Determination of the Concentration of a Monoprotic Acid
Purpose:
In this lab, you will be working on your own to determine the concentration of a strong monoprotic acid,
“the unknown.” This lab will be conducted in two parts. In part I, you will determine the exact molarity (to
four decimal places) of a NaOH solution by standardizing the NaOH against a primary standard, potassium
hydrogen phthalate (KHP = KC8H5O4, 1 mole = 204.22 grams). This NaOH solution will be used in part II
to titrate the solution containing your monoprotic acid of unknown concentration. In this experiment, the
major portion of your grade will be based on the accuracy and precision of your results, i.e., how close you
come to the actual concentration of your unknown strong acid.
Background:
Titration is a technique that chemists use to determine the unknown concentration of a known solution (we
know what chemical is dissolved, but not how much in a solution). Because we know what the chemical is,
we know how it will react with other chemicals and we can use that reaction to determine the concentration
of the solution by measuring the formation of product(s). In the case of an unknown concentration of acid,
we can use a known concentration of hydroxide base. This type of reaction is a neutralization reaction,
where salt and water are products of the reaction:
Acid + Base → salt + H2O
We can use a pH indicator, a chemical that changes color depending on the pH, to show us when the
reaction has completely neutralized. The point, where all acid was consumed and there is no excess base, is
called the equivalence point. We can use this equivalence point to determine the initial concentration of
acid using a series of calculations. The goal of the titration is to get as close as possible to the equivalence
point by careful addition of the base; this will ensure the calculated acid concentration is as close to the true
value as possible. The terms below will help you understand the terminology used throughout the
experiment:
· Titrant—the solution of known concentration, in this lab, the titrant is sodium hydroxide.
· Buret—a long, cylindrical piece of glass that can be used to determine small, accurate quantities of a
solution. A buret is controlled by a stopcock, a white Teflon piece that can be turned to deliver the solution.
The markings on the buret are such that you must subtract the initial reading (where the titrant level is
initially) from the final reading to determine the volume of base delivered. The buret measures 2 digits
after the decimal point accurately.
· Volumetric pipette/pipette bulb—a thin glass tube with only one marking used to measure a very specific
volume of liquid. You will use a pipette bulb to pull the liquid into the pipette.
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· Phenolphthalein—a pH indicator. In acidic and neutral solutions, the indicator is colorless, but in a basic
solution, the color is a vibrant pink. The high the pH is, the stronger the pink color is. The end point will be
when the color is a very faint pink color. Keep your flask with acid and indicator over a white piece of
paper to ensure you can see the color change.
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· End point versus equivalence point - An end point is a color change that indicates when the right amount
of titrant is added. The end point is observable. The equivalence point is when the stoichiometric amount of
titrant is added to the analyte. In an acid-base titration, it is when an equal number of moles of H+ and OH–
react.
Apparatus – Part I
Analytical balance (min. 4 decimal places)
Pure KHP (pre-dried)
Clean 125 mL Erlenmeyer flask
Clean 50 mL buret
Weighing paper
Small funnel
Wash bottle filled with DI water
~ 50 mL of 0.1 M NaOH solution (dispensed by
instructor)
Phenolphthalein indicator
Buret clamp
Ring stand
Clean spatula
Plastic funnel
Procedure Part 1 : Preparation of Standardized NaOH Solution
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In part I, you will determine the concentration of a sodium hydroxide solution to a high degree of accuracy.
This process is called standardization and the resulting solution is a standard solution. That is, a standard
solution is one having an accurately known concentration.
In order to determine the concentration of the sodium hydroxide solution, one must have an especially pure
acid so that an accurately measured amount of acid can be weighed out on the analytical balance. The
weight of this acid is the starting point for all subsequent calculations and it is therefore called the primary
standard. In general, a primary standard is any especially pure chemical that can be used as the starting
point to quantify an analysis. Few chemicals are pure enough and stable enough to be used as primary
standards. For example, solid sodium hydroxide cannot be used as a primary standard because it absorbs
atmospheric moisture and carbon dioxide during storage and also during a weighing operation. A primary
standard should have the following qualities:
a. It must be easily prepared, purified and dried.
b. It must be stable and easily stored.
c. So it can be weighed in open air, it must not be hygroscopic. It must not react with any of the
components of air such as carbon dioxide, oxygen or water.
d. Suitable methods must be available to test it for impurities. Generally, the total impurities must be less
than 0.01-0.02%. The exact assay (i.e., the percent purity) must be known.
e. The reaction for which the primary standard is to be used must be quantitative and must be fast enough
that it goes to completion in a reasonable period of time. To determine the concentration of a sodium
hydroxide solution through a titration, the primary standard must be an acid. In the present experiment,
potassium hydrogen phthalate (KHP = KC8H5O4 – 204.22 g/mole) will be used.
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From the stoichiometry of the reaction one can see that one mole of KHP reacts with one mole of OH- ions.
Procedure for Standardization of NaOH
1. Prepare data tables and obtain the necessary glassware. 1 Wash all glassware with hot soapy water, rinse
thoroughly with tap water, and then rinse 3-4 times with deionized water.
2. Condition the buret with the ~ 0.1 M NaOH solution and check that the buret is flowing freely. To
condition a piece of glassware, rinse it so that all surfaces are coated with solution, then drain. Conditioning
at least two times will insure that the concentration of titrant is not changed by a stray drop of water.
3. To fill the buret, close the stopcock at the bottom and use a funnel. You may need to lift up on the funnel
slightly, to allow the solution to flow in freely. Fill the buret with NaOH solution.
