Acid-Base Titration Study Notes
Acid-Base Titration Overview
1. Definition and Purpose
Titration: A technique used to determine the concentration of an unknown solution called the analyte.
- It employs a titrant, a solution of known concentration, to react with the analyte.
- The process includes the addition of titrant from a burette to the analyte solution until the reaction completes.
- The completion is typically indicated by a color change, referred to as the endpoint.
2. Key Terms
Titrant: Solution in the burette with a precisely known concentration, also known as a standard solution.
Titrand or Analyte: The unknown concentration solution present in the conical flask.
End-Point: The point in titration when the indicator changes color (does not necessarily coincide with the equivalence point).
Equivalence Point: The theoretical point where the amount of titrant added is stoichiometrically equivalent to the amount of analyte present, defined as:
.Indicator: An organic compound, either a weak acid or weak base, added to the solution that changes color at or near the equivalence point.
Blank Titration: A titration process performed without the analyte to estimate titration error.
- All reagents are included except the analyte, and the result is used to adjust the volume measured during the real titration.
3. Titration Curves
3.1 Definition
A titration curve is a plot illustrating the relationship between the pH of the analyte solution and the volume of titrant added (in cm³ or mL).
3.2 Example
When titrating 20.00 mL of 0.100 M HCl (strong acid) with 0.100 M NaOH (strong base), the titration curve displays a significant change in pH around the equivalence point.
Indicators in Acid-Base Titrations
1. Importance of pH Range and Transition
The pH range of indicators signifies the extent of color change at certain pH levels, typically covering ±1 of its pKa value.
- For instance, an indicator is yellow when pH is lower than pKa - 1, and turns red to greater than pKa + 1.
Indicator | Colour Range | pKa |
|---|---|---|
Litmus | Red at < pH 7, Blue at > pH 7 | 6.5 |
Methyl Orange | Red (pH < 3.1) to Yellow (pH > 4.4) | 4.4 |
Phenolphthalein | Colourless (pH < 8) to Pink (pH > 10) | 9.6 |
|
2. Transition Mechanism
Acid-base indicators can be represented generally as weak acids or bases, for example, HIn (acid form) changes form to In− (conjugate base), influencing the solution's colour.
The general dissociation reaction is:
3. Color Change Mechanism
At low pH, high concentration of H3O+ drives the equilibrium left, leading to the acid form's color appearing (Colour A).
At high pH, a low concentration of H3O+ shifts the equilibrium to the right leading to the base form's color (Colour B).
4. Effect of Concentration Ratios
Human perception of color differing in a mixture of [In−] and [HIn] is affected by their ratio, where colors appear constant when 10 times greater or less.
Indicators exhibit an acid color if:
Conversely, an indicator appears its base color if:
5. Analysis of Specific Indicators
Methyl Orange:
- Changes from red (acidic) to yellow (basic) across the pH range of 3.1 to 4.4, as expressed in the equilibrium reaction:
Phenolphthalein:
- Normally colourless in acidic solutions; becomes pink in basic solutions.
- The equilibrium represented as:
Litmus: Turns from red in acidic conditions to blue in basic conditions, typically around pH 7, making it useful for general acidity/basicity testing.
6. Conclusion
Understanding the mechanisms behind titrations, indicators, and the significance of various pH levels is crucial for accurate analytical results in chemistry. Indicators play a vital role in determining endpoints during acid-base titrations, where their color change corresponds to the reaction progress.