Name_______________________________________
Date_______________
Acid Base Titration
Background
Titration is a method by which an unknown concentration of a solution can be determined by comparing
it with a solution having a predetermined or standard concentration. To titrate an acid versus a base,
one can add a basic solution of unknown concentration to an acid until neutralization occurs, or when
the number of moles of hydronium (H3O+) ions equals the number of moles of hydroxide (OH-) ions. A
substance that changes color when ion balance is achieved called an indicator is used to determine
when this balance in ions occurs and is known as the equivalence or end point of the titration. By
knowing the volume of both the acid and base used in the titration as well as the molarity of the acid,
one can find the unknown concentration of the base through stoichiometry of the acid-base
neutralization reaction. Phenolphthalein solution is an indicator that works well with the acid and base
chosen for this investigation. Phenolphthalein is colorless in acid solution and pink in basic solution.
Burettes are used to dispense the acid and base (see picture below).
http://suhow.info/sh-procedure-for-complexometric-titration
Objectives
1. After completing this laboratory investigation, the student will be able to describe the acid base
titration process including the use of burettes and titration endpoints.
2. After completing this laboratory investigation, the student will be able determine molarities of
acids or bases from titration data.
Materials/Equipment
(per team)
Beakers, 250 mL (2)
Burettes, 50 mL (2)
Distilled water
Double burette clamp
Erlenmeyer flask, 125 mL (1)
Funnel (1)
Hydrochloric acid (0.15 M HCl, 100 mL)
Phenolphthalein solution (2 mL)
Name_______________________________________
Date_______________
Ring stand
Sodium hydroxide solution (unknown concentration NaOH)
Stirring rod
Procedure
1. Use a ring stand with the double burette clamp and two burettes to set up the titration
apparatus. Label one burette acid and the other base. Place a 250 mL beaker under each
burette tip.
2. Add 5-10 mL of 0.15 M HCl to the acid burette to rinse the burette and drain into the 250 mL
beaker. Complete the same process with the base burette using the unknown concentration of
NaOH. Discard the acid and base.
3. Fill the acid burette with the 0.15 M HCl carefully until the acid is above the zero mark.
Dispense acid from the burette until all air is removed from the burette tip and the level of acid
is within the graduated portion of the burette. Record the burette level as the initial burette
reading of HCl for trial 1. Use the bottom of the meniscus for your burette level reading.
Complete the same process with the base burette using the unknown concentration of NaOH.
4. Into a clean 125 mL Erlenmeyer flask, dispense approximately 10 mL of the acid. Add
approximately 50 mL of deionized water and three drops of phenolphthalein solution to the
flask and swirl to mix.
5. Titrate slowly by adding the base into the flask containing the acid with stirring until a pink color
starts to persist in the beaker. Decrease the flow of base to a slow drop by drop process and
continue until the pink color no longer disappears.
6. If you overshoot the titration endpoint, slowly add acid drop by drop until the pink color
disappears. Then add base again drop by drop until a faint pink color persists. Repeat this
titration process until one drop of base will cause the faint pink color to remain.
7. Record the final acid and base burette readings in the data table.
8. Repeat steps 4-7 to complete two more trials. Since you are using the same acid and base, you
will be able to complete step 5 faster because you will have a reasonable estimate of the
amount of base necessary to complete the titration.
Name_______________________________________
Date_______________
Data Table
Molarity of HCl solution _______________________
Trial 1
Trial 2
Trial 3
Trial 1
Trial 2
Trial 3
Initial burette reading
for HCl (mL)
Final burette reading
for HCl (mL)
Initial burette reading
for NaOH (mL)
Final burette reading
for NaOH (mL)
Calculations Table
Volume of HCl
dispensed (mL)
Volume of NaOH
required (mL)
Moles of HCl
dispensed
Moles of NaOH
required
Molarity of NaOH
solution
Data Analysis
1. Write a balanced chemical equation for the reaction between HCl and NaOH.
2. Calculate the volume of HCl dispensed and NaOH required from each trial and place in the
calculations table.
3. Using the molarity formula, calculate the moles of HCl dispensed in each trial and place in the
calculations table.
Name_______________________________________
Date_______________
4. Using the mole ratio from your balanced chemical equation, calculate the moles of NaOH
required from each trial to neutralize the HCl and place in the calculations table.
5. Calculate the average molarity of NaOH from your three trials.
Conclusions
1. Compare your average molarity of NaOH to the one supplied by your teacher as the accepted
value. Determine a percentage error for your average molarity of NaOH.
2. Investigate the transition interval for phenolphthalein and determine what pH interval its color
change corresponds to.
3. Describe two potential errors in the titration process.