CHEM 2124
General Chemistry II Laboratory
Week 3 – DETERMINING THE CONCENTRATION OF CITRIC ACID IN
SOFT DRINKS VIA ACID/BASE TITRATION (MAJOR LAB #1)
OBJECTIVE
In this laboratory exercise, the student will determine the amount of citric acid in various soft drinks.
This will be determined via acid-base titration with the use of phenolphthalein indicator.
INTRODUCTION
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 citrus
flavored sodas such as Sprite, Sierra Mist, and 7-Up the acids are carbonic acid and citric acid.
As we have discussed in class, acids can be classified as monoprotic, diprotic or triprotic depending on
the number of ionizable hydrogen ions. Examples of each class of acid undergoing a neutralization
reaction with sodium hydroxide are shown below.
Monoprotic Acid:
Diprotic Acid:
Triprotic Acid:
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
H2SO4 (aq) + 2 NaOH (aq) → 2 H2O (l) + Na2SO4 (aq)
H3PO4 (aq) + 3 NaOH (aq) → 3 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, with the chemical formula of H3C6H5O7 is a weak, polyprotic acid that undergoes a
neutralization reaction with sodium hydroxide to form water and a salt.with sodium hydroxide.
As you know, soft drinks are carbonated or “fizzy.” The chemical responsible for this carbonation is
carbonic acid which decomposes into H2O and CO2. Prior to the laboratory exercise, the majority of
the carbonic acid will have been removed by allowing the soft drink to go flat.
A buret 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, and other acid-base
neutralization reactions.
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.
In order to analyze unknown acids/bases, we must have a "standard" solution to react with the
unknowns. A standard solution is one in which the concentration is known accurately. We will first
prepare a standard solution of NaOH. One way to prepare a standard solution is to dissolve an
accurately massed amount of the substance and dilute it to a measured volume. In this way, the
concentration can be calculated exactly. However, it is usually impossible to obtain NaOH of sufficient
purity to use it as a primary standard. An indirect method is more practical for obtaining a standard
solution of NaOH. We will prepare a solution of an approximate molarity and standardize it against a
primary standard of known purity.
Many compounds can be used as primary standards. However, a primary standard should:
a. be of high purity
b. remain unchanged in air during massing and remain stable during storage
c. have a high molar mass to reduce massing errors
d. react with the solution to be standardized in a direct, well-defined reaction
Potassium hydrogen phthalate (KHP) will serve as our primary standard. This is a large molecule
(KHC8H4O4) with a molar mass of 204.2 g/mol. KHP is a monoprotic acidic substance and will react with
NaOH in a simple 1 to 1 relationship according to the following equation:
NaOH(aq) + KHC8H4O4(aq)  KNaC8H4O4(aq) + H2O(l)
After we standardize the sodium hydroxide with KHP, we will use this standard base to determine the
concentration of the citric acid in various brands of soda.
PROCEDURE
A. STANDARDIZATION OF NaOH
1. Clean your buret by rinsing it with deionized water.
2. Rinse your buret twice with 10 mL of the NaOH solution that you will be standardizing.
3. Fill the buret with the NaOH solution.
4. Accurately weigh approximately 0.5 g KHP and transfer it to one of your 250 mL Erlenmeyer
flasks.
5. Add approximately 25 mL of deionized water and 2 to 3 drops of the phenolphthalein
indicator. Swirl the flask until all the KHP has dissolved.
6. Read the initial volume of NaOH solution in your buret to the nearest tenth of a mL.
7. Slowly add the NaOH solution to the dissolved KHP until the first faint lasting pink color
appears. Placing a white piece of paper under your Erlenmeyer flask helps you to better see
this color change.
8. Once a lasting pink color (at least 30 seconds) has been reached, read your final volume
from your buret to the nearest tenth of a mL.
9. Calculate the molarity of the NaOH solution.
10. Repeat steps 1 through 9.
11. Calculate the average molarity of the NaOH solution from the two trials.
B. TITRATION OF A SOFT DRINK
NOTE - The citric acid content in many soft drinks is quite low, so good technique is critical.
1. Obtain 75 mL of one soft drink sample as assigned by the instructor. The soft drink samples
in your lab have been opened several days in advance to assure they are decarbonated.
2. Pipette 25 mL of the soft drink into a clean 250 mL Erlenmeyer and add 2 to 3 drops of
phenolphthalein.
