Determining the Molarity or Normality of an Acid Using Titration
Titrations are a critical analytical technique used in chemistry to determine the concentration of a solution, particularly in acid-base chemistry. This article will discuss how to determine the molarity or normality of an acid using an example titration with sodium hydroxide.
Understanding the Basics of Titration
Titration involves the gradual addition of a solution of known concentration (the titrant) to a solution of unknown concentration (the analyte). In the case of acid-base titrations, a base of known molarity, such as sodium hydroxide (NaOH), is used to neutralize the acid. By tracking the volume of titrant added, the concentration of the acid can be accurately determined.
Assessing the Nature of the Acid
The first step in titrating an acid is to determine whether it is diprotic (can donate 2 protons) or monoprotic (can donate 1 proton). This knowledge is crucial for ensuring the correct stoichiometry is used during the titration. For example, a diprotic acid like sulfuric acid (H?SO?) will react with sodium hydroxide as follows:
H?SO?(aq) 2NaOH(aq) → Na?SO?(aq) 2H?O
While a monoprotic acid like hydrochloric acid (HCl) will react as:
HCl(aq) NaOH(aq) → NaCl(aq) H?O
The Titration Process
To begin the titration, a few drops of a suitable indicator, such as phenolphthalein, are added to the acid solution. Phenolphthalein is an acid-base indicator that remains colorless in acidic solutions and turns pink in basic solutions. This indicator is particularly useful for detecting the endpoint of the titration, where the pH changes significantly.
The titration itself involves slowly adding the sodium hydroxide solution to the acid solution while continuously stirring. The titration continues until the solution changes color, indicating the endpoint of the reaction. Typically, a single drop of the titrant near the endpoint is sufficient to cause a visible color change. This is often referred to as the “equivalence point” or “endpoint” in titrations.
Calculating the Molarity or Normality
The concentration of the acid can be calculated using the following formula:
M?V? M?V?
Where M? is the molarity of the titrant (NaOH), V? is the volume of titrant used, M? is the molarity of the acid, and V? is the volume of the acid. This equation can be rearranged to solve for M?.
It is important to note that for a diprotic acid, the concentration will be calculated as if twice as much base is being neutralized.
Example Calculation
Suppose you are titrating 25 mL of an unknown concentration of HCl with 0.1 M NaOH. Before reaching the endpoint, you add 1.2 mL of 0.1 M NaOH. At the endpoint, you add an additional 0.01 mL (one drop) of 0.1 M NaOH to achieve the color change.
Net NaOH added: 1.2 mL 0.01 mL 1.21 mL
Firstly, convert the volume of NaOH from mL to L:
1.21 mL 0.00121 L
Now, use the titration formula:
M?V? M?V?
(0.1 M) × (0.00121 L) M? × (0.025 L)
M? (0.1 × 0.00121) / 0.025 4.84 × 10?? M
Therefore, the molarity of the HCl solution is 0.000484 M.
Ensuring Precise Results
To ensure accurate results, it is essential to perform several trials and take the average of the results. Small errors in measurement can significantly affect the accuracy of the titration. Care should also be taken to ensure that the indicator is used in the correct amount and is not present in excess, as this can affect the endpoint detection.
Conclusion
Titrating with a known molarity base allows for the determination of the molarity or normality of an acid. Through proper titration techniques and calculations, a precise concentration can be determined. Understanding the nature of the acid and the appropriate use of indicators are crucial steps in obtaining accurate results.