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Practical # 3

The document outlines a procedure for standardizing a hydrochloric acid (HCl) solution using a 0.5 M sodium hydroxide (NaOH) solution through acid-base titration. It describes the theory behind titration, the chemical equation, the mole ratio, and the use of phenolphthalein as an indicator. The procedure includes rinsing equipment, measuring volumes, conducting multiple titrations, and calculating the molarity of HCl based on the volumes used.

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mfurqankhattak
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57 views1 page

Practical # 3

The document outlines a procedure for standardizing a hydrochloric acid (HCl) solution using a 0.5 M sodium hydroxide (NaOH) solution through acid-base titration. It describes the theory behind titration, the chemical equation, the mole ratio, and the use of phenolphthalein as an indicator. The procedure includes rinsing equipment, measuring volumes, conducting multiple titrations, and calculating the molarity of HCl based on the volumes used.

Uploaded by

mfurqankhattak
Copyright
© © All Rights Reserved
We take content rights seriously. If you suspect this is your content, claim it here.
Available Formats
Download as PDF, TXT or read online on Scribd
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Experiment # 3

Standardize the given HCl solution using 0.5 M NaOH

Theory: It is an acid-base titration. An acid–base titration is a method of quantitative analysis for


determining the concentration of an acid or base by exactly neutralizing it with a standard
solution of base or acid having known concentration. A pH indicator is used to monitor the
progress of the acid–base reaction. Alkalimetry is the specialized analytic use of acid-base
titration to determine the concentration of a basic (alkaline) substance; acidimetry is the same
concept applied to an acidic substance.

Equation: HCl + NaOH NaCl + H2O

Mole ratio: 1:1

Indicator: Phenolphthalein End point: Pink

Titration Procedure:

1. Rinse the burette with the standard solution, the pipette with the unknown solution, and the
conical flask with distilled water.
2. Place an accurately measured volume of the analyte into the Erlenmeyer flask using the
pipette, along with a few drops of indicator. Place the standardized solution into the burette
and indicate its initial volume in a lab notebook. At this stage, we want a rough estimate of
the amount of known solution necessary to neutralize the unknown solution. Let the
solution out of the burette until the indicator changes color and record the value on the
burette. This is the first titration and it is not very precise; it should be excluded from any
calculations.
3. Perform at least three more titrations, this time more accurately, taking into account where
the end point will roughly occur. Record the initial and final readings on the burette, prior
to starting the titration and at the end point, respectively. (Subtracting the initial volume
from the final volume will yield the amount of titrant used to reach the endpoint.)
4. The end point is reached when the indicator permanently changes color.
Calculations:

Mean volume of NaOH used = V2 =……… cm3

M1V1/n1 (HCl) = M2V2/n2( NaOH)

Molarity of HCl = M1= M2V2/n2 x n1/V1

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