Analytical Chemistry Methods for Water Quality and Spectroscopy

EDTA Titration and Water Hardness Calculation

EDTA Structure

EDTA (C₁₀H₁₆N₂O₈) is a hexadentate ligand that binds metal ions through four carboxyl (–COOH) groups and two amine (–NH₂) groups.

Titration Procedure (Water Hardness Test)

1. Take a 50 mL water sample.
2. Add buffer (pH 10) and Eriochrome Black T indicator (resulting in a wine-red color).
3. Titrate with EDTA until the color changes to sky blue (the end point).
4. Note the volume of EDTA used (V).

Reaction

M²⁺ + EDTA⁴⁻ → [M-EDTA]²⁻

Calculation Formula

Hardness (ppm) = (V × M × 1,000,000) / Vₛₐₘₗₔₗₑ

• M = EDTA molarity
• V = Volume of EDTA used (mL)
• Vₛₐₘₗₔₗₑ = Sample volume (mL)

Water Impurities and Boiler Problems

Hardness

The presence of calcium (Ca²⁺) and magnesium (Mg²⁺) ions in water. Causes: Dissolved minerals like CaCO₃ and MgCl₂.

Temperature

Affects the solubility of salts and gases. Higher temperature can lead to increased scale formation.

Permanent Hardness

Cannot be removed by boiling. Caused by: Chlorides and sulfates of Ca²⁺ and Mg²⁺ (e.g., CaSO₄, MgCl₂).

Alkalinity

The capacity of water to neutralize acids (due to HCO₃⁻, CO₃²⁻, OH⁻). Causes: Dissolved bicarbonates, carbonates, and hydroxides.

Reverse Osmosis (RO)

A water purification method using a semipermeable membrane to remove ions. Removes: Hardness, salts, and other impurities.

Priming

The carrying of water droplets with steam in boilers. Causes: High water levels, high steam velocity, or impurities.

Foaming

Stable bubbles formed on the water surface in boilers. Causes: Presence of oil, alkali, or organic matter in the water.

Scales

Hard, adherent deposits on boiler surfaces. Causes: Precipitation of calcium sulfate or magnesium hydroxide at high temperatures.

Sludges

Soft, loose, non-adherent deposits that settle in low-flow areas. Causes: Precipitation of substances like MgCO₃ and CaCO₃.

Limitations of Analytical Water Methods

EDTA Method Problems

1. pH Sensitive: Requires a buffer (pH ≈ 10) for accurate results.
2. Interference: Other metal ions (Fe³⁺, Cu²⁺) may interfere with the titration.
3. Indicator Error: Wrong color change if the indicator is not fresh.
4. Air Contamination: CO₂ from the air affects alkalinity measurements.

Alkalinity Determination Problems

1. Indicator Error: The endpoint may not be sharp.
2. CO₂ Presence: May give false high readings.
3. Mixture of Alkalinities: Difficult to distinguish OH⁻, CO₃²⁻, and HCO₃⁻ easily without stepwise titration.

Zeolite Softening Problems

1. Limited Removal: Only removes Ca²⁺ and Mg²⁺; does not remove Na⁺, Fe³⁺, or CO₂.
2. Regeneration Needed: Requires regular regeneration using NaCl.
3. Efficiency Drops: Reduced efficiency with suspended solids or oils in the water.
4. Sludge Formation: Occurs if hardness is not fully removed.

Electrolysis and Demineralization (DM)

1. Electrolysis (in Water Treatment)

• Definition: Breaking down water into hydrogen and oxygen using electric current.
• Use (Limited): Removes some impurities by oxidation/reduction reactions at the electrodes. Can assist in disinfection or metal ion removal in specialized setups.
• Problems: Not practical for large-scale water softening; high energy cost.

2. Demineralization (DM Plant)

• Definition: Removal of all dissolved salts (ions) using ion-exchange resins.
• Process:
  • Cation Exchanger (H⁺ form): Removes positive ions (Ca²⁺, Mg²⁺, Na⁺).
  • Anion Exchanger (OH⁻ form): Removes negative ions (Cl⁻, SO₄²⁻, NO₃⁻).
  • Result: H⁺ + OH⁻ → H₂O (Pure water)
• Advantages: Produces very high-purity water; removes both hardness and salts.
• Problems: Expensive resins; requires acid/alkali for regeneration; sensitive to turbidity or oil.

Electrodes and Potentiometric Titration Techniques

1. Glass Electrode

   • Use: Measures the pH of a solution.
   • Working: A special glass bulb allows H⁺ ions to interact, creating a voltage related to the pH.
   • Used With: A reference electrode to form a complete pH meter.

2. Ion-Selective Electrode (ISE)

   • Use: Measures specific ion concentration (e.g., Na⁺, K⁺, F⁻).
   • Working: A membrane selectively binds to the target ion, producing a measurable voltage.
   • Types: Glass (for H⁺), crystal, or polymer membrane.

