Chemistry 151
Lab 9: Standardizing a Solution (Titrations, Pt I)
Last Updated Dec. 2013
Introduction
Sometimes, it’s necessary to verify the concentration of a solution or determine the concentration with a
higher degree of certainty (i.e., more significant figures) than what you originally measured while preparing
them. This is known as standardizing the solution. The most common method of standardization is a
titration, where a solution of one reactant (the titrant) is slowly added to the other. In an acid-base
titration, the titrant can be either reactant (whichever is more convenient or makes the most sense).
A titration is typically performed with a buret (or burette), which is essentially a long, narrow graduated
cylinder with a value attached to the bottom. This allows you to add the titrant in dropwise amounts.
Furthermore, burets are usually graduated in 0.1mL marks, so you can usually determine the necessary
amount of titrant with a fairly high degree of certainty (i.e., more significant figures).
One common issue that’s encountered with acid-base titrations is that all chemical species in the reaction,
both reactants and products, are often colorless in aqueous solution. This can make it difficult to determine
when the reaction has reached completion. To resolve this issue, an indicator solution is used, which
changes color at or near the point where the titrated solution has been consumed by the titrant. This color
change is often called the endpoint of the titration. Usually, only 2-3 drops of an indicator solution is
required for a typical titration, so the indicator itself should have little or no adverse effect on the reaction
being studied.
It this lab, you will prepare and standardize a 1 M solution of sodium hydroxide and determine its
concentration to at least three significant figures. To do this, you’ll use it to titrate a solution of potassium
hydrogen phthalate (KHC8H4O4, often abbreviated KHP). Hydrogen phthalate, HC8H4O4-, is an organic
anion that contains an acidic hydrogen. It will react with NaOH via a neutralization reaction, the same way
it would with other acids you’ve encounter in the class.
HC8H4O4-(aq) + OH-(aq)  C8H4O42-(aq) + H2O(l)
or
KHP(aq) + NaOH(aq)  NaKP(aq) + H2O(l)
One advantage to using KHP for our standardization is that it’s a solid a room temperature, allowing us to
measure our acid with a higher degree of accuracy than you could with a solution.
The indicator used for this reaction is phenolphthalein, which is colorless in acidic solution, but will begin
turning pink to magenta in basic solution. The phenolphthalein will be added to your KHP solution and the
NaOH will be your titrant. Once the KHP has been completely neutralized, any excess base you add will
remain in solution, causing the solution to turn pink in color.
Procedure
Part I: Preparing your NaOH solution
Prepare a 400 mL solution of 1 M NaOH using the amount you calculated in the Pre-lab assignment. If
neither you nor your lab partner thought to record this amount, you’ll need to decide which one of you will
have to make that “walk of shame” to the instructor’s table—or a neighboring group—to find out how
much you’ll need (and I hope you’ve finally learned the importance of reading the entire lab ahead of time;
I mean, seriously, you’ve been doing this for two months). Transfer your solution to a 500 mL bottle,
labeled with your and your partner’s name, as well as your lab section.
Part II: Preparing the Buret
1. Attach a buret clamp to a ring stand.
2. Add 20 mL of your sodium hydroxide solution to a buret and place the buret in the clamp. Inspect the
stopper for leaks. If you see any, inform your instructor before continuing.
3. Slowly pour the sodium hydroxide into the sink, rotating the buret as you do in order to clean as much
of the inside as possible. Repeat this step with an additional 20mL of sodium hydroxide.
4. Clamp the buret and fill it to the 0.0 mL mark with sodium hydroxide. If air bubbles form, try to
remove as many as you can by gently tapping the side of the buret.
Part III: Standardizing Your NaOH Solution
1. Using an analytical balance, measure approximately 5 g of KHP and transfer it to an Erlenmeyer flask.
2. Add 80 mL of water to the flask and mix to dissolve the KHP. If it doesn’t all dissolve, that’s OK.
3. Add 3 drops of phenolphthalein to the solution.
4. Begin titrating sodium hydroxide to the solution. As soon as you see a pink color appear in the
solution, turn off the valve.
