362 DIY Science: Illustrated Guide to Home Chemistry Experiments
dISoALp S :
Return the unused chlorine bleach from part I to the
original bottle. The waste solutions from this lab can be
flushed down the drain with plenty of water.about 0.0007 moles of sodium hypochlorite. In the first redox
equation, one mole of sodium hypochlorite reacts with two
moles of iodide ion to form one mole of molecular iodine
(I 2 ), so the stoichiometric ratio of hypochlorite:iodine is 1:1.
In the second redox equation, one mole of molecular iodine
reacts with two moles of sodium thiosulfate to produce two
moles of iodide ions, so the stoichiometric ratio of iodine:
thiosulfate is 1:2. It follows, then, that the stoichiometric ratio
for hypochlorite:thiosulfate is 1:2, which means that two moles
of thiosulfate ion react stoichiometrically with one mole of
molecular iodine.
Our sodium thiosulfate titrant contains 0.1 mol/L or 0.0001
mol/mL. Because two moles of thiosulfate are required per mole
of hypochlorite and the chlorine bleach contains about 0.0007
moles of hypochlorite, about 14 mL of titrant should be required
to neutralize 1 g of chlorine bleach. As we determined in Part I,
the density of chlorine bleach is slightly more than 1 g/mL (my
observed value was 1.0841 g/mL), so just over 15 mL of titrant
should be required per mL of chlorine bleach. Using a 3 mL
sample of bleach would require 45+ mL of titrant, uncomfortably
close to the 50 mL capacity of our burette. We decided to use a
2.0 mL sample, which should require just over 30 mL of titrant,
well within the capacity of our burette.
- If you have not already done so, put on your splash
goggles, gloves, and protective clothing. - Weigh about 2.0 g of potassium iodide and add it to
about 25 mL of distilled water in a 125 mL Erlenmeyer
flask. (Alternatively, you can use 25 mL of a 0.5 M bench
solution of potassium iodide.) - Swirl the contents of the flask until the potassium iodide
dissolves and then, with swirling, add about 5.0 mL of
glacial (concentrated) acetic acid to the flask. - Use the pipette to transfer 2.00 mL of the chlorine bleach
sample to the flask. The solution should immediately turn
dark brown as the chlorine bleach oxidizes the iodide ions
to free iodine. Record the volume of the sample to 0.01 mL
on line E of Table 20-2. - Rinse the 50 mL burette with a few mL of 0.1 M sodium
thiosulfate, and allow it to drain into a waste container. - Transfer 50 mL or slightly more of 0.1 M sodium
thiosulfate to the burette. Drain a couple mL into the
waste container, until the level of the solution is below the
zero index mark on the burette, and make sure that there
are no bubbles in the burette, including the tip. Record
the initial volume as accurately as possible, interpolating
between the graduation marks on the burette, and record
that volume on line F of Table 20-2. - We calculated that we should need about 30 mL of
sodium thiosulfate titrant, so begin by running 25 mL or
so of titrant into the receiving flask with constant swirling.
As you near the endpoint, the intense brown color of the
solution in the flask will begin to fade noticeably. When
you reach that point, slow the addition rate to a fast drip,
and continue swirling.- When the solution in the flask reaches a pale brown color,
stop adding titrant. - Add about 2 mL of the starch indicator to the flask and
swirl the flask to mix the contents, which immediately
take on an intense blue-black color due to the formation
of a starch-iodine complex.
Continue adding sodium thiosulfate titrant dropwise, with
swirling. When the color of the solution fades to a slight
bluish tint, continue swirling the flask for 30 seconds.
(See Figure 20-1.) If the blue tint persists, add one more
drop of titrant and swirl the solution. The endpoint occurs
when the blue tint disappears.
Record the final volume as accurately as possible,
interpolating between the graduation marks on the
burette, and record that volume on line G of Table 20-2.
Determine the volume of titrant used by subtracting
the initial volume reading from the final volume reading.
Record that value on line H of Table 20-2.
Using the actual molarity of your nominal 0.1 M sodium
thiosulfate titrant, calculate the number of moles of
sodium thiosulfate required to reach the endpoint and
enter that value on line I of Table 20-2.
Sodium thiosulfate reacts stoichiometrically with iodine
at a 1:2 mole ratio, so the number of moles of iodine (and
therefore hypochlorite) present in the sample is half the
number of moles of sodium thiosulfate needed to reach
the endpoint. Calculate the number of moles of sodium
hypochlorite present in the sample by halving the number
of moles of thiosulfate required, and enter that value on
line J of Table 20-2.
Determine the mass of sodium hypochlorite present in
the sample by multiplying the number of moles present
by the gram molecular weight of sodium hypochlorite
(74.4422 g/mol). Enter your calculated value on line K of
Table 20-2.
Using your values for the density of chlorine bleach (line
D) and the volume of the chlorine bleach sample (line E),
calculate the mass of the chlorine bleach sample and
enter the result on line L of Table 20-2.
Calculate the mass percentage of sodium hypochlorite in
the chlorine bleach sample, and enter the result on line M
of Table 20-2.