Illustrated Guide to Home Chemistry Experiments

Chapter 20 Laboratory: Quantitative Analysis 361

SBSTITUTIU oNS ANd modIfICATIoNS

• You may substitute any suitable container of similar
  size for the 125 mL Erlenmeyer flask.


• You may substitute a 1.00 mL or 10.0 mL pipette for the
  2.0 mL pipette at the expense of some loss in accuracy.


• Ideally, the 0.1 M sodium thiosulfate solution should
  be standardized, but doing so is time-consuming and
  requires the relatively expensive chemical potassium
  iodate. You can make do with a nonstandardized
  sodium thiosulfate solution. Make a 0.1 M solution
  of sodium thiosulfate by adding 1.58 g of anhydrous
  sodium thiosulfate or 2.48 g of sodium thiosulfate
  pentahydrate to your 100 mL volumetric flask and
  making up the solution to 100.0 mL.


• Starch indicator solution does not keep for more than
  few days even if refrigerated, so you should make it up
  fresh for each laboratory session. To so so, mix about
  1 g of cornstarch or other starch with about 25 mL
  of cold water until it forms a paste. Add this paste to
  about one liter of boiling water and stir until it forms a
  clear suspension. Allow the liquid to cool and use it as
  is. This indicator is used in such small quantity that it’s
  fine to make it up with tap water rather than distilled
  water. You’ll need only a couple mL for each titration.
  Alternatively, you can use water in which potatoes,
  macaroni, or other pasta has been boiled.

CUTIOA nS

Chlorine bleach is corrosive and a strong oxidizer. Glacial

acetic acid is corrosive. Wear splash goggles, gloves, and

protective clothing.

z

dAR. Rm y CHERvENAk CommENTS:

Back titration is often used when an appropriate indicator

isn’t available. For example, aqueous silver ions react

strongly with halides, but indicators don’t respond to the

presence of silver ions. Back-titrating the excess silver

ions with thiocyanate will enable the determination of

the halide ion concentration. Ferric chloride acts as an

indicator for the back titration—the ferric ion forms a

colorless complex in water, but the ferric thiocyanate

complex is bright red.

very low concentrations, aqueous iodine reacts with a starch

solution to produce an intense blue color. By using starch as an

indicator, we can judge the endpoint very precisely. We’ll begin

the titration without the indicator. When we’ve added enough

sodium thiosulfate titrant to reduce the brown coloration to a

faint tint, we know the titration is near the endpoint. We’ll then

add starch indicator, which turns the solution bright blue, and

continue adding sodium thiosulfate titrant dropwise until the blue

color disappears.

POCEDURER

This lab session is in two parts. In Part I, we’ll determine the

density of chlorine bleach solution, which we’ll need later to

calculate the mass percentage of sodium hypochlorite in the

sample. In Part II, we’ll react the chlorine bleach sample with a

solution of potassium iodide in acetic acid, and titrate the iodine

formed by that reaction with sodium thiosulfate titrant.

PRTI: A TERNEdE mI THE dENSITy of CHLoRINE BLEACH

ATERNATIL vE pRoCEdURE

If you don’t have a 100 mL volumetric flask, or if the

mass of the filled flask would exceed the capacity of your

balance, you can substitute a 25 mL or 50 mL volumetric

flask. Alternatively, tare a small beaker or a foam cup on

your balance and use a 10 mL or larger pipette to transfer

a measured volume of bleach to the beaker. The value you

obtain for density will be less accurate using the smaller

volume, but is sufficiently accurate for this experiment.

In Part I, we’ll determine the mass of a 100.00 mL sample of

chlorine bleach, which allows us to calculate its density.

1. If you have not already done so, put on your splash
   goggles, gloves, and protective clothing.


2. Weigh a clean, dry, empty 100 mL volumetric flask and
   record its mass to 0.01 g on line A of Table 20-2.


3. Use the funnel to fill the volumetric flask with the chlorine
   bleach sample nearly to the index line. Use the dropper or
   Beral pipette to add the last mL or two of chlorine bleach,
   bringing the level of the solution up to the index mark on
   the flask.


4. Reweigh the filled flask and record its mass to 0.01 g on
   line B of Table 20-2.


5. Subtract the empty mass of the flask from the combined
   mass of the flask and chlorine bleach, and enter the mass
   of the chlorine bleach on line C of Table 20-2.


6. Calculate the density of chlorine bleach in g/mL and enter
   that value on line D of Table 20-2.

PRTII: A dNETERE mI HypoCHLoRITE CoNCENTRATIoN

Before we begin the analysis, we need to decide what sample

size is appropriate for our quantitative test. We know that 1 g

of chlorine bleach nominally contains 0.0525 g of sodium

hypochlorite. The gram molecular mass of sodium

hypochlorite is 74.44 g/mol, so 1 g of chlorine bleach contains