5. Swirl the flask to mix the solution then continue adding sodium hydroxide dropwise to flask. Continue
swirling the flask as you add titrant until the pink color persists for at least 30 seconds. Any
undissolved KHP from step 2 should be dissolved by this point in the titration.
6. Repeat steps 1-5 four times to give you a total of five determinations. Do not refill the buret unless it
appears that you don’t have enough to perform the next determination.
Your remaining NaOH solution will be set aside (your instructor will show you where) and used for next
week’s lab. Make sure to add the standardized concentration to the label.
Waste Disposal
All waste can be poured down the sink with running water. Rinse you buret with water to remove any
remaining traces of NaOH, which can dry up and clog the buret tip.
Name: _____________________________
Section: ________
Data
Trial
1
2
3
4
5
Mass of KHP, g
_______
_______
_______
_______
_______
Initial buret reading, mL
_______
_______
_______
_______
_______
Final buret reading, mL
_______
_______
_______
_______
_______
Volume NaOH titrated, mL
_______
_______
_______
_______
_______
mol KHP reacted
_______
_______
_______
_______
_______
mol of NaOH titrated
_______
_______
_______
_______
_______
Concentration of NaOH, M
_______
_______
_______
_______
_______
Average Concentration
_______
Show your work for each of the following calculations from Trial #1.
1) Volume NaOH titrated
2) Moles of KHP reacted
3) Moles of NaOH titrated
4) Concentration of NaOH
Name: _____________________________
Section: ________
Post-lab Questions
1. How would each of the following affect your calculated molarity (too high, too low, or no effect)? In
each case explain your answer.
a) When reading the buret, the top of the meniscus was read instead of the bottom.
b) Several air bubbles formed near the bottom of the cylinder and in the tip of the dropper. These bubbles
left the buret with the solution as the sample was titrated.
c) The jar of KHP was left open overnight and the solid (which is hygroscopic) had become hydrated by
absorbing moisture in the air.
d) The KHP was dissolved with 100 mL of water instead of 80 mL.
2. After dissolving the KHP with water, a student noticed some undissolved solid, but continued with the
titration as instructed by the procedure. He titrated until the pink color persisted for 30 sec and recorded
the volume of titrant required. Preparing for the second determination, he set the reaction solution aside
and added some more NaOH to the buret. He was about to pour the reaction solution down the drain, so he
could rinse and reuse the flask, when he saw that the pink color had disappeared!
a) What happened?
b) What can the student do to correct for this error and salvage this trial? How should the data sheet be
modified to reflect this correction?
3. Another student began adding titrant to the acid solution, but stopped when she suddenly realized that
she forgot to add the 3 drops of phenolphthalein. When she did, the solution immediately turned dark red!
a) What happened?
b) What can the student do to correct for this error and salvage this trial? How should the data sheet be
modified to reflect this correction?
Name: _____________________________
Section: ________
Pre-lab Questions
1. Calculate the molar mass of KHP.
2. Calculate the mass of NaOH required to prepare 400.0 mL, 1.0 M solution.
3. Estimate the amount of 1.0 M NaOH (in mL) required to neutralize 5.0 g of KHP.
4. Given your previous experience with sodium hydroxide, why is it a good idea to standardize NaOH
solutions, even if they were prepared using solid NaOH?
5. Why is using an acidic solution for this titration less accurate than using an acid such as KHP, which is a
solid at room temperature.
6. How will you know when all of the acid has been neutralized by the sodium hydroxide?
7. A 3.0 g sample of sodium hydroxide was dissolved in water and titrated with a sulfuric acid solution.
a) Write a balanced equation (with phases) for the reaction that’s taking place.
b) If 22.8 mL of sulfuric acid was required to neutralize the sodium hydroxide, what was the concentration
of the acid?