Buffer Solutions and pH Determination: Electrochemical and Colorimetric Analysis

The experiment aimed to achieve the following objectives: the preparation of diverse buffer solutions, determination of pH using both colorimetric methods with various liquid indicators and electrometric methods with a pH meter. A primary phosphate buffer solution was created using phosphoric acid (H3PO4) as a weak acid and primary sodium phosphate dihydrate (NaH2PO4.2H2O) as a conjugate base. The pH of the buffer solution was measured using a pH meter, and adjustments were made using 6M HCl and 6M NaOH to alter acidity or basicity.

The pH meter, relying on electrical potential, accurately measured pH levels and exhibited sensitivity to small changes with the addition of HCl or NaOH.

This precision surpassed the accuracy of pH paper readings. Colorimetric pH determination involved the use of acid-base indicators (Thymol blue, Bromophenol blue, Bromocresol green, Bromocresol purple, Phenol red, Methyl orange, Phenolphthalein). The varying color changes observed during colorimetric measurements allowed for the identification of substances based on their pH range.

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Different substances exhibit distinct pH levels, and acid-base indicators aid in narrowing down this range. The color-change intervals for each indicator were determined as follows: Thymol blue (2-3/8-11), Bromophenol blue (3-5), Bromocresol green (3-5), Bromocresol purple (5-7), Phenol red (5-8), Methyl orange (5-7), Phenolphthalein (8-11).

In our daily lives, we often encounter acids and bases present in various substances like fruits and cleaning agents. The definitions of acids and bases, proposed by Thomas M. Lowry and Johannes N. Bronsted in 1923, describe acids as proton releasers and bases as proton acceptors.

This distinction is crucial in understanding the behavior of different substances in solution. Strong acids release protons readily, while weak acids have less than 1% ionization in dilute solutions. Similarly, strong bases readily accept protons, while weak bases are poor proton acceptors.

To simplify the expression of low concentrations of hydrogen ions in weak acid solutions, S.P.L. Sorensen proposed the concept of pH in 1909, defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH=−log10​[H+].

The pH of solutions holds significance in biomedical sciences due to its influence on biomolecule functioning and the potential for small pH changes to cause metabolic disturbances. Controlling pH is vital in biomedical laboratory procedures involving biomolecule separation, purification, and assay for biological activity.

Buffer solutions play a critical role in maintaining pH stability. Buffers resist changes in pH by containing high concentrations of weak acids or bases along with their conjugate partners. These systems adhere to the Le Chatelier Principle, which governs equilibrium conditions.

Compounds Tested:

1. Distilled water
2. Phosphoric acid (H3PO4)
3. Primary sodium phosphate dihydrate (NaH2PO4.2H2O)
4. 6M HCl
5. 6M NaOH
6. Acid-base indicators (Thymol blue, Bromophenol blue, Bromocresol green, Bromocresol purple, Phenol red, Methyl orange, Phenolphthalein)

Procedure:

1. Preparation of Reagents:
   • 120mL of concentrated HCl (12.2M) was diluted to 250mL to obtain 6M HCl.
   • The computation: (12.2M)(xL)=(6.0M)(0.250L) yielded x=0.12L or 120mL of concentrated HCl.
2. Preparation of Buffer Solution:
   • A buffer solution was prepared using 0.95mL (19 drops) of Phosphoric acid (H3PO4) and 1.716g of Primary sodium phosphate dihydrate (NaH2PO4.2H2O) in a 250mL solution.
   • Moles of the buffer were calculated as (0.250L)(0.10M)=0.025 moles.
   • Henderson Hasselbach equation was applied: pH=pKa+log⁡pH=pKa+log, where pKa=2.00pKa=2.00.
   • This yielded a buffer molarity of 1.76M, with 0.011 moles of H2PO4−H2PO4− and 0.014 moles of H3PO4H3PO4.
   • The volume of H3PO4H3PO4 required was calculated as 0.95mL, and the mass of NaH2PO4.2H2ONaH2PO4.2H2O was found to be 1.716g.

Results and Discussion:

The experiment aimed to prepare buffer solutions and determine their pH both colorimetrically and electrometrically. The colorimetric determination involved using various acid-base indicators, each with its distinct pH range. The pH meter provided accurate readings, showcasing its superiority over pH paper.

The significance of pH in biomedical sciences was emphasized, stressing its influence on biomolecule functions and the sensitivity of various laboratory procedures to pH changes. Buffers were highlighted as crucial systems resisting pH changes, with practical applications in maintaining stability during experimental procedures.

This laboratory experiment successfully demonstrated the preparation of buffer solutions and the determination of pH using both colorimetric and electrometric methods. The importance of pH in biomedical sciences and the role of buffers in maintaining pH stability were highlighted. The results obtained contribute to a better understanding of pH control and its applications in various scientific processes.

The pH meter calibration was performed at pH 7 using neutral solutions such as distilled water. After calibration, the electrode was immersed in the primary phosphate buffer solution. For accurate readings, the standard buffer used should have a pH within two units of the expected pH of the test solution. If the pH reading was outside the desired range, adjustments were made using 6.0M HCl or 6.0M NaOH. The process was repeated until the pH of the buffer solution fell within the target range.

Seven test tubes were labeled and filled with 5mL of the prepared buffer solution. Each test tube received 2 drops of a specific acid-base indicator, and after shaking, the color change was noted.

Results and Discussion:

1. Electrometric pH Measurement: The pH meter, functioning based on electrochemical properties, accurately measured the hydrogen ion concentration in the prepared buffer solution. Slight additions of HCl or NaOH induced fluctuations in readings, with HCl decreasing and NaOH increasing the pH. The sensitivity of the electrode to changes in [H+] and [OH-] ions demonstrated the pH meter's superior accuracy compared to pH paper.

2. Colorimetric pH Determination: Organic substances exhibit color changes in dilute solutions at specific hydrogen ion concentrations. Acid-base indicators, such as phenolphthalein, change color depending on the pH. Methyl orange, a weak acid, shifts its equilibrium between the red and yellow forms based on changes in hydrogen ion concentration. The color-change interval of an indicator, such as methyl orange, is the pH range where a pronounced color change occurs. Acid-base indicators also reveal molecular characteristics, as color changes result from alterations in electron confinement.

Colorimetric Analysis, utilizing variations in color intensity, is employed to determine pH. The inherent properties of a substance or the formation of a product due to the addition of a suitable reagent or acid-base indicator contribute to the solution's color. pH determination is achieved by comparing the color intensities of the solution with unknown pH to those of solutions with known pH values.

In summary, the electrometric pH measurement with the pH meter showcased precision and sensitivity to slight pH changes. On the other hand, colorimetric pH determination with acid-base indicators demonstrated the dynamic nature of color changes corresponding to varying pH levels, providing a visual representation of the solution's acidity or basicity. Both methods contribute valuable insights into the characteristics and behavior of substances in solution.

The laboratory report has been condensed to maintain brevity while addressing the essential elements of the experiment. Additional details and expansions can be included based on specific requirements.

Table 4. Results of the colorimetric determination of pH