Purpose:
The purpose of this lab is to determine the molar mass of solid tartaric
acid. To do this, a solution of sodium hydroxide will first be standardized
against a weighed sample of potassium hydrogen phthalate (KHP), an acid
whose mass and composition is consistent over time. The sodium hydroxide
will then be used to titrate a solution of tartaric acid so that the molar mass
of tartaric acid can be accurately calculated.
Background:
An acid/base titration is performed by carefully adding one solution from
a buret, the titrant, to another substance in a flask until all of the substance in
the flask has reacted. To show when the titration is complete, a dye called an
indicator is added to the solution to show when the end point, or the color
change indicated when chemical equilibrium between moles of the acid and
moles of the base is reached. In the titration performed in this experiment, the
reactant NaOH must be standardized against a primary standard such as KHP
because it is a compound that changes chemical composition in storage. Solid
NaOH absorbs water vapor so that the mass of pellets of NaOH is partly NaOH
and partly water. Thus, it is impossible to know the exact molarity of any
solution of NaOH without standardizing it by titration because of the variable
water content that will distort and invalidate a mass measurement. A heat-dried
sample of KHP can be relied on to be 100% KHP; thus, an accurately determined
mass of KHP will reveal an accurate number of moles of sodium hydroxide.
Materials
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Potassium hydrogen
phthalate (KHP), 204. g/mol
NaOH solution (approx.
0.100M)
Tartaric acid
Phenolphthalein (Phth)
indicator
Buret and clamp
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250ml Erlenmeyer flask
Graduated cylinder
Distilled water bottle
Analytical balance
Magnetic stirrers
Bottle container
Procedure:
Part 1: NaOH and KHP
1) Obtain a clean buret washed with soap and rinsed with distilled water.
2) Calculate mass of NaOH in a 0.1M solution of NaOH, and obtain the
calculated mass using weigh boats. Be sure to “tare” the balance with
only the weigh boat on it before adding the NaOH.
3) Put massed NaOH in a Florence flask, and add water to the 1L mark.
4) Place a magnet into the flask, and the flask on the stirrer. Turn on the
stirrer and dissolve the NaOH.
5) Pour the solution into a bottle container.
6) Measure out 0.5g-1.0g of KHP, and record the exact mass. Put the KHP in
an Erlenmeyer flask, add about 100mL of distilled water, and dissolve it
using the magnet stirrers.
7) Add 2-3 drops of Phenolphthalein indicator into the Erlenmeyer flask.
8) Rinse the buret with a small portion of the NaOH solution. Clamp it so
that the buret will drip into the Erlenmeyer flask. Then, using a funnel, fill
the buret with the solution to the “0” mark, making sure not to let
solution spill over the top.
9) Calculate the amount of NaOH theoretically needed to reach the
equivalence point, and add 10mL less than that from the buret to the
flask.
10) Open the buret slightly so that NaOH will drip slowly.
11) When the color of the solution in the flask turns slightly pink, tighten the
buret, calculate the amount of NaOH used, and record the data.
12) Calculate the molarity and label bottle container containing NaOH
solution.
13) Repeat titration twice.
Part II: NaOH and Tartaric Acid
14) Measure 0.1g-0.2g of tartaric acid, and record exact mass.
15) Add about 100mL of water, the measured tartaric acid, 2-3 drops of Phth
indicator, and a magnet to a clean Erlenmeyer flask.
16) Dissolve the acid.
17) Record the initial volume of NaOH in your buret. Titrate the tartaric acid
solution with the standardized NaOH. Record the final volume of NaOH.
Calculate and total volume of NaOH used in the titration.
18) Calculate moles of tartaric acid in your sample as well as the molar mass
of tartaric acid in g/mol.
19) Repeat titration until results are within 10% of each other.
Conclusion:
The molar mass of solid tartaric acid is 150.0g/mol. This mass was
found by first multiplying the volume of NaOH required to reach the end
point (avg. of trials was 21.34mL) and the molarity of the NaOH solution
(calculated by dividing moles of KHP by volume of NaOH needed to
complete the titration) to calculate moles of NaOH (0.0974M). Since the
molar ratio of tartaric acid to NaOH is 1:2, dividing the calculated moles of
NaOH by 2 yielded the moles of tartaric acid in the sample (avg. of trials was
0.00104). Lastly, dividing the measured mass of tartaric acid (0.155g) by the
calculated moles of acid yielded the molar mass of tartaric acid. The average
of the three trials showed the molar mass of tartaric acid to be 150.0g/mol.
The calculations used in this experiment largely relied on stoichiometry and
the definition of molarity (moles divided by liters) to allow the manipulation
of data for the desired information – molar mass of tartaric acid. The
chemical process that occurred during the titration of NaOH and tartaric acid
involved hydroxide (OH-) ions reacting with the hydrogen (H+) ions floating
in the solution of dissolved H2C4H4O6 in the beaker. So, for every mole of
acid 2 moles of hydrogen ions reacted with the hydroxide ions dissolved in
the solution, as tartaric acid is a diprotic acid. When the equivalence point
was reached, signifying chemical equilibrium between moles of H2C4H4O6
acid and moles of the NaOH base, the end point, or the pink color change,
indicated that the titration was complete.