EXPERIMENT 13
Volumetric Analysis II:
Determination of Active Ingredients in Commercial Bleach
and Vinegar
Outcomes
After completing this experiment, the student should be able to:
1. Demonstrate the analysis of common chemicals using titration technique.
2. Explain the concept of quantitative analysis.
3. Quantitatively determine the amounts in an unknown using standard solutions.
4. Demonstrate the understanding of redox titrations.
Part A: Analysis of commercial bleach
Introduction
Companies selling chemical consumer products do extensive advertising in an attempt to convince the
consumer that their detergent is stronger, gets clothes whiter, and is safer for the environment. Of all
the formulations that have been placed on the market that claim to remove stains from most clothing,
perhaps the most effective is still simply bleach, such as Javex, Clorox, Purex, or some similar generic
(store) brand. In addition many laundry detergents also add ''whiteners'' which fluoresce in the
presence of ultraviolet radiation, giving the appearance of a whiter wash.
Commercial bleaching agents contain the hypochlorite ion, CIO−, as the ''active ingredient'' for
making clothes whiter or removing stains. This ion is generally in the form of the sodium salt, NaClO,
or the calcium salt, Ca(ClO)2.
The concentration of the oxidizing agent in a bleaching agent is determined in this volumetric titration
based on redox reactions. The oxidation-reduction analysis of a bleach involves the reaction of the
hypochlorite ion, an oxidizing agent, with iodide ion in the presence of an acid. The I− ion is oxidized
to I2. The reaction equation (1) is represented as follows:
ClO− (aq) + 2I− (aq) + 2H+ (aq) → I2 (aq) + Cl−(aq) + H2O (l)
(1)
Iodide ion (in the form of KI) is added in excess, so that all ClO− reacts (in other words, ClO− is the
limiting reagent). As given in the equation, 1 mol is formed for each mol of ClO-. and for each 1 mol
ClO− 1 mol I2 is formed. Some of the I2 will react with the excess I− to form the dark brown triiodide
ion:
I2 + I− → I3−
(2).
The liberated I2 is now titrated with standard sodium thiosulfate, Na 2S2O3. The reaction equation (3)
for the titration is represented as follows:
I2(aq) + 2 S2O32− (aq) → 2I− (aq) + S4O62− (aq)
(3)
These equations serve as the basis for determining the strengths of bleaching agents in this
experiment. From equations (1) and (3) we see that the net ionic equation indicates 1 mole ClO −
liberates 1 mol I2, which in turn reacts with 2 moles thiosulfate, S2O32−:
ClO− ≡ I2 ≡ 2 S2O32−
(4)
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Safety Precautions
Handle concentrated HCl in the fumehood only, wear gloves and safety glasses (as always!). Rinse
any spills immediately.
Materials and Equipment
Bleach, burette, 100 mL volumetric flask, 25 mL pipette, 250 mL Erlenmeyer flask, 2 g of KI, 10 mL
of 3 M H2SO4 Standard Na2S2O3,
Procedure
Part A: Analysis of commercial bleach solution
1. Pipet 10.0 mL of bleach (the density of liquid bleach is 1.084 g/mL) into a 100 mL volumetric
flask and add about 50 mL of distilled water. Shake it well and fill it to the mark with distilled
water. Mix thoroughly by inverting the stoppered volumetric flask a number of times. Pipet 25.0
mL of this diluted bleach solution into a 250 mL Erlenmeyer flask and add 20 mL of distilled
water, 2 g of KI (use approximate weight), and 10 mL of 3 M H2SO4. A yellow color indicates the
presence of I2 as I3−.
2. Fill the buret with standard Na2S2O3 solution (provided), record the molarity of the Na2S2O3
solution and the initial buret reading V1. Titrate the liberated I2 with your standard Na2S2O3
solution until the solution becomes light yellow. Now add 2 mL of starch solution to form the
deep-blue color. While swirling the flask continue titrating very slowly until the deep-blue color
disappears and the solution becomes colorless. Record the final buret reading V2 and the volume
delivered, Vburet = V2 − V1 (±0.02 mL).
3. Repeat the titration with a second 25.0 mL sample from the solution in the 100 mL volumetric
flask.
4. Calculate the amount of available hypochlorite ion in the sample, and from that the weight% of
NaClO in the commercial bleach.
