Laboratory Investigation:
Titration of an Unknown Acid
Purpose:
To prepare a solution of NaOH, and to standardize that solution.
To use titration to determine the concentration of a solution of hydrochloric acid of unknown
concentration.
Introduction:
In acid-base titration, an acid and base are reacted in order to determine the concentration of one
or the other. In this lab we will react a solution of NaOH that we prepare with solid oxalic acid,
H2C2O4, and with acid of an unknown concentration. The titration with the solid oxalic acid will
be used to determine our NaOH solution concentration. Once we have standardized the NaOH
solution, we will then determine the unknown acid concentration.
Materials:
Solid NaOH
Volumetric Flask
Indicator Solution
Oxalic Acid Dihydrate(H2C2O4. 2H2O)
Erlenmeyer flask
Hydrochloric Acid of unknown concentration
Buret
Wash Bottle
Procedure:
A. Preparation of a Solution of NaOH
1. Obtain a 100 mL volumetric flask and fill it half way with deionized water.
2. Calculate the mass of NaOH needed to make 100.0 mL of a 0.250 M solution of NaOH.
Check your answer with your instructor before proceeding.
3. Mass out as close to this value as you can using a piece of weighing paper for the solid.
Record the exact mass of NaOH used on your data sheet. You will use your this value as your
“true value” to determine the percent error of your NaOH solution.
4. Immediately add your solid NaOH to the water in the volumetric flask. Put the top on the flask
and mix as directed until the base has dissolved.
5. Add deionized water until the meniscus of the solution reaches the line on the flask.
6. Put the top on the flask and mix the solution again.
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B. Standardization of the NaOH
1. Mass out as close to 0.20 g of oxalic acid dihydrate (H2C2O4·2H2O) as you can using a piece
of weighing paper. Record the exact value used. Do not exceed .25 g of solid acid!
2. Add the oxalic acid dihydrate to an Erlenmeyer flask. Add approximately 25 mL of deionized
water to the flask, and dissolve the acid by swirling gently. If not all of it dissolves, it will
dissolve as the titration proceeds.
3. Rinse the buret on the “B” side with deionized water. Add the NaOH solution (from Part A)
to the buret. Fill the buret to as close to the 0.00 mark as you can. Open the stopcock and let the
tip fill with NaOH solution. Read the buret to the closest 0.01 (remember to read the buret
directly from the top down) and record.
4. Add 3 DROPS (too much indicator will cause cloudy solutions!) of the indicator to the flask,
and titrate the solid acid with your NaOH solution. Try to obtain the lightest permanent shade of
pink in the flask that you can. Remember to record starting and final volumes for the base.
5. Repeat the titration using a second sample of oxalic acid dihydrate. Again, record the mass of
the solid, and initial and final volumes of the base. If these two trials are off by more than 10%,
repeat the titration, and then average the two closest trials.
C. Determination of the Molarity of an Unknown
1. Obtain a solution of HCl with an unknown molarity from your instructor.
2. Rinse the buret on the “A” side with deionized water. Add the HCl solution to this buret until
it is about half full. Again, open the stopcock to fill the tip with acid solution. Read the buret to
the closest 0.01. Read and record the level on your base buret (refill if needed).
3. Put between 10 and 11 mL of HCl into the Erlenmeyer.
4. Add 3 drops of the indicator to the flask, and titrate HCl with your NaOH. Try to obtain the
lightest permanent shade of pink in the flask that you can.
5. If your solution turns dark pink, add acid drop-wise until it turns clear again. Then, titrate
with base until you achieve a light pink color.
6. Remember to record starting and final volumes for the acid and base. Complete a second trial
of part C.
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Calculations:
From Part A and B Data:
1. From the mass of the solid oxalic acid dihydrate, calculate the number of moles of acid used
in each titration. Hint: since this is a dihydrate, remember to include the mass of water!
2. Complete and balance the equation for the neutralization (below) and calculate the number of
moles of base used in each titration.
___H2C2O4 + ___NaOH  ______________________
3. From your data, calculate the molarity of the NaOH for each trial.
4. Average the two trials to find the average experimental Molarity for your NaOH.
5. Calculate the Molarity of the NaOH using the mass of NaOH you measured. Calculate
percent error using this as your true value.
From Part C Data:
1. For each trial, calculate the volumes of acid and base used.
2. Use these volumes and the molarity of your base obtained in the calculations above to
calculate the Molarity of the unknown HCl.
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Sample DATA TABLE:
Part A.
Mass of NaOH used in solution
Part B.
Trial #1
Trial #2
Trial #3
Mass of Oxalic Acid
Initial Buret reading
Final Buret reading
Volume of NaOH required
Part C.
Trial #1
Trial #2
Initial Buret reading-acid
Final Buret reading-acid
Initial Buret reading-base
Final Buret reading-base
Lab Notebook Checklist:
1. Headings
2. Introduction:
a. Background
b. Equations (Molarity & percent error)
c. Safety (HCl, H2C2O4, NaOH, phenolphthalein)
3. References: handout, online safety sources
4. Purpose: summarize from handout
5. Materials: list from handout
6. Procedure: reference handout
7. Data Table: see sample above
8. Calculations: numbered and titled showing all work
9. Discussion: chemistry learned, techniques used and comment on possible experimental
errors
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