Determining the Concentration of Citric Acid in 7-up using Acid/Base Titration
Background:
The acid content of many foods and beverages contributes significantly to their taste. Soft
drinks often contain varying quantities of several acids, which give sodas their tart flavor. In
cola products, these acids are predominantly carbonic acid (from the carbonated water) and
phosphoric acid. In sodas such as Squirt and 7-Up, the acids are carbonic acid and citric acid.
Acids can be classified as monoprotic, diprotic, triprotic, etc. A monoprotic acid has one
proton that can undergo a reaction with a base, a diprotic acid has two such protons, and a
triprotic acid has three. Shown below are examples of each type of acid undergoing a reaction
with sodium hydroxide.
Monoprotic: HCl(aq) + NaOH(aq) → H2O(l) + NaOH(aq)
Diprotic: H2SO4(aq) + 2NaOH(aq) → H2O(l) + Na2SO4(aq)
Triprotic: H3PO4(aq) + 3NaOH(aq) → H2O(l) + Na3PO4(aq)
Any acid that has more than one proton that undergoes a reaction with a base is called a
polyprotic acid. Citric acid is a weak, polyprotic acid that undergoes the following reaction.
H3C6H5O7(aq) + 3 NaOH(aq) → 3 H2O(l) + Na3C6H5O7(aq)
In this experiment you will be performing a titration to determine the concentration of citric
acid in 7-Up. Prior to the titration the majority of the carbonic acid was removed by allowing
the 7-Up to go flat so we do not have to take it into consideration. A burette is used in a
titration to dispense measured increments of one solution into a known volume of another
solution. Careful technique will allow you to detect the point where the reaction is complete; in
this case, when all of the citric acid has been reacted with the base. The technique of titration
can be applied to other types of reactions such as oxidation-reduction, precipitation,
complexation and other acid-base neutralization reactions.
Indicators
Indicators are used to provide visual information about the amount of acid or base in a solution.
You will be using an indicator called phenolphthalein for this experiment. In a basic solution (a
solution that contains more base than acid), phenolphthalein is pink. In solutions that are acidic
(a solution that contains more acid than base), phenolphthalein is colorless. During the titration
performed in this experiment, the point where the indicator first changes from colorless to pink
corresponds to the point where the reaction is complete. This point is called the endpoint of
the titration.
Preparing a standard base
Prior to performing the titration to determine the amount of acid in a 7-Up sample, you need to
know the exact concentration of the base that you will be using. Standardizing the base
involves performing a titration with a known amount of a stable acid, which in this case is the
monoprotic acid potassium hydrogen phthalate (abbreviated KHP). The reaction of KHP and
NaOH is
NaOH(aq) + KHC8H4O4(aq) → H2O(l) + KNaC8H4O4(aq)
Note that one mole of NaOH reacts with one mole of KHP. Note that the chemical formula of
KHP is not KHP! It is KHC8H4O4
Standardization of NaOH
You will first need to make ~250 mL of a ~0.020 M NaOH solution. You will then standardize it
by titrating a known mass of KHP (~0.05 g) in ~25 mL of water into a 125 mL flask and don't
forget some indicator. Your end point is a faint persistent pink color. Placing a white piece of
paper under your Erlenmeyer flask helps you better see this color change. Calculate the
molarity of the NaOH solution. Repeat this process two more times and calculate the average
molarity of the NaOH solution. Formulas have been provided
1) In a 250 ml Erlenmeyer flask place a mass of ~ 0.2 g of NaOH and dissolve in ~250 mL of
Distilled Water (This will be your ~0.020 M solution of NaOH)
a. To assist in the dissolving process place stopper in flask and invert several times
2) Use funnel and rinse burette (twice) with your NaOH solution.
a. This cleans burette of previous contaminants
3) Set up titration as shown by instructor
4) Fill burette with NaOH solution (use funnel, make sure that the stopcock is closed)
5) Place a clean 125 mL Erlenmeyer flask on balance and Zero the balance.
6) Add ~ 0.05g of KHP to the flask and record mass in data table. Add ~25 mL of Distilled
Water and dissolve KHP.
a. MAKE SURE TO ADD 2-3 DROPS OF PHENOLPHTHALEN
7) Record initial volume from burette in data table and begin titration
8) Slowly add NaOH solution, stop and swirl occasionally. When it remains pink,
neutralization has been reached. (record final burette reading in data table)
9) Repeat steps (5-8) for trial #2 & #3
10) Complete data table (FORMULAS HAVE BEEN PROVIDED)
Standardization of NaOH
Trial #1
Trial #2
Trial #3
Mass of KHP (g)
Final burette reading NaOH (mL)
Initial burette reading NaOH (mL)
Total Volume NaOH used (mL)
(Final Volume – Initial Volume)
Molarity of NaOH
((Mass of KHP/204.22)/ Total Volume NaOH used)X1000
Average Molarity of NaOH Solution [NaOH]
Titration of 7-Up
Now that you have your [NaOH], it is time to titrate the soda. The citric acid content in Squirt
and 7-Up is quite low, so good technique is critical. You will be titrating 10 mL samples of the
supplied soda samples. The sodas in your lab have been opened days in advance to assure they
are decarbonated. Slowly add the NaOH solution to the 7-Up until a faint pink color is obtained.
Calculate the molarity of the citric acid in 7-Up. Repeat this process two more times and
calculate the average molarity. I UNDERSTAND WE HAVE NOT BEEN INTRODUCED TO SOME
OF THE CONCEPTS NEEDED FOR CALCULATION; THEREFORE I HAVE PROVIDED FORMULAS
FOR THESE CALCULATIONS WITHOUT EXPLANATION.
1)
2)
3)
4)
5)
6)
7)
8)
Refill your burette with NaOH solution (Using a funnel)
Dispense ~50 mL of 7-Up into a 100 mL beaker
Measure ~10 mL of 7-Up into a graduated cylinder and record volume in data table
Dispense 7-Up into a clean 125 mL Erlenmeyer flask and add 2-3 drops of
pheoolphthalen
Begin titration (record initial volume from burette in data table)
When solution remains pink, neutralization has been reached. (record final burette
reading in data table)
Repeat steps (3-7) for trial #2 & #3
Complete data table (FORMULAS HAVE BEEN PROVIDED)
Titration of 7-Up
Trial #1
Trial #2
Trial #3
Volume of 7-Up (mL)
Final burette reading NaOH (mL)
Initial burette reading NaOH (mL)
Total Volume NaOH used (mL)
(Final Volume – Initial Volume)
Molarity of citric acid
(Vol. NaOH X Average molarity of NaOH X 3)/Vol. 7-Up
Average molarity of citric acid
Average moles of citric acid in can of 7-Up (355 mL)
(Average molarity of citric acid X 0.355)
Average grams of citric acid in can of 7-Up (355 mL)
(Average moles of citric acid in can of 7-Up X 192.12)
1) We titrated the base into the acid and not the acid into the base? Could you perform
the experiment the other way? Titrating an acid of known concentration into a base of
unknown concentrations? Explain. (Think of the indicator you would use)
2) In the calculation for Molarity of Citric acid you notice that I have you multiply by 3 in
the numerator. Yet it the preparation of the standard we did not. Offer an explanation.
(Hint: look at the balanced equations for the two reactions!!!)
3) Why was it essential that we standardized the base (needed to know the EXACT
concentration) before we used it to titrate our citric acid?