3. Indicator Electrode

   • Function: Participates in redox or ion-specific measurement.
   • Examples: Glass electrode (for pH), platinum electrode (for redox).

4. Reference Electrode

   • Function: Provides a stable, known potential.
   • Common Types: Calomel electrode (Hg/Hg₂Cl₂) or Silver-silver chloride (Ag/AgCl).

5. Conductometric Titration

   • Principle: Measures the change in conductivity as the titrant is added.
   • Graph: X-axis: Volume of titrant; Y-axis: Conductivity. Shows a V-shaped curve; the intersection gives the end point.

6. pH Metric Titration

   • Principle: Measures the change in pH during titration using a pH meter.
   • Graph: X-axis: Volume of titrant; Y-axis: pH. Shows a sigmoidal (S-shaped) curve with a steep jump near the equivalence point.

Conductometry and pH-metry Applications and Definitions

Applications

pH-metry Applications

1. Acid-base titrations (to detect the end point precisely).
2. Monitoring fermentation (pH changes indicate progress).
3. Water quality testing (checks for acidity/alkalinity).
4. Soil pH analysis (for agricultural productivity).
5. Pharmaceuticals (drug formulation and quality control).

Conductometry Applications

1. Titrations (acid-base, precipitation, redox titrations).
2. Purity check (of distilled or deionized water).
3. Salinity measurement (in water samples).
4. Ion concentration detection (in industrial and medical fields).
5. Monitoring reactions where ions are involved.

Definitions

1. Conductance (G)

   The ability of a solution to conduct electricity. Unit: Siemens (S) or mho.

2. Resistance (R)

   Opposition to the flow of electric current. Unit: Ohm (Ω). Depends on the length and area of the conductor and the nature of the medium.

3. Cell Constant (K)

   Ratio of the distance between electrodes (l) to their area (A): K = l/A. Used to calculate specific conductance.

4. Specific Conductance (κ) or Conductivity

   Conductance of 1 cm³ of solution between two electrodes 1 cm apart. Unit: S/cm.

5. Molar Conductance (Λₘ)

   Conductance of all the ions produced by 1 mole of electrolyte in a given volume. Unit: S·cm²/mol.

6. Equivalent Conductance (Λₑₔ)

   Conductance of all the ions produced by 1 gram equivalent of electrolyte.

The Beer-Lambert Law and Mathematical Expression

The Beer-Lambert Law (also known as Beer's Law or the Beer-Lambert-Bouguer Law) relates the absorption of light to the properties of the material through which the light is traveling.

Mathematical Expression

A = ε · c · l

Where:

• A = Absorbance (unitless, as it is a logarithmic ratio)
• ε = Molar absorptivity (or molar extinction coefficient) in L·mol⁻¹·cm⁻¹
• c = Concentration of the absorbing species in solution (mol/L)
• l = Path length of the sample (cm)

Alternate Logarithmic Form

A = log₁₀( Iₒ / I )

Where:

• Iₒ = Intensity of incident light
• I = Intensity of transmitted light

This law is used widely in spectroscopy for quantitative analysis.

Spectroscopic Shifts and Wave Properties

1. Red Shift (Bathochromic Shift)

   Definition: A shift of the absorption or emission spectrum to longer wavelengths (lower energy). Cause: Conjugation, solvent effects, or substituent effects. Example: A compound absorbing at 250 nm shifts to 270 nm.

2. Blue Shift (Hypsochromic Shift)

   Definition: A shift of the absorption or emission spectrum to shorter wavelengths (higher energy). Cause: Removal of conjugation, changes in polarity, or structural changes. Example: A compound absorbing at 270 nm shifts to 250 nm.

3. Hyperchromic Effect

   Definition: An increase in the absorbance (intensity) of absorption or emission. Cause: Often due to structural changes or increased conjugation. Example: DNA denaturation leads to hyperchromicity in UV absorbance.

4. Hypochromic Effect

   Definition: A decrease in the absorbance (intensity) of absorption or emission. Cause: Base stacking in DNA, or reduced conjugation. Example: Double-stranded DNA shows hypochromicity compared to single-stranded.

5. Wavelength (λ)

   Definition: The distance between two consecutive peaks of a wave. Units: Nanometers (nm) or meters (m). Relation to Energy: Inversely proportional to energy.

6. Wave Number (ṷ)

   Definition: The number of waves per centimeter. Formula: ṷ = 1/λ (in cm⁻¹). Units: cm⁻¹ (used often in IR spectroscopy). Higher wave number = higher energy.

7. Frequency (ν)

   Definition: The number of wave cycles per second. Formula: ν = c/λ, where c = speed of light. Units: Hertz (Hz or s⁻¹).