Basics of Analytical Chemistry

Basics of Analytical Chemistry

• Course title: Basics of Analytical Chemistry
• Course codes: chm201, chm219B
• Instructor: Dr Ibrahem M A Hasan

Examples

• Example 1: Calculating volume of HCl needed to neutralize NaOH. 25.0 mL of 0.1 N sodium hydroxide requires 12.5 mL of 0.2 N HCl.
• Example 2: Determining strength of ferrous sulfate solution. A 25 mL solution of ferrous sulfate completely reacts with 30 mL of 0.125 M KMnO4. The solution's concentration is 22.78 g/L.
• Example 3: Calculating volume of reagent needed for dilution. Preparation of 1000 mL of 0.1 N solution from 0.127 N reagent requires 787.4 mL of the latter.

The Equipment

• Volumetric analysis uses specific equipment.
• Pipette: Precisely measures solution volumes.
• Burette: Delivers measured volumes of solutions.
• Conical flask: Mixing solutions.
• Wash bottle: Contains distilled water for cleaning.
• Funnel: Transferring liquids with minimal spillage.
• Volumetric flask: Preparing solutions of known concentration with specific volumes.

Requirements of primary standard substances

• Easy to obtain, purify, dry, and preserve: Preferably at 110-120°C.
• Stable in air: Should not be altered by air, humidity, or carbon dioxide.
• Testable for impurities: Should be capable of being tested for impurities.
• High equivalent weight: To minimize weighing errors.
• Soluble under conditions of use: Facilitates solution preparation.
• Examples: Sodium carbonate, borax, potassium hydrogen phthalate, and sodium chloride.

Preparation of standard solutions

• Weigh a precise amount of pure reagent.
• Dissolve it in the minimum volume of solvent (usually water).
• Transfer the solution to a volumetric flask.
• Fill the flask to the mark with solvent and mix thoroughly.
• 1.0 M NaCl solution example: Weigh 58.45 grams of NaCl and add it to 1.0 L volumetric flask, then dissolve with water to the mark.

Theory of acid-base titration

• 1. Arrhenius theory:
  • Acid: Solution with excess H+ ions.
  • Base: Solution with excess OH- ions. H+ + OH- → H₂O.
• 2. Brønsted-Lowry theory:
  • Acid: Proton (H+) donor.
  • Base: Proton (H+) acceptor. HCl(g) + H₂O(l) → H₃O+(aq) + Cl-(aq).
• 3. Lewis theory:
  • Acid: Electron-pair acceptor.
  • Base: Electron-pair donor.

Conjugate acid-base pairs

• Conjugate acid-base pairs differ by one proton.
• When an acid donates a proton, it forms the conjugate base.
• HCI(g) → Cl-(aq) (acid → conjugate base).
• When a base accepts a proton, it forms the conjugate acid. H₂O(l) → H₃O+(aq) (base → conjugate acid).
• Acid + Base → Conjugate base + Conjugate acid.

Relative Strengths of Conjugate Acid-Base Pairs

• List of acids and their conjugate bases, ordered by relative strengths.

Equivalent weight

• Acid: The weight containing one replaceable hydrogen (MW/basicity). -Monobasic acid have Equivalent weight and molecular weight same.
• Base: The weight containing one replaceable hydroxyl group (MW/acidity).
• Compound: Molecular weight/number of electrons charges.
• A normal solution of monobasic acid contains 1 MW per Liter.
• example calculations for dibasic and tribasic acids.

Hydrogen-ion concentration and the pH scale

• Self-ionization of water: H₂O + H₂O ↔ H₃O⁺ + OH⁻.
• Equilibrium constant (Kw): Kw = [H₃O⁺][OH⁻] = 1.0 x 10⁻¹⁴

pH and pOH

• pH: Negative logarithm of the hydronium ion concentration
• pOH: Negative logarithm of the hydroxide ion concentration.
• Relationships among hydrogen and hydroxide concentration and pH.

Examples (Calculations)

• Example 1: Calculate pH from given hydronium ion concentration.
• Example 2: Calculate hydronium ion concentration from pH.
• Example 3: Calculate pH of a given acetic acid solution (including dissociation).

Calculating [H⁺] or pH from ionization constants

• Henderson-Hasselbalch equation: For weak acids, pH = pKa + log([salt]/[acid]).
• Titration: Finding [H+] (e.g. in an acid-base titration.