Acids, Bases, and pH - SS2 Chemistry Past Questions and Answers - page 3

Acid-base indicators are substances that exhibit different colours in acidic and basic environments, allowing us to determine the pH of a solution. They are usually weak acids or bases that undergo a reversible colour change as the pH of the solution changes.

Acid-base indicators function through the concept of conjugate acid-base pairs. In their acidic form, indicators are protonated and have one colour, while in their basic form, they are deprotonated and have a different colour. The equilibrium between the acidic and basic forms shifts as the pH of the solution changes, resulting in a visible colour change.

The choice of an appropriate indicator depends on the desired pH range. Different indicators have different pH ranges over which they undergo a colour change. It is crucial to select an indicator that changes colour near the pH of the solution being tested. For example:

●     Methyl orange has a range of pH 3.1-4.4 and is suitable for determining the pH of acidic solutions.

●     Phenolphthalein has a range of pH 8.2-10.0 and is commonly used for neutralisation reactions or to determine the endpoint of acid-base titrations.

●     Bromothymol blue has a range of pH 6.0-7.6 and is useful for determining the pH of slightly acidic to slightly basic solutions.

The choice of indicator is based on the principle that the indicator's pKa (the pH at which it changes colour) should be close to the pH of the solution being tested. This ensures that the indicator is sensitive to the pH change and provides an accurate representation of the solution's acidity or basicity.

While acid-base indicators are useful tools for approximating the pH of a solution, they have certain limitations that should be considered:

1.    Limited pH Range: Each indicator has a specific pH range over which it changes colour. Outside this range, the indicator may not provide accurate results. Using an indicator that is not suitable for the pH range being tested can lead to incorrect pH readings.

2.    Subjectivity of Colour Perception: Colour changes in indicators are often subtle and subject to individual interpretation. Different observers may perceive colour changes differently, leading to variations in recorded pH values.

3.    Interference from Other Substances: Some substances present in the solution can interfere with the colour change of the indicator, resulting in inaccurate readings. Substances that absorb light in the same wavelength range as the indicator can affect the observed colour.

In situations where higher accuracy is required, pH metres or universal indicators may be preferred over specific acid-base indicators:

●     pH Metres: pH metres offer precise and quantitative measurements of pH by measuring the electrical potential difference between electrodes immersed in the solution. They provide real-time and digital pH values, eliminating subjective interpretation and offering a wider pH range for accurate readings.

●     Universal Indicators: Universal indicators are mixtures of several acid-base indicators that provide a colour spectrum over a wide pH range. They offer a rough estimation of pH across a broader range of acidity or basicity, making them suitable for quick assessments without the need for precise measurements.

In the given acid-base titration scenario, where a strong base (NaOH) is being titrated against hydrochloric acid (HCl), the suitable indicator is phenolphthalein. Phenolphthalein is commonly used for strong acid-strong base titrations and undergoes a colour change from colourless to pink at a pH range of approximately 8.2 to 10.0, which is suitable for the equivalence point of the titration.

In an acid-base titration, the balanced equation for the reaction between acetic acid (CH3COOH) and sodium hydroxide (NaOH) is:

CH3COOH + NaOH → CH3COONa + H2O

From the given information, the volume of NaOH required to reach the endpoint is 27.5 mL. The stoichiometry of the reaction is 1:1, indicating that the moles of NaOH used are equal to the moles of CH3COOH present in the solution.

Moles of CH3COOH = Moles of NaOH

Moles = (Volume in litres) x (Molarity)

Moles of CH3COOH = (27.5 mL / 1000 mL/L) x (Molarity of NaOH)

Assuming the molarity of NaOH is 0.1 M, we can calculate the molarity of the acetic acid solution as follows:

Moles of CH3COOH = (27.5 mL / 1000 mL/L) x (0.1 M)

Moles of CH3COOH = 0.0275 mol

The volume of the acetic acid solution is given as 25.0 mL, which is equivalent to 0.025 L. Therefore, the molar concentration of the acetic acid solution is:

Molarity = Moles / Volume in litres

Molarity = 0.0275 mol / 0.025 L ≈ 1.1 M

Rounded to two significant figures, the molar concentration of the acetic acid solution is approximately 0.3

Acid-base titrations are analytical techniques used to determine the concentration of an unknown acid or base solution by reacting it with a standardised solution of known concentration. The process involves gradually adding the titrant solution (of known concentration) to the analyte solution (of unknown concentration) until a chemical reaction between the acid and base is complete.

In an acid-base titration, a burette is used to deliver the titrant solution in measured volumes to the analyte solution in a flask. The reaction between the acid and base is monitored using an indicator, which changes colour when the equivalence point is reached. The equivalence point occurs when the stoichiometrically equivalent amounts of acid and base have reacted. At this point, the reaction is complete, and the indicator undergoes a visible colour change.

The titration curve is a graph plotting the pH (or volume of titrant) against the volume of titrant added. It exhibits characteristic shapes based on the type of acid and base involved. The curve typically shows a gradual change in pH at the beginning, followed by a steeper slope near the equivalence point, and finally levels off after the equivalence point.

Indicators play a crucial role in acid-base titrations by providing a visual indication of the equivalence point. They undergo a colour change within a specific pH range, typically close to the equivalence point. Common indicators include phenolphthalein (pH range around 8.2-10) for weak acid-strong base titrations and methyl orange (pH range around 3.1-4.4) for strong acid-strong base titrations. By observing the colour change of the indicator, the endpoint of the titration can be determined, which corresponds to the equivalence point.

Several factors can impact the accuracy and precision of acid-base titrations, and it is important to control or minimise them to obtain reliable results:

1.    Instrument Calibration: Accurate calibration of the equipment, including burettes and pH metres, is crucial for precise volume and pH measurements. Regular calibration using appropriate standards ensures reliable readings and minimises errors.

2.    Indicator Selection: The choice of indicator should be appropriate for the titration, ensuring that the pH range of the indicator's colour change matches the expected equivalence point. Selecting the correct indicator minimises errors associated with misinterpretation of colour changes.

3.    Analyte and Titrant Concentrations: Accurate determination of analyte and titrant concentrations is crucial for calculating the unknown concentration. Precise measurements of the initial volumes and concentrations of the solutions minimise errors in calculations and improve accuracy.

4.    Titration Speed: The speed at which the titrant is added can affect the accuracy of the endpoint determination. Rapid or slow addition of the titrant can lead to overshooting or undershooting the equivalence point, respectively. Controlling the titration speed ensures accurate endpoint detection.

5.    Stirring and Mixing: Effective stirring and mixing of the reaction mixture ensure homogeneity and uniformity, facilitating rapid and complete reactions. Proper mixing minimises errors arising from uneven distribution of reactants and allows for a more precise determination of the equivalence point.

6.    Replicates and Averaging: Performing multiple titrations and calculating the average helps minimise random errors and improves precision. The more replicates performed, the more reliable the results become.

7.    Environmental Conditions: Temperature and atmospheric conditions can influence the accuracy of titrations. Temperature variations affect reaction rates and volumes, while exposure to atmospheric gases can alter the pH of solutions. Controlling and maintaining stable environmental conditions minimises errors due to such factors.

By controlling and minimising these factors, the accuracy and precision of acid-base titrations can be improved, resulting in reliable and consistent results. Adhering to good laboratory practices, proper technique, and attention to detail are essential in achieving accurate and precise titration measurements.