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1. To determine the concentration of hypochlorite ion (ClO-) in a solution of bleach, 25.0 cm³ of the bleach solution were first diluted to exactly 500 cm³ - Leaving Cert Chemistry - Question 1 - 2021
Question 1
1. To determine the concentration of hypochlorite ion (ClO-) in a solution of bleach, 25.0 cm³ of the bleach solution were first diluted to exactly 500 cm³. Then a ... show full transcript
Worked Solution & Example Answer:1. To determine the concentration of hypochlorite ion (ClO-) in a solution of bleach, 25.0 cm³ of the bleach solution were first diluted to exactly 500 cm³ - Leaving Cert Chemistry - Question 1 - 2021
Step 1
Describe how the 25.0 cm³ sample of the original bleach solution, provided in a beaker, was diluted to exactly 500 cm³.
Answer
To dilute the 25.0 cm³ sample of the bleach solution, a rinsed volumetric flask is utilized. First, transfer the bleach solution into a 500 cm³ volumetric flask using a funnel. After that, rinse the beaker with deionized water, and pour this rinse into the volumetric flask to ensure all the bleach is transferred. Fill the flask to the 500 cm³ mark with deionized water, allowing the bottom of the meniscus to touch the mark. Finally, stopper the flask and gently invert it several times to mix the solution thoroughly.
Step 2
Why must excess KI be added to the dilute bleach?
Answer
Excess potassium iodide (KI) should be added to ensure that all the hypochlorite ions (ClO-) react completely. The addition of KI serves as a reducing agent, ensuring that the reaction proceeds to completion and that the amount of iodine (I₂) produced can be accurately measured during the titration with sodium thiosulfate.
Step 3
Step 4
Identify a suitable indicator for use in the titration stage of the analysis. In this laboratory work, what final colour change was observed in the conical flask at the end point?
Answer
A suitable indicator for the titration is starch, which forms a blue-black complex with iodine. At the endpoint of the titration, the blue-black colour disappears, indicating that all iodine has reacted and the solution changes to colourless.
Step 5
Find by calculation: (i) the number of moles of sodium thiosulfate used up in one of the titrations.
Answer
To find the number of moles of sodium thiosulfate ( ext{Na}_2 ext{S}_2 ext{O}_3) used:
Number of moles = Concentration (M) × Volume (L)
= 0.09 ext{ mol/L} × rac{21.4 ext{ cm}^3}{1000} ext{ L}
= 0.001926 ext{ mol}
Step 6
Find by calculation: (ii) the number of moles of iodine in 25.0 cm³ of the conical flask.
Answer
From the balanced reaction, we know:
I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆.
Thus, for every 2 moles of sodium thiosulfate, 1 mole of I₂ is required. Therefore,
Moles of I₂ = rac{0.001926 ext{ mol}}{2} = 0.000963 ext{ mol (of I₂) in 25.0 cm³ of the conical flask.}
Step 7
Find by calculation: (iii) the concentration of ClO- in the original bleach, expressed as a % (w/v).
Answer
To find the concentration of ClO- in the original bleach solution:
Firstly calculate the moles of ClO- based on the stoichiometry: 1 mole of ClO- produces 1 mole of I₂; thus:
Moles of ClO- = 0.000963 ext{ mol} (in the 25.0 cm³ conical flask) multiplied by the dilution factor of 20 (since the original sample was 25 cm³ diluted to 500 cm³).
= 0.000963 × 20 = 0.01926 ext{ mol in original bleach.}
Finally, we can convert this to a % (w/v) concentration:
Concentration (g/100 cm³) = moles × molar mass of ClO- (approximately 51.5 g/mol) / total volume × 100.
= rac{0.01926 imes 51.5}{100} = 0.0099 ext{ g/100 cm³} ext{ or } 4 ext{% (w/v)}.