LAB 6:
DETERMINATION OF THE MOLAR MASS OF AN ACID BY TITRATION
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This lab requires a lab report.
You will work with a partner for this lab.
I Introduction
Salicylic acid, an ingredient in the synthesis of aspirin, is a monoprotic acid,
meaning that it can dissolve and release only one H+ ion (thus the mono- in
monoprotic). Salicylic acid is an organic compound, meaning that it’s made of carbon,
hydrogen, and other possible elements (in this case, oxygen). In other words, it has a
slightly more complex formula than our traditional HCl, HNO3, etc.
Like all acids, salicylic acid reacts with bases, such as sodium hydroxide
(NaOH). Since we do not know the formula of salicylic acid, let us represent it by the
generic formula HA, where H is the removable H+ ion and A is the rest of the acid. We
can then write the following balanced equation:
HA(aq) + NaOH(aq) → H2O(l) + NaA(aq)
In this lab, you will titrate a solution containing a known mass of salicylic acid
with a 0.10 M solution of sodium hydroxide. Using the volume of NaOH used, its
molarity, the mass of acid used, and some simple stoichiometry, you will determine
the molar mass of salicylic acid.
Although reasonable results can be achieved by using the concentration of
NaOH written on the bottle (i.e. 0.10 M) in your calculations, even better results are
obtained when the NaOH solution is standardized, or reacted with an acid solution of
known concentration so as to calculate the concentration of NaOH. Standardization of
NaOH is especially important because solid sodium hydroxide is hygroscopic: it
absorbs moisture from the air. Thus, when the 0.10 M solution of NaOH is made, solid
NaOH is used that contains an unknown amount of water. This makes the actual
sodium hydroxide concentration in the prepared solution less than 0.10 M.
To standardize our solution of NaOH, we will use the monoprotic acid
potassium hydrogen phthalate (KHC8H4O4, often abbreviated as KHP). KHP is a nonhygroscopic solid, and so solutions of exact KHP concentration can be prepared quite
easily.
Equipment and Reagents
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Buret and ringstand
Graduated cylinder
250-mL Erlenmeyer flask
100-mL beaker
Stirring rod
Magnetic stirrer and stirbar
Analytical balance
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KHP
~ 0.10 M NaOH solution
Salicylic acid
Ethanol (C2H5OH), 95%
Distilled water
Phenolphthalein indicator
Litmus paper
! Warnings!
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Though in relatively mild concentrations, the acids and bases used in this lab
ARE harmful. Wear your goggles and aprons. If you spill any substance on
yourself, rinse it off with plenty of water. If you spill any substance on the
tabletop, immediately clean it up with wet paper towels, Acid-Lok®, or Alky®,
as appropriate.
Procedure
Part 1: Standardizing the NaOH
Preparing the Acid
1) Prepare your lab notebook.
2) Measure out about 0.5 g of KHP. Record the exact amount of KHP used.
3) Transfer ALL of the KHP into a 250-mL Erlenmeyer flask and add about 25
mL of distilled water to dissolve.
4) Test the acid with litmus paper and a stir rod; record your observations.
5) Add 2-3-4 drops of phenolphthalein indicator.
Your KHP solution is now ready for titration.
Preparing the Titrant (base solution)
1) Rinse out your buret with a few milliliters of the base solution: pour some in,
rotate the buret around to coat all the surfaces, and then rinse your stopcock
by opening it and allowing the base to drain into the sink. Flush the base down
the sink with running water. Do NOT rinse your buret with water.
2) Fill your buret to the 0.0 mL mark with your base solution. (Make sure your
stopcock is closed before you start to pour!)
Your base solution is now ready for titration.
Performing the Titration
1) Begin your titration by opening the stopcock; do not remove your hand from it.
Once you begin to see a lingering color change in the analyte (acid) solution,
slow down the rate of flow until the titrant is added drop by drop. Once the
color change is persistent, you have reached the endpoint (and equivalence
point) of your titration. (NOTE: Do NOT empty your buret past the 50.0 mL
mark! Instead, stop at the 50.0 and add more NaOH if you need to). Record
your observations.
2) Record the final volume of base in the buret.
3) Using the mass of KHP and volume of NaOH used, calculate the concentration
of the NaOH solution. Record your answer on the board for class comparison.
4) Flush your analyte solution down the drain with running water and rinse out
your flask. Leave the buret and its contents alone.
Part 2: Titrating the Unknown Acid
Preparing the Acid
1) Measure out about 0.5 g of salicylic acid. Record the exact amount of acid
used.
2) Add about 10 mL of ethanol to the acid and dissolve with stirring (feel free to
use the magnetic stirrer). Then add about 10 mL of distilled water and stir.
Why ethanol? Ethanol’s not JUST for burning; it’s also a respectable solvent. Although water is
considered the “universal solvent”, it can’t dissolve everything (think oil). To dissolve some
things (such as salicylic acid), you need a less polar solvent. Ethanol is nonpolar enough to
dissolve the acid but polar enough to mix well with the water added in the NaOH solution.
3) Add 2-3-4 drops of the phenolphthalein indicator. Turn on your magnetic
stirrer.
Your acid solution is now ready for titration.
Preparing the Titrant (base solution)
1) Prepare the buret of NaOH solution as described in Part 1. Note that there is no
need to rinse the buret.
Your base solution is now ready for titration.
Performing the Titration
1) Perform the titration as described in Part 1, recording all necessary data and
observations.
2) Using the concentration of NaOH, the volume of NaOH used, and the mass of
the acid used, calculate the molar mass of salicylic acid. Record your result on
the board for class comparison.
3) Flush your analyte solution down the drain with running water. Leave your
buret and its contents alone.
Data and Calculations (to be recorded on your lab Data Sheet)
Part 1: Standardizing the NaOH
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Title and Purpose
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Procedure used
Masses and volumes of all reagents used
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Observations of all tests and reactions
Moles of KHP used (MM of KHP = 204.22 g/mol)
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Concentration of NaOH and all supporting calculations
Class average of [NaOH]
Part 2: Titrating the Unknown Acid
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Procedure used
Masses and volumes of all reagents used
Observations of all tests and reactions
Balanced molecular equation for the reaction
Molar mass (in units of g/mol) of salicylic acid and all supporting calculations
Class average of molar mass
? Questions to Answer
1) Elemental analysis of salicylic acid reveals that it is 60.9% C (by mass), 4.39% H (by
mass), and the rest oxygen. What are the empirical and molecular formulas of salicylic
acid?
2) Why do you not want to drain the NaOH past the 50-mL mark during the titration?
3) Why does it NOT matter how much ethanol and water you use to dissolve the acid?
4) What was the percent error of your molar mass determination? What are some possible
sources of error that you would correct/prevent if you were to repeat the exercise?
5) How did your percent error compare to the percent error of the class average? What
are some advantages and disadvantages of using a class average?
6) Oops! You rinsed your buret with water instead of NaOH. Some water remained in the
buret when you filled it with NaOH.
a. Is your calculated moles of NaOH added too big, too small, or unchanged?
Explain.
b. Is your calculated molar mass of the acid too big, too small, or unchanged?
Explain.
7) Oops! You overshot your titration and used too much NaOH.
a. Is your calculated moles of NaOH added too big, too small, or unchanged?
Explain.
b. Is your calculated molar mass of the acid too big, too small, or unchanged?
Explain.