Titration with Oxalic Acid
Abstract
Titration is a versatile analytical procedure that can be used for a wide variety of chemical
analyses. For example, when your town’s water supply is tested for purity, or pond water is
tested for dissolved oxygen and contaminants, chances are a titration is carried out. Some
tests essential for a medical diagnosis require a titration of various body fluids.
A titration makes use of a known reaction between two chemicals. A solution of unknown
concentration is reacted with a precisely measured amount of another chemical. An
appropriate indicator must be used to determine when chemically equivalent amounts of
each chemical are present, that is, when no excess of either reactant is present. This is
known as the equivalence point. To measure solution volumes accurately, finely calibrated
pipets and burets are used. Titrations are commonly used to determine the strength of acids
or bases.
Acid-base titrations follow a relatively standard procedure for analysis of acid or base
strength. The concentration of either an acid or base solution can be determined. A measured
amount of acid is neutralized by reacting it with a basic solution titrated from a buret.
Consider the following example:
HCl + NaOH → NaCl + H2O
(known) (unknown)
The example shows that the reaction is a special case of a double replacement reaction,
where the products are water and a corresponding salt. At the equivalence point, all of the
NaOH has been reacted with the HCl and a pH of 7 (a neutral solution) would result. Some
reactions (ones with weak acids or weak bases) do not reach their equivalence point at a pH
of 7 because of equilibria occurring in the solution.
In this experiment, you will determine the concentration of a solution of NaOH by reacting it
with a known concentration of oxalic acid, H2C2O4 – a weak acid, which will have an
equivalence point at approximately pH 8. The indicator phenolphthalein will be used to find
the endpoint of the titration because it is colourless below pH 8, but turns pink just above pH
8.
Purpose
To determine the concentration of a sodium hydroxide solution using the process of titration.
Titration with Oxalic Acid
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Materials
safety goggles
laboratory apron
2 Erlenmeyer flasks, 125mL
balance
oxalic acid dihydrate (H2C2O4 · 2H2O)
distilled water
phenolphthalein
wash bottle
buret
buret clamp
ring stand
sodium hydroxide solution (NaOH)
2 beakers, 100mL
white paper
Procedure
A: Preparation
1. Put on your safety goggles (and lab apron if desired). While one partner prepares the
oxalic acid solution (steps 2 + 3), the other should prepare the buret for titration (steps 4 –
9).
2. In a 125 mL Erlenmeyer flask, weigh out approximately 0.20g of oxalic acid (this will be
a very small amount). Record the exact mass of oxalic acid used in the data table. Repeat
this step in another Erlenmeyer flask for trial 2.
3. Dissolve the oxalic acid in the flask with approximately 50 mL of distilled water. Add 23 drops of phenolphthalein indicator to the flask.
4. Obtain a buret and buret funnel.
5. Rinse the buret thoroughly with tap water, then rinse it once with a small amount of
distilled water, draining the final rinses through the tip. Clamp the clean buret to the ring
stand.
6. Obtain 50-60 mL of NaOH solution in a 100 mL beaker. Position the buret so that the top
is below eye level and ensure it is closed!
7. Pour approximately 5 mL of the base into the buret. Drain this solution through the tip to
remove water and coat the inside of the buret with base.
8. Fill the buret to slightly above the zero line. With a waste beaker below the buret, drain
some of the base through the tip to clear the buret of air. Stop between 0.0 and 2.0 mL.
Remove the hanging drop by touching the tip to the inside of the waste beaker.
9. Read the initial volume of the buret and record it in the data table.
B: Titration
1. Place the flask with acid and phenolphthalein under the buret. The buret tip should be
down about 1 cm inside the mouth of the flask to avoid losing any of the base. Place a
sheet of white paper under the flask to help highlight the pink indicator colour.
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2. Drip the base into the flask while swirling the flask to stir it. Add base quickly at first, but
as the pink colour starts to last longer, slow the drip rate. It is better to take longer on this
part than to need to start again. When the whole flask flashes pink before turning clear
again, add only one drop at a time and swirl the flask before adding more. Occasionally
rinse down the splashes on the side of the flask using a little distilled water from a wash
bottle (this water will not change your results).
3. When the faintest pink colour persists for 30 seconds, stop and record the final volume in
the buret.
4. Repeat the titration for trial 2.
5. Flush all chemicals down the sink and clean out the buret as described in step 5.
6. Have your instructor sign your data.
7. Clean up your work area and wash your hands.
Observations:
Mass of oxalic acid used:
Trial 1:____________g
Trial 2:____________g
Table 1: Titration of H2C2O4 with unknown concentration of NaOH
Volume of NaOH (mL):
Initial
Trial 1
Trial 2
Final
Used
Questions:
1. Why is it difficult to see whether you have added the phenolphthalein to the flask
solution before you have titrated it?
2. Why did you need to rinse the buret with base before you completely fill it with NaOH?
3. Write and balance the neutralization equation for the reaction performed in this
experiment.
4. Does the amount of water in which you dissolved the oxalic acid affect the outcome of
the experiment? Explain.
5. Why is it important to wash down the sides of the reaction flask with water as you near
the equivalence point?
6. In this experiment, how many moles of base are needed to neutralize one mole of acid?
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7. Determine the molarity of the unknown NaOH:
a. Calculate the number of moles of oxalic acid used.
b. Determine the number of moles of base needed to neutralize the moles of oxalic acid.
c. Knowing the volume of NaOH you used in each trial, calculate the molarity of the
NaOH.
d. Find the average [NaOH] from the two trials.
Conclusion:
What is the concentration of NaOH that was used in this experiment?
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