Preparing a standardized chemical solution requires precision, especially when dealing with concentrated reagents. Creating 0.1 N Hydrochloric Acid (HCl) is a common laboratory procedure used for various quantitative analyses. HCl is frequently used as a titrant in acid-base titrations to accurately determine the unknown concentration of a basic substance. Achieving the correct concentration is important for reliable results in analytical chemistry.
Understanding Normality and Starting Materials
Normality (N) is a measure of concentration that represents the number of equivalents of solute per liter of solution. For Hydrochloric Acid, a monoprotic acid, the normality is equal to its molarity (M) because each molecule of HCl donates only one hydrogen ion in a reaction. Thus, a 0.1 N HCl solution is the same as a 0.1 M HCl solution.
The starting material for this preparation is concentrated hydrochloric acid, which is an aqueous solution of hydrogen chloride gas. Commercial concentrated HCl is typically sold with a concentration around 36% to 38% by weight (w/w). The specific gravity, or density, of this concentrated acid is usually in the range of 1.18 to 1.20 g/mL.
Since concentrated acid slowly loses hydrogen chloride gas over time, its exact concentration is variable. The precise concentration of the concentrated stock must always be taken directly from the bottle’s label or the accompanying Certificate of Analysis. This stated percentage and density are necessary to calculate the required volume for dilution.
Calculating the Required Dilution
The first step in calculating the dilution is to convert the concentrated acid’s weight/weight percentage into Normality. This conversion uses the density of the concentrated acid, its percentage purity, and the molar mass of HCl, which is approximately 36.46 g/mol. For example, a 37% w/w HCl solution with a density of 1.18 g/mL has a concentration of about 12.0 N.
Once the initial normality (\(\text{N}_1\)) is known, the dilution formula, \(\text{N}_1\text{V}_1 = \text{N}_2\text{V}_2\), is used to find the necessary volume (\(\text{V}_1\)) of the concentrated acid. Here, \(\text{N}_2\) is the target normality (0.1 N), and \(\text{V}_2\) is the final volume of the solution being prepared, often 1 liter.
Using the example of a 12.0 N concentrated acid, the volume required to make 1 liter of 0.1 N HCl is approximately 8.3 milliliters (\(\text{V}_1 = \frac{0.1 \text{ N} \times 1000 \text{ mL}}{12.0 \text{ N}} \approx 8.3 \text{ mL}\)). This calculation provides the theoretical volume of concentrated acid needed to achieve the target concentration. Measuring this small volume with high accuracy is important for creating a solution close to the intended 0.1 N concentration. Although the resulting solution will be close to the target, it is considered an approximate concentration until it is verified through a separate analytical procedure.
Step-by-Step Preparation Protocol
The practical preparation of 0.1 N HCl requires strict adherence to safety protocols. Mandatory personal protective equipment (PPE) includes chemical splash goggles, a lab coat, and appropriate chemical-resistant gloves. The entire dilution process must be performed inside a functioning fume hood to manage the highly pungent and corrosive hydrogen chloride fumes released by the concentrated acid.
The calculated volume of concentrated acid should be measured using a precise measuring device like a graduated pipette or a small graduated cylinder. Before adding the acid, a portion of the final required water volume, such as 100 to 200 milliliters, should already be in the volumetric flask. It is a firm rule to always add acid to water, never the reverse, because the dilution of concentrated acid is a highly exothermic process.
Adding the acid slowly to the water allows the heat generated to dissipate safely within the larger volume of the solvent. After the acid has been added and the solution mixed gently, the flask should be allowed to cool to room temperature.
Finally, the remaining distilled or deionized water is added carefully to bring the solution volume exactly to the etched calibration mark on the neck of the volumetric flask. Ensure the bottom of the meniscus aligns with the mark.
The volumetric flask is then sealed and inverted several times to ensure the homogeneous mixing of the acid and water. This careful mixing is necessary because the density difference between the concentrated acid and water can lead to layering. The prepared solution is then transferred to a clean, sealed storage container.
Standardization and Storage
The prepared 0.1 N HCl solution is considered a secondary standard because the exact concentration of the starting material was uncertain, and the volume measurements inherently contain a degree of error. Therefore, the solution must be standardized to determine its true, precise normality.
Standardization involves performing a titration against a primary standard, which is a substance of high purity and known concentration that is stable and non-hygroscopic. Common primary standards used for standardizing HCl include anhydrous sodium carbonate (\(\text{Na}_2\text{CO}_3\)) or Tris(hydroxymethyl)aminomethane (THAM). The titration reaction allows for the determination of the exact concentration of the prepared HCl solution, which is then recorded, often to four significant figures, for accurate analytical work.
For storage, the standardized 0.1 N HCl must be kept in a tightly sealed glass bottle to prevent evaporation and minimize contact with air. The container must be clearly labeled with the name of the solution, the exact standardized normality, the date of preparation, and the initials of the preparer. Proper labeling and storage ensure the solution remains reliable for use in precise laboratory procedures.