3. Fill your buret with the standardized NaOH. Read the initial volume of NaOH solution in
your buret to the nearest tenth of a mL.
4. Slowly add the NaOH solution to the soft drink until the first faint lasting (at least 30
seconds) pink color appears.
5. Once a lasting pink color has been reached, read your final volume from your buret to the
nearest tenth of a mL.
6. Calculate the molarity of the citric acid in the soft drink sample.
7. Repeat this procedure (steps 1 through 6) using a second 25 mL sample of the same soft
drink.
8. Calculate the average molarity of citric acid from the two trials.
9. From your average molarity of citric acid, calculate the number of moles of citric acid per
can or bottle of soft drink.
10. Finally, calculate the number of grams of citric acid there are in one can or bottle of soft
drink.
11. Repeat steps 1 through 10 using a second brand of soft drink as assigned by the course
instructor.
CHEM 2124
General Chemistry II Laboratory
Week 3 – DETERMINING THE CONCENTRATION OF CITRIC ACID IN
SOFT DRINKS VIA ACID/BASE TITRATION (MAJOR LAB #1)
Name
Date
PRE-LABORATORY ASSIGNMENT – Must be completed prior to class.
SHOW ALL WORK ON CALCULATION PROBLEMS!
1. Write out the balanced neutralization reaction between citric acid and sodium hydroxide.
2. A solution is created by dissolving 25.0 g of NaCl in 150 mL of water. What is the molarity of
this solution?
3. What is the molarity of a diluted solution if 30.0 mL of 1.25M acetic acid is diluted to a final
volume of 175 mL?
4. Given the following reaction: 2Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g)
How many grams of hydrogen are produced when 75.0 mL of 0.825 M hydrochloric acid is
reacted with an excess of aluminum?
5. How do you know when a titration of citric acid with a base has reached the endpoint?
Data and Calculations
A. Standardization of NaOH with Potassium Hydrogen Phthalate (KHP)
Trial 1
A
Mass of KHP (g)
B
Initial Buret Reading (mL NaOH)
C
Final Buret Reading (mL NaOH)
D
Total Volume of NaOH used (mL) (C-B)
E
Molarity of NaOH (mol/L) (see below)
Trial 2
Calculate the molarity of NaOH as follows:
M NaOH = g KHP
mL NaOH
x 1mol KHP x
204.23 g KHP
Calculations of NaOH Molarity:
Trial #1:
Trial #2:
Average Value of Molarity of NaOH (mol/L):
1 mol NaOH
1 mol KHP
x
1000 mL
1L
B. Titration of Soft Drink with NaOH
Brand Name of Soda Used:
Trial 1
A
Volume of soda used (mL)
B
Initial buret reading (mL NaOH)
C
Final buret reading (mL NaOH)
D
Total Volume of NaOH used (mL) (C-B)
E
Molarity of citric acid in sample (mol/L)
(see below)
Calculate the molarity of the citric acid as follows:
M Acid = mol NaOH x mL NaOH x 1 mol Acid x
1L
3 mol NaOH
1
mL Soda
Show the calculations of NaOH Molarity:
Trial #1:
Trial #2:
Average Value of Molarity of Citric Acid in sample (mol/L):
Volume of container of soda (mL and L):
Moles of Citric Acid per bottle of soda:
Calculations:
Molar mass of Citric Acid (g/Mole):
Mass of Citric Acid per bottle of soda (g):
Calculations:
Trial 2
Brand Name of Soda Used:
Trial 1
A
Volume of soda used (mL)
B
Initial buret reading (mL NaOH)
C
Final buret reading (mL NaOH)
D
Total Volume of NaOH used (mL) (C-B)
E
Molarity of citric acid in sample (mol/L)
(see below)
Calculate the molarity of the citric acid as follows:
M Acid = mol NaOH x mL NaOH x 1 mol Acid x
1L
3 mol NaOH
1
mL Soda
Show the calculations of NaOH Molarity:
Trial #1:
Trial #2:
Average Value of Molarity of Citric Acid in sample (mol/L):
Volume of container of soda (mL and L):
Moles of Citric Acid per bottle of soda:
Calculations:
Molar mass of Citric Acid (g/Mole):
Mass of Citric Acid per bottle of soda (g):
Calculations:
Trial 2