4. Check the tip of the buret for an air bubble. To remove an air bubble, gently tap the side of the buret tip
while solution is flowing. If an air bubble is present during a titration, volume readings may be in error.
6. Go to the analytical balance with your flask, funnel and bottle of KHP. Place a piece of creased
weighing paper on the balance pan. Place the tare button to cancel out the weight of the weighing paper.
Carefully add at least about 0.26 and not more than 0.35 grams of KHP to the balance pan. Close the glass
door and record the mass of KHP to a minimum of 4 places.
7. Quantitatively transfer the KHP to the clean 125 mL flask. This can be done be carefully tipping the
creased weighing paper to pour the solid into a plastic funnel inserted at the top of the flask. Tapping the
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Refer to pages 6 & 7 for sample data tables.
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5. Set-up the ring stand with buret clamp and secure buret to clamp.
paper with a spatula will knock particles into the funnel and the flask. Finally, the paper and funnel should
be rinsed with DI water to ensure that all KHP is transferred to the flask.
8. Return to your bench and add about 50 mL of DI water to the beaker. Using a clean glass stirring rod
gently stir the KHP until it completely dissolves (this may take a couple of minutes).
9. Add 1 drop of phenolphthalein indicator to the KHP solution. Stir the solution until the indicator is
completely solubilized in the KHP solution.
10. When your buret is conditioned and filled, with no air bubbles or leaks, take an initial volume reading.
Read the bottom of the meniscus. Be sure your eye is at the level of meniscus, not above or below. Reading
from an angle, rather than straight on, results in a parallax error.
11. Begin the titration. Place the flask containing the acid solution and indicator under the buret. Add
NaOH from the buret to the flask with swirling until the color of the solution in the flask is a faint pink.
This faint pink color should last only 45 to 60 seconds. There should be a one-drop difference between
when the solution is colorless and when it is pink. If too much base is added (that is, if you "over-shoot" the
endpoint), discard the solution and repeat the titration. A white piece of paper placed under the flask will
aid in the color detection. When you have reached the endpoint, read and record the final NaOH volume to
within + 0.02 mL.
12. Subtract the initial volume to determine the amount of NaOH solution delivered. Use this, the mols of
KHP, and the stoichiometry of the titration reaction to calculate the concentration of the NaOH solution.
13. Discard the contents of the Erlenmeyer flask into the waste container located under the hood.
14. Repeat the titration procedure a second time by following steps 6 - 15.
15. Calculate the molarity of the NaOH from the two titrations. If the calculated base concentrations from
the first and second titration vary by more than 2%, perform a third titration.
Apparatus – Part II
Clean 125 mL Erlenmeyer flask
15 or 20 mL volumetric pipet
50 mL unknown monoprotic strong acid
Clean 50 mL buret
Ring stand
Buret clamp
Small clamp (for securing pH probe)
Glass stirring rod
Small funnel
Wash bottle with DI water
Phenolphthalein indicator
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Standardized NaOH solution (from Part I)
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Procedure for Titration of Strong Acid
In part II, you will use your standardized NaOH solution from part I to titrate the solution containing your
monoprotic, strong acid, of unknown concentration.
1. Condition the buret with the ~ 0.1 M NaOH solution and check that the buret is flowing freely. Refill and
burette and record the NaOH level to within + 0.02 mL on the data sheet.
2. Obtain a clean, but not necessarily dry, 125 mL Erlenmeyer flask and using a volumetric pipette, transfer
15.00 mL of your unknown acid solution into the flask. Add 1 drop of phenolphthalein indicator to the HCl
solution in the Erlenmeyer flask.
3. Place the flask containing the acid solution and indicator under the buret and add NaOH from the buret
to the flask with swirling until a phenolphthalein endpoint is reached. There should be a one-drop
difference between when the solution is colorless and when it is pink. If too much base is added (that is, if
you "over-shoot" the endpoint), discard the solution and repeat the titration.
5. When the proper end point is reached, read and record the final NaOH volume to within + 0.02 mL.
6. Discard the contents of the Erlenmeyer flask into the waste container located under the hood. Repeat the
titration procedure a second time by following steps 2 - 5.
7. Calculate the molarity of the unknown mon0protic acid for each titration. If the calculated acid
concentrations from the first and second titrations vary by more than 2%, perform a third titration.
Calculate an average acid molarity using the two closest values.
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8. Before returning the buret to the lab bench, please rinse it out with a couple of water rinses.
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Report Sheet – Part I
Unknown Number (Base) ________________
Name_______________________________
Data Table for Standardization of Base
Trial 1
mL NaOH, final reading 2
2
mL NaOH, initial reading
3
mL NaOH to reach endpoint
4
Liters NaOH to reach endpoint
5
Mass of KHP 3
6
Molar mass of KHP
7
Mols KHP
8
Mols of NaOH (same as mols KHP)
9
Calculated NaOH molarity
10
Average NaOH molarity
Trial 3
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1
Trial 2
2
3
Record buret volume readings to 2 decimal places
Record mass to 4 decimal places
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Report Sheet – Part II
Unknown Number (Acid) ________________
Data Table for Titration of Acid with Standardized NaOH (Base)
Trial 1
mL NaOH, final reading
2
mL NaOH, initial reading
3
mL NaOH added to reach endpoint
4
Liters NaOH to reach endpoint
5
Concentration of NaOH (from Part I)
6
Mols NaOH to neutralize acid
7
Mols of acid (same as mols NaOH)
8
Initial volume of acid (Liters) in flask
9
Calculated molarity of acid unknown
10
Average molarity of unknown acid
Trial 3
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1
Trial 2
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Extra Space for Calculations.
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