206 DIY Science: Illustrated Guide to Home Chemistry Experiments
HCl has reacted with the sodium bicarbonate to form one mole
each of sodium chloride, carbon dioxide, and water. The pH at
the equivalence point of this reaction happens to correspond
very closely to the pH range where methyl orange changes
color, but is well below the pH range of phenolphthalein.
Because these reactions occur with exactly a 2:1 proportion
of hydrochloric acid, the amount of titrant needed to reach
the indicated equivalence point with methyl orange is exactly
twice as much as the amount of titrant needed to reach the
indicated equivalence point with phenolphthalein. In other
words, if you use methyl orange, you miss the first equivalence
point completely. If you use phenolphthalein, you are misled
into believing that the first equivalence point is the final
equivalence point. For this reason, the neutralization of sodium
carbonate with hydrochloric acid is often used as a (literal)
textbook example of the importance of choosing the proper
indicator.
Note that for titrations with multiple equivalence points, it
makes a difference which solution is used for the titrant. For
example, if we titrate a solution of sodium carbonate by adding
hydrochloric acid, we could observe the first (higher pH)
equivalence point by using phenolphthalein as an indicator.
When sufficient HCl has been added to convert the sodium
carbonate to sodium bicarbonate, the phenolphthalein
changes from pink to colorless. We could then add some
methyl orange indicator to observe the color change from
yellow to red at the second (lower pH) equivalence point,
when the sodium bicarbonate is converted to sodium
chloride. Conversely, if we titrate a solution of hydrochloric
acid with sodium carbonate as the titrant, the acid is always
in excess until sufficient sodium carbonate has been added
to completely neutralize the acid. In this case, we use
phenolphthalein as the indicator, because only one equivalence
point exists for this reaction, and it is at the higher pH when
sodium carbonate is in slight excess.
In this lab, we’ll standardize an approximately 1 M solution of
hydrochloric acid by titrating a known volume of the acid with a
sodium carbonate solution of known molarity.
POCEDURER
This laboratory has three parts. In Part I, you’ll make up
a stock reference solution of 1.500 M sodium carbonate.
In Part II, you’ll use serial dilution to make up a working
reference solution of 0.1500 M sodium carbonate solution
to use as a titrant. In Part III, you’ll use that titrant to
standardize an approximately 1 M bench solution of
hydrochloric acid by titration.
PRTI: A mA kE UpA SToCk REfERENCE SoLUTIoN
of ~1.500 m SodIUm CARBoNATE
1.f you have not already done so, put on your splash I
goggles, gloves, and protective clothing.- Place a weighing paper on the balance and tare the
balance to read 0.00 g. - Weigh out about 15.90 g of anhydrous sodium carbonate
powder, and record the mass to 0.01 g on line A of
Table 11-4. - Using the funnel, transfer the sodium carbonate to the
100 mL volumetric flask. - Rinse the funnel with a few mL of distilled or deionized
water to transfer any sodium carbonate that remains in
the funnel into the volumetric flask. - Fill the volumetric flask with distilled or deionized water
to a few cm below the reference line. - Stopper the flask and invert it repeatedly until all of the
sodium carbonate dissolves. - Finish filling the volumetric flask with water until the
bottom of the meniscus is just touching the reference line. - Stopper the flask and invert it several times to mix the
solution thoroughly.
Transfer the sodium carbonate solution to the 100 mL
storage bottle and cap the bottle.
Use the actual mass of sodium carbonate from step 3
to calculate the actual molarity of the sodium carbonate
solution to the appropriate number of significant figures,
and record that molarity on line B of Table 11-4.
Label the bottle with its contents, molarity, and the date.
10.
11.
12.
PRTII: A EUERIALS S dILUTIoN To mAkE Up A
woRkING REfERENCE SoLUTIoN of ~0.1500 m
SodIUm CARBoNATE
1. If you have not already done so, put on your splash
goggles, gloves, and protective clothing.- Rinse the 100 mL volumetric flask, first with tap water
and then with distilled or deionized water. - Use the 10 mL pipette to transfer about 10 mL of the
~1.500 M sodium carbonate solution you made up in
Part I to the volumetric flask. Record the volume to
0.01 mL on line C of Table 14-4. - Fill the volumetric flask with water until the bottom of
the meniscus just touches the reference line. - Stopper the flask and invert it several times to mix the
solution thoroughly. - Use the actual volume of sodium carbonate solution
from step 3 to calculate the actual molarity of the
working sodium carbonate solution to the appropriate
number of significant figures, and record that molarity
on line D of Table 11-4.