Analía Bellizzi – Chemistry Classes

Ronald Reagan High School

AS % OF NaClO in Commercial Bleach

Determining the Percent Sodium Hypochlorite in Commercial Bleach

Objectives

  • Additional practice in the technique of titration, this time a more complicated titration not involving acid-base chemistry.
  • Introduction a new piece of precision volumetric glassware, the volumetric flask; additional practice in the use of the pipet and buret.
  • Additional illustration of the use of an indicator, a substance which produces a detectable change when a reaction is complete.

Titration

All titrations are essentially exercises in stoichiometry. Titrations allow one to measure how much of a given substance is present in a sample by reacting that substance quantitatively with a known quantity of another reagent.

Conceptually, titration with an indicator is rather simple:

  1. Clean glassware
  2. Put carefully measured sample into flask
  3. Add indicator
  4. Carefully add titrant from buret until endpoint is reached and measure amount of titrant added.
  5. Relate moles of titrant used to moles of analyzed substance in sample through reaction stoichiometry.

Titrating Vinegar was conceptually and procedurally simple. The sample (vinegar) contained an unknown quantity of acetic acid, to which you carefully added base until the reaction was complete. The reaction in that acid-base titration was:

CH3COOH + NaOH –> CH3COO + Na+ + H2O .

When the reaction was complete, the number of moles of base used equals the number of moles of acid originally present; knowing the number of moles of acid allowed you to compute the concentration of acid in the vinegar sample.

Conceptually, this week’s lab is no different, although the procedure is a bit more elaborate. This week’s sample (bleach) contains an unknown quantity of sodium hypochlorite (NaOCl), which we convert completely to iodine (I2). We determine how much iodine is formed (and therefore how much hypochlorite was in the bleach) by carefully adding sodium thiosulfate (Na2S2O3) solution to the iodine until its reaction is complete.

OCl + 2 I + 2 H3O+ –> I2 + Cl + 3 H2O
I2 + 2 S2O32- –> 2 I + S4O62- .

When the reaction is complete, the number of moles of thiosulfate added equals twice the number of moles of hypochlorite originally present; knowing the number of moles of hypochlorite will allow us to compute the concentration of hypochlorite in the bleach sample.

 

Procedure Overview

  1. Clean glassware.
    Rinse 10-ml pipet with commercial bleach, 25-ml pipet with diluted bleach solution (see next step), and buret with sodium thiosulfate.
  2. Put carefully measured sample into flask.
    This week, the sample must be prepared before it can be titrated with thiosulfate. This preparation involves two steps:

    • dilution of the commercial bleach solution
    • production of iodine from the hypochlorite in the bleach.

    The production of iodine is necessary because iodine is readily titrated with thiosulfate, whereas the substance we are really interested in (hypochlorite) is not readily titrated. Since we can produce one mole of iodine for every mole of hypochlorite in the sample, we can relate moles of titrant used to moles of hypochlorite present. The dilution of the commercial bleach solution is necessary because iodine is not very soluble, and we want all the iodine we are going to produce to stay in solution. But since we are interested in the concentration of hypochlorite in the original bleach solution, we must dilute that solution carefully and in a controlled manner. So we prepare the diluted bleach by pipetting 10 ml of commercial bleach into a 100-ml volumetric flask. Then we pipet 25 ml of the diluted bleaching solution to our sample flask, where we react it with KI and HCl solutions. (Note: we don’t have to measure the KI and HCl exactly as long as hypochlorite is the limiting reagent.)

  3. Add indicator.
    The indicator is added to signal the endpoint of the titration, that is, the endpoint of the reaction of thiosulfate with iodine. Starch serves as an indicator because it forms a dark blue complex with iodine, but that complex disappears (turning the blue solution colorless) when all the iodine is used up. This week, though, we cannot add indicator until the titration is nearly done; otherwise, so much of the starch-iodine complex will form that it will form a precipitate, and effectively remove iodine from the reach of thiosulfate.
  4. Carefully add titrant from buret until endpoint is reached and measure amount of titrant added.
    As mentioned just above, the starch indicator must be added just before the endpoint is reached, that is, when the solution’s brownish-orange color fades to pale yellow.
  5. Relate moles of titrant used to moles of analyzed substance in sample through reaction stoichiometry.
    moles hypochlorite titrated = moles iodine titrated = (moles thiosulfate used)/2.

 

Procedural tips

You may work in groups (as usual), but everyone should do at least one titration.

Sequence of operations:

  • dilute bleach
  • add diluted bleach to KI in sample flask
  • add HCl to sample only after buret is filled and you are ready to titrate
  • add starch to sample only when titrated sample is pale yellow.

Calculations

  • moles thiosulfate used = (molarity of thiosulfate)x(volume thiosulfate used)
  • mass NaOCl titrated = (moles hypochlorite titrated)x(molar mass NaOCl)
  • volume of commercial bleach titrated = (volume dilute bleach titrated)x(dilution ratio (10/100))
  • mass of commercial bleach titrated = (density of commercial bleach)x(volume of commercial bleach titrated)
  • % NaOCl in commercial bleach = (mass NaOCl titrated/mass commercial bleach titrated) x 100%
  • Cost comparison as set out in handou