab 6a.
• Activated magnetic stirrer, and set to a mid-range speed.
• Activated the Titrator, running beyond the equilibrium point, until the pH curve had flattened, ensuring that the system was recording at the time.
• Repeated with a second and third HCl solution.
• Recorded data, saved graph, and flushed titrator with deionized water.
• Cleaned all glassware, set them to dry in their designated areas, and returned all equipment to its point of origin.

Data Analysis:

Average concentration of the HCl Solution: 0.09081 M

Conclusion:

The goal of this experiment is to learn how to properly standardize an acidic solution of which we do not know the exact concentration. By titrating the unknown solution with a standardized titrant, one can then use the data gathered by this to calculate exactly how much titrant was needed to reach equilibrium, and many moles of the titrant were required to completely react with the solution of unknown concentration. By taking the stoichiometric ratios of these compounds into account, one can then calculate the concentration of the unknown solution by using the formula:

Average concentration of the HCl Solution: 0.09081 M

As one began with an expected concentration of 0.099 M (given that one began with 16.5 mL of a 6 M solution, and diluted it to 1 L, with deionized water), given the resulting average experimental concentration of 0.09081 mol/L, this indicates error within the experiment that should be further examined.

Trial 1: 9.63%

Trial 2: 9.18%

Trial 3: 5.99%

Possible sources of error include:

• Failure to properly measure (or standardize) the concentration of the NaOH solution.
• Failure to titrate beyond the equivalence point (making determining the point at which the two solutes had completely reacted impossible).
• Failure to keep the solution mixed while titrating, introducing the possibility that the solution may not completely react in areas, and the possibility of inaccurate measurement.
• Failure to properly flush the titrator before beginning the experiment, either contaminating the solution, or diluting it with residual deionized water.
• Failure to properly handle the solutions, introducing the likelihood of contamination (NaOH can react with carbon dioxide in the air, while the HCl can leave solution, and return to its gaseous state), thus disrupting measurements.
• Human error is always in effect, given that the laboratory does not function under ideal conditions. As such, there is always the possibility of inaccuracies with measurement, perception of measurement, inaccuracies of equipment, and other such errors. (However, this is not likely to be the sole cause of the inaccuracies within this experiment, though it may contribute to it.)

Possible improvements that one could make to the experiment include using a more accurate pH probe, using more accurate balances (relating back to Lab 6a, and the standardization of the NaOH titrant), using a more accurate titrator (the syringe could be made more accurate), ensuring that lab partners read the experiment beforehand (increasing familiarity with the procedure, and minimizing human error), preparing the solutions in an atmosphere that lacks carbon dioxide (to avoid reaction of the NaOH solution), keeping the HCl solution covered (to minimize inaccuracies caused by it leaving solution) and repeating the experiment multiple times (to minimize the impact of an anomalous result).