Analytical Chemistry Techniques, Reagent Preparation, and Acid-Base Theory

Fundamentals of Analytical Chemistry Techniques

Analytical techniques are methods used to identify, quantify, and understand the chemical composition and structure of substances. These techniques are broadly classified into qualitative (what is present) and quantitative (how much is present) methods.

Major Analytical Methods

Common analytical techniques include:

• Gravimetric Analysis: Involves measuring the mass of a substance to determine the amount of analyte.
• Titrimetric (Volumetric) Analysis: Based on measuring the volume of a standard solution required to react with the analyte.
• Spectroscopic Methods: Measure the interaction between light and matter.
• Electrochemical Methods: Based on the measurement of electrical properties.

Preparation and Standardization of Reagents

Hydrochloric Acid (HCl) Preparation

• Molarity (M) = 1 mol/L
• Molecular weight of HCl = 36.5 g/mol
• Concentrated HCl is approximately 37% w/w and has a specific gravity = 1.18 g/mL.
• Approximate molarity of concentrated HCl ≈ 11.8 M.

To prepare 1 L of 1 M HCl:

Sulfuric Acid (H₂SO₄) Preparation

Preparation of 1 N H₂SO₄

• Equivalent weight of H₂SO₄ = 49 g/equiv (as it yields 2 H⁺ ions).
• Concentrated H₂SO₄ is approximately 98% w/w and has a specific gravity = 1.84 g/mL.
• Approximate normality of concentrated H₂SO₄ ≈ 36 N.

To prepare 1 L of 1 N H₂SO₄:

Procedure:

Measure 27.8 mL of concentrated H₂SO₄, slowly add it to water, and dilute the solution to 1 liter.

Standardization of 1 N H₂SO₄

Use standard sodium carbonate (Na₂CO₃) for standardization.

Reaction:

H₂SO₄ + Na₂CO₃ → Na₂SO₄ + CO₂ + H₂O

Procedure:

Titrate 25.0 mL of H₂SO₄ with standard 0.1 N Na₂CO₃ using methyl orange indicator.

Calculate the exact normality using the formula:

N₁V₁ = N₂V₂

Principles of Acid-Base Titration

Acid-base titration is a quantitative analytical method used to determine the concentration of an unknown acid or base by reacting it with a base or acid of known concentration.

It involves the gradual addition of one solution (the titrant) from a burette into another solution (the analyte) until the reaction reaches the equivalence point.

An indicator (e.g., phenolphthalein or methyl orange) is often used to visually show the end point of the titration by a color change.

Key Acid-Base Theories

1. Arrhenius Theory (Svante Arrhenius)

   • Acid: Produces H⁺ ions in aqueous solution.
   • Base: Produces OH^- ions in aqueous solution.
   • Limitation: Only applies to aqueous solutions; does not explain acid-base behavior in non-aqueous systems.

   Example:

   HCl → H⁺ + Cl^-

   NaOH → Na⁺ + OH^-

2. Brønsted-Lowry Theory (1923)

   • Acid: Proton (H⁺) donor.
   • Base: Proton (H⁺) acceptor.
   • Application: Applies to both aqueous and non-aqueous systems.

   Example:

   NH₃ + H₂O ↔ NH₄⁺ + OH^-

   (NH₃ acts as a base by accepting H⁺)

3. Lewis Theory (Gilbert Lewis)

   • Acid: Electron pair acceptor.
   • Base: Electron pair donor.
   • Scope: Most general theory; includes reactions that do not involve H⁺ ions.

   Example:

   BF₃ + NH₃ → F₃B←NH₃

   (BF₃ is the Lewis acid; NH₃ is the Lewis base)