Introduction:

The goal of this experiment is to determine the concentration of a known solution, by titrating it with another strong electrolyte that will react with the first substance, and enter a non-ionic state. By doing this, we change the number of ions in the solution, and thus change its conductivity. One can accomplish by reacting a Barium Hydroxide solution of an unknown concentration with a titrant of Sulfuric Acid of known molarity. Then, by measuring the conductivity of the solution as one continues to add titrant to, and beyond the equivalence point, we can locate the point at which the conductivity is lowest, due to the lack of ions, as those from both compounds will have reacted completely. Then, as we know the chemical composition of both compounds, we can then calculate their stoichiometric ratio, and then use that, coupled with the known molarity of the Sulfuric Acid solution, to calculate the molarity of the Barium Hydroxide solution.

Equipment:

• Data Collection System
• Conductivity Sensor
• Magnetic Stirrer
• Ring Stand
• Clamp
• Beaker (2), 100-mL
• Beaker, 50-mL
• Hot Plate/Magnetic Plate
• Automatic Titrator
• Jar with custom-made hockey puck-derived lid, 180-mL
• Hydrochloric Acid, unknown concentration (approximately 6M), 70 mL
• 0.0200 M Sulfuric Acid, 50-mL
• Barium Hydroxide, Unknown Concentration, 50-mL
• Water, deionized, 50 mL
• Volumetric Pipet, 50-mL
• Wash Bottle
• Tissue

Method:

• Gathered all equipment and materials. Inspected all equipment for damage or contamination.
• Set up Data Collection System, assembled GLX, and conductivity sensor.
• Set up Automatic Titrator, ring stand, Hot Plate/Magnetic Plate, and attached the Jar with custom-made hockey puck-derived lid.
• Placed the conductivity sensor so that it would be suspended in the jar, and attached the magnetic stirrer to the end of the sensor.
• Flushed Automatic Titrator with deionized water, to minimize contamination, before filling with 50 mL of 0.0200 M Sulfuric Acid.
• Acquired 50 mL of Barium Hydroxide, and poured into Jar.
• Added deionized water until the solution covered the tip of the conductivity sensor.
• Turned on Magnetic Stirrer, checked all connections to ensure that they were secure, checked the Automatic Titrator to ensure that no bubbles were present within it, then began titrating.
• Recorded the Data via the GLX, generating the graphs used.
• Disposed of solutions in designated containers. Cleaned all glassware, set them to dry in their designated areas, and returned all equipment to its point of origin.

Data Analysis:

Trial 1:

Trial 2:

The concentration of the Barium Hydroxide solution was:

Trial 1: 0.00385 M

Trial 2: 0.00370 M

Conclusion:

The goal of this experiment is to determine the concentration of a known solution, by titrating it with another strong electrolyte that will react with the first substance, and enter a non-ionic state. By doing this, we change the number of ions in the solution, and thus change its conductivity. One can accomplish by reacting a Barium Hydroxide solution of an unknown concentration with a titrant of Sulfuric Acid of known molarity. Then, by measuring the conductivity of the solution as one continues to add titrant to, and beyond the equivalence point, we can locate the point at which the conductivity is lowest, due to the lack of ions, as those from both compounds will have reacted completely. Then, as we know the chemical composition of both compounds, we can then calculate their stoichiometric ratio, and then use that, coupled with the known molarity of the Sulfuric Acid solution, to calculate the molarity of the Barium Hydroxide solution.

The concentration of the Barium Hydroxide solution was:

Trial 1: 0.00385 M

Trial 2: 0.00370 M

(Calculation of error not possible, due to a lack of a provided concentration for the Barium Hydroxide)

Possible sources of error include:

• Failure to properly measure the volumes of the solutions used.
• 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, thus disrupting measurements. 
• Failure to remove all bubbles from the Automatic Titrator, thus disrupting the measurement of volume for the Sulfuric Acid.
• 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 conductivity probe, using more accurate pipets, 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), and repeating the experiment multiple times (to minimize the impact of an anomalous result).