Standardization of a Sodium Hydroxide Solution
Frederick R. Wight
Manatee Community College
Objective: To determine the molarity of a sodium hydroxide solution
Principles:
Acid base titrations are among the most common laboratory procedures. In any
titration a solution of unknown composition is analyzed quantitatively by adding a second
solution of known concentration (primary standard solution) until the reaction is
complete. The reaction is complete when the mole ratio of the two reacting substances
are the same as what appears in the balanced equation. This is the stoichiometric point or
equivalence point in the titration. In this experiment the equivalence point for an acidbase titration is detected using phenolphthalein indicator, colorless in acidic and red in
basic solutions. The point at which the phenolphthalein indicator changes color is the
endpoint of the titration. Indicators are selected so that the equivalence point and the
endpoint occur at essentially the same point in the titration.
In this experiment we will determine the molarity of a sodium hydroxide solution
using the primary standard potassium acid phthalate (KHP):
O
O
OH
-
O K
O
+
+ NaOH
-
O K
O
+
Na
+
O
KHP
At the equivalence point equal moles of KHP and NaOH have been combined.
Since the number of moles of NaOH and KHP are equal all we need is the number of
moles of KHP that was neutralized and the volume of NaOH used to get to the
equivalence point.
Molarity of NaOH, M =
moles KHP
Liters of NaOH
Once the sodium hydroxide solution is standardized it can be used to determine the
molarity of any acid solution. Consider the following example. A 50.00 mL solution of
hydrochloric acid is titrated with a 0.100M solution of sodium hydroxide. The
phenolphthalein end point was found at 37.50 mL of NaOH. What is the molarity of the
sodium hydroxide? Solving such problems (of which there will be many in this course)
requires a three step strategy:
1) Figure out how many moles of the titrant (in this case, the base) were needed.
2) Use the balanced chemical equation to calculate the moles of analyte (in this case, the
acid) present.
3) Use the volume of analyte to find the concentration of the analyte.
Use M x V (moles/liter x liters) to get moles of titrant:
0.100 x .0375 = 0.00375 moles NaOH
HCl + NaOH  NaCl + H2O
1 mole base per mole of acid: 0.00375 moles HCl
Molarity HCl = moles/Liter
0.00375/.050 = 0.075M HCl
Procedure:
You are to complete at least three trials in this experiment. The three determinations
should be within 3%, if they are not you must do another determination.
1.
2.
3.
4.
5.
6.
7.
8.
9.
Accurately weigh out 0.7-0.9 g of KHP (molar mass = 204.2 g/mol) in a labeled
125 mL Erlenmeyer flask. Add 50 ml of DI water and 2 drop of phenolphthalein
indicator. Place a stir bar in the bottom of the flask and stir gently with a magnetic
stirrer until dissolved. (Rapid stirring introduces carbon dioxide into the water
making the water slightly acidic).
Get approximately 80 mL of the unknown NaOH in a labeled 150-200 mL beaker.
Attain a 50 mL buret and thoroughly wash, rinse with two portions of DI H2Oand
then with 5 mL of NaOH. Make sure to rinse the sides well.
Fill the buret with the NaOH solution and record it’s volume to 0.01 mL.
Place a piece of white paper under the KHP flask which should now be under the
NaOH filled buret.
Slowly add the NaOH solution to the KHP, swirling the flask after each addition.
As the rate of the indicator color change decreases, decrease the rate of NaOH
addition; proceed with drop addition of NaOH until the indicator endpoint is
reached. This is when 1 drop causes the pink color to last for 30 seconds.
Read the volume of the NaOH to 0.01 mL and record.
Repeat the titration with two additional KHP solutions.
Data Sheet
1
Initial Buret
Reading, mL
Final Buret
Reading, mL
Volume NaOH
Dispensed
Grams of KHP
Moles of KHP
Moles of Base
Molarity of Base
Average Molarity of
Base
2
3
Pre-Lab:
1.
What is a primary standard?
2.
What is the purpose of placing a white piece of paper under the flask during this
titration?
3.
A mass of 0.497 g sulfamic acid, NH2SO3H, dissolved in 50.0 mL of water is
neutralized by 28.4 mL of NaOH at the phenolphthalein endpoint. What is the
molarity of the NaOH solution? The formula weight of NH2SO3H is 97.1 g/mol:
NH2SO3H + NaOH  NH2SO3-Na+ + H2O
4.
If the endpoint in the titration is surpassed (too pink) what effect does this have on
the calculated molarity of the NaOH solution? Explain.
5.
Why does the phenolphthalein color change fade with continual stirring?
Post Lab Questions:
1) A 25.00 mL sample of HBr is titrated with a 0.150 M standardized sodium hydroxide
solution. The endpoint was reached when 18.80 mL of titrant had been added.
Calculate the molar concentration of the HBr.
2) A 20.00 mL sample of sulfuric acid (H2SO4) is titrated with a 0.100 M solution of
sodium hydroxide. The endpoint was reached when 45.65 mL of titrant was added.
Calculate the molar concentration of sulfuric acid.
3) A 1.00 gram sample of an unknown acid HA is dissolved in 50.0 mL of water and
titrated with 0.150 M sodium hydroxide. The endpoint was observed after 24.50 mL
of titrant had been added. Calculate the molecular weight of the acid HA.