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Part B: Determination of weight% acetic acid in commercial vinegar
Introduction
Acetic acid, CH3COOH, is the acid component of vinegar. Its concentration varies slightly in
different samples but must be at least 4% by mass (acetic acid in water) to meet the minimum
commercial standards. Concentration of acetic acid may even exceed 5% in some vinegar products.
Caramel flavoring and coloring may also be added to make the product more pleasing to the
consumer
The percent by mass of CH3COOH in vinegar is determined with an acid-base titration using standard
NaOH. Phenolphthalein is used as an indicator to determine the end point which gives pink color in
basic solution and is colorless in acidic solution. The amount of CH3COOH present in the vinegar
is calculated from the balanced equation:
CH3COOH(aq) + OH−(aq) → H2O(l)+ CH3COO−(aq)
(5)
−
At the equivalence point the moles of OH dispensed from the buret equals the moles of
CH3COOH in the vinegar.
mol OH− = mol CH3COOH
Mass of CH3COOH = mol CH3COOH x molar mass of CH3COOH
% CH3COOH =
Mass of CH 3 COOH
x100%
Mass of vinegar sample
Procedure Part B: Weight% acetic acid in a commercial vinegar
1. Determine the mass (±0.01 g) of a 125 mL (or 250mL ) Erlenmeyer flask. Do not tare the
balance, but record the actual weight in your notebook. Add 5.0 mL (use Mohr pipet) of a selected
brand of vinegar to the Erlenmeyer flask and weigh again (± 0.01 g). Record the mass of the 5.0
mL vinegar. Be sure to use the same balance.
2. Add 2 drops of phenolphthalein to the vinegar in the Erlenmeyer flask and wash down the wall
of the flask with 20 mL of de-ionized water.
3. Obtain standardized NaOH solution (provided, make sure to record the M NaOH). Rinse and fill the
the buret.
4. Titrate the sample to the phenolphthalein endpoint.
5. Determine the weight% acetic acid in the vinegar and then compare with that shown on the label
on the vinegar bottle.
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Experimental General Chemistry 1
Experiment 12: Applications of Volumetric analysis
Laboratory Data Sheet
Name: ______________________________________________ Section:
_________
Part A: Analysis of commercial bleach
Trial
mL
diluted
bleach
1
25.0
2
25.0
3
25.0
V1
V2
Volume
(mL) of
Na2S2O3
(V2-V1)
Volume in L
of Na2S2O3
(Vburet)
Moles
ClO− in
25.0 mL
% of
NaClO in
bleach
Average weight % NaClO:
Calculation (repeat for each of the three titrations)
1. From titration, we find # moles of ClO− in 25.0 mL pipetted and titrated sample
equals 0.5 x moles Na2S2O3 = 0.5 x MNa2S2O3 x Vburet
= ________ moles
2. Total moles ClO− in 10.0 mL original sample = (100/25)x(moles in #1 above)
= _________ mol
3. Mass of NaClO in 10.0 mL original sample
= ___________ g
4. Mass of 10.0 mL bleach (mass = dxV), d = 1.084 g/mL)
= ___________
5. Wt% of NaClO =
mass of NaClO ( from 3)
x100%
mass of bleach sample ( from 4)
= ___________
Calculate the Wt% values in the table for each titration, calculate the average Wt% and enter in
the table above.
Calculate:
Standard Deviation, SD =
Show your work:
% Standard Deviation, SD% =
Show your work:
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Part B: Weight percent of acetic acid in vinegar
Trial
Mass of
Vinegar
V1 (mL)
V2 (mL)
Volume
of NaOH
(mL)
(V2-V1)
Volume
of NaOH
(L)
Moles
NaOH
Moles
CH3CO
OH
Mass of
CH3CO
OH
% of
CH3CO
OH in
vinegar
1
2
3
Average % CH3COOH
Calculations: (repeat for each of the three titrations)
1. Moles of NaOH = M of NaOH x Volume in L
2. Moles of CH3COOH in sample = moles of NaOH used
3. Mass of CH3COOH = molar mass of CH3COOH x # of moles
4. Wt% of CH3COOH Vinegar =
(from #2 above)
mass of CH 3COOH
x100%
mass of vinegar sample titrated
Calculate the Wt% values in the table for each titration, calculate the average Wt% and enter in
the table above.
Calculate:
Standard Deviation, SD =
Show your work:
% Standard Deviation, SD% =
Show your work:
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