Determination of a Sodium Hydroxide Solution Concentration by Titration
Purpose:
To determine the concentration of a sodium hydroxide solution using a standardized solution.
Background:
Sodium hydroxide is stored in solid pellets. To create a solution of sodium
hydroxide, a mass is weighed and dissolved in distilled water. However, solid
sodium hydroxide is incredibly hygroscopic. This means that it readily absorbs
water vapour from the air.
This means that as soon as the sodium hydroxide is exposed to air, it absorbs
water and the mass weighed on a balance is a mix of both water and sodium
hydroxide. This changes the concentration of the solution made from what was
theoretically determined.
This means that if we must use other chemistry techniques to
determine the exact concentration of a sodium hydroxide solution.
This is often done with a titration. A titration is when a known
amount of standard solution is reacted with a solution of unknown
concentration. A standard solution is one with a precisely known
concentration and a carefully measured mass.
During a titration, the standardized and unknown solution are
mixed. When the reaction is completed, the colour of the reaction
changes using an indicator.
The value obtained from this then double checked with a second
standard to ensure that the concentration obtained was correct.
This experiment will use this method to determine the concentration of a sodium hydroxide solution.
This experiment will take 4 main stages:
1. Making a Primary Standard – a standard solution potassium hydrogen phthalate is prepared
2. Standardizing a Solution - sodium hydroxide solution is standardized with a titration using the
primary standard
3. Performing a Serial Dilution – a secondary standard of sulphuric acid is prepared using a serial
dilution
4. Titrate to Verify the Concentration of a Prepared Solution – the concentration of the sodium
hydroxide solution is double checked using the secondary standard
Part I: Making a Primary Standard
A primary standard is a solution containing a carefully measured mass and an accurately known
concentration.
1. Find the molar mass of potassium hydrogen phthalate (KHC8H4O4).
2. At one of the weigh stations, weigh exactly 2.00 g of KHC8H4O4 into a weigh boat, and carefully put
this into a small beaker.
3. Add about 50 mL of distilled water to the KHC8H4O4. Use the glass stir rod to stir until it is all
dissolved, and then transfer the solution to the 100 mL volumetric flask. Wash the beaker carefully with
additional distilled water and pour this into the volumetric flask also. Add distilled water to the
volumetric flask so that the meniscus sits on the line in the neck. Stopper the flask and mix thoroughly
by inversion (no milkshake mixing!) **If you add too much water you must RESTART**
4. Calculate the molarity of the solution (keep 3 sig figs):
Mass of KHC8H4O4
Molar mass of KHC8H4O4
2.00 g
Moles of KHC8H4O4
Volume of flask
Molarity of KHC8H4O4
*THIS IS “CALCULATION 1”
Part II: Standardizing a Solution
100.00 mL
Sodium hydroxide (NaOH) absorbs moisture out of the air, so we can not be sure of the mass that was
used (i.e. some of what we weighed out was actually water.) To find the precise molarity of a solution of
NaOH we will do a titration to measure it against our primary standard, made in part I. This is called
standardizing a solution.
1. Fetch from the front bench a 150 mL beaker that has a sample of NaOH solution in it.
2. Set up your titration apparatus.
3. Rinse the burette with the NaOH solution and then fill with some NaOH. 4. Record the starting place
on the burette.
5. Using a volumetric pipette, measure EXACTLY 10 mL of your primary standard into an Erlenmyer flask
(calculate NOW how many moles of KHC8H4O4 this represents!) *You MUST understand, include &
show this calculation in your lab write up! *CALCULATION 2
Moles KHC8H4O4:
Add 4 drops of phenolphthalein indicator. Phenolphthalein is clear in an acid solution, but pink in a
basic solution.
6. Titrate with NaOH until the pink colour stays for 20 seconds. Make sure you are swirling the
Erlynmyer flask gently as you titrate into it. **Slowly titrating will provide more accurate results**
7. Record the final reading for the NaOH.
- The final pink-colored titrated solution is neutralized (done reacting) and can be safely dumped down the drain
8. Repeat the titration twice more.
Trial 1
Trial 2
Trial 3
NaOH start
NaOH end
NaOH used
9. Balance the equation:
_____KHC8H4O4 + _____NaOH  _____Na KC8H4O4 + _____H2O
10. Using the moles of KHC8H4O4 that you determined in step #4, determine how many moles of NaOH
you just neutralized. *You MUST understand, include & show this calculation in your lab write up!
*CALCULATION 3
11. Use the average of the three volumes you titrated, together with the moles of NaOH you just
determined, and calculate the molarity of the NaOH. *You MUST understand, include & show this
calculation in your lab write up! *CALCULATION 4
Part III: Perform a Serial Dilution
Using only a 10 mL graduated pipette and the 100 mL volumetric flask, make a 0.03 M solution of H2SO4 from the
3.00 M H2SO4 stock solution (found in beakers on the front bench). You may not use a measurement less than 5.0
mL. *Show a sample calculation – Calculation 5
Dilution 1
Dilution 2
M1
3.00 M
V1
M2
V2
0.03 M
Be careful with the sulfuric acid. It wrecks your jeans. And you skin. And your eyes. You get the idea.
Part IV: Titrate to Verify the Concentration of a Prepared Solution
We will use the concentration of the sulfuric acid you just prepared in Part III to check the standardized sodium
hydroxide solution’s concentration.
1. Fetch from the front bench a 150 mL beaker that has a sample of NaOH solution in it. You already know the
molarity of this solution: you determined it in Part II.
2. Fill the burette with the NaOH. Record the starting place on the burette.
3. Using a pipette, measure exactly 10.0 mL of your H2SO4 into an Erlynmyer flask. Add four drops of bromothymol
blue. Bromothymol blue is clear in an acid, but blue in a base.
4. Titrate with NaOH until the blue colour stays for 20 seconds.
5. Record the final reading for the NaOH.
- The final blue-colored titrated solution is neutralized and can be safely dumped down the drain
6. Repeat the titration twice more.
Trial 1
Trial 2
Trial 3
NaOH start
NaOH end
NaOH used
7. Balance the equation:
_____NaOH + _____H2SO4  _____Na2SO4 + _____H2O
8. From the molarity of the NaOH and the average of the three volumes you titrated, calculate how many moles of
NaOH it took to neutralize the H2SO4. *You MUST understand, include & show this calculation in your lab write
up! Calculation 6
9. Use the molar ratio between NaOH and H2SO4 to determine how many moles of H2SO4 you neutralized. *You
MUST understand, include & show this calculation in your lab write up! Calculation 7
10. Using the number you just calculated together with the volume of H 2SO4 you used (10.0 mL) calculate the
molarity of the H2SO4. Compare it to what it was supposed to be, i.e. 0.03 M. *You MUST understand, include &
show these calculations in your lab write up! Calculation 8
*Use a percent error calculation to compare your results to the 0.03. Calculation 9
IMPORTANT** Disposal of excess or left over Chemicals
-
Acids (H2SO4) on their own are dangerous & should not be poured down the drain
Bases (NaOH) on their own are dangerous & should not be poured down the drain
When acids and bases are mixed together they NEUTRALIZE each other and the resulting solution is chemically
safe
To safely neutralize and dispose of your excess chemicals, do the following:
o Find a large beaker
o Pour your left over H2SO4 from your volumetric flasks into the beaker
o Pour your left over NaOH from your burette into the beaker
o The solution is now neutralized and can safely be poured down the drain with excess water