general chemistry -3

pH and Ionization of Water

• Pure water can ionize weakly, producing hydrogen ions (H+) and hydroxide ions (OH−) as shown in the equation:[ H_2O \rightleftharpoons H^+ + OH^- ]

• However, free hydrogen ions in water are not present; instead, they are hydrated and called hydronium ions (H₃O). The dissociation of water may be rewritten as:[ H_2O + H_2O \rightleftharpoons H_3O^+ + OH^- ]

• The equilibrium constant (K_eq) for this reaction is given by:[ K_{eq} = \frac{[H^+][OH^-]}{[H_2O]} ]

Equilibrium Constant and Ion Product of Water

• The dissociation of water is minimal, allowing us to assume the concentration of undissociated water is constant (approximately 55.5 M). Therefore, the equation simplifies to:[ Kw = [H^+][OH^-] = (55.5 M)K_{eq} ]

• Here, K_w represents the ion product of water, which at 25°C equals 1 × 10^-14 M², showing that the concentrations of H⁺ and OH− are each 1 × 10⁻⁷ M. At equal concentrations, an aqueous solution is defined to be neutral pH.

The pH Scale

• pH is derived from the ion product of water:[ [H^+][OH^-] = 10^{−14} ]

• Taking logarithms results in:[ -\log[H^+] - \log[OH^-] = 14 ]

• Therefore, we define:

• [ pH = -\log[H^+] ]

• [ pOH = -\log[OH^-] ]

• This leads to the relationship:[ pH + pOH = 14 ]

Characteristics of the pH Scale

• The pH scale ranges from 0 (most acidic) to 14 (most basic).

• A pH of 7 is neutral; values below 7 signify acidity, while values above 7 indicate basicity.

• pH plays a crucial role for living organisms as it influences biomolecule function, where most biomolecules are active within a narrow pH range.

Acids and Bases: Arrhenius Concept

• According to the Arrhenius definition:

  • Acids release H+ ions in water.

  • Bases release OH− ions in water.

• This definition is limited to protic acids and hydroxide bases.

Acids and Bases: Bronsted-Lowry Concept

• Bronsted-Lowry defines acids as proton donors and bases as proton acceptors.

• This introduces the concept of conjugate acids and bases, where:

  • Donating a proton converts an acid to its conjugate base.

  • The weaker the acid, the stronger its conjugate base.

Acids and Bases: Lewis Concept

• Lewis definitions describe acids as electron pair acceptors and bases as electron pair donors.

• For example, BF₃ acts as a Lewis acid, while tertiary amines are Lewis bases.

Strength of Acids and pKa

• The strength of an acid is determined by its dissociation in water, expressed by pKa.

• Lower pKa values indicate stronger acids, while higher values indicate stronger bases.

• The relationship between acid dissociation constant (K_a) and pKa is given as:[ pKa = -\log Ka ] [ K_w = K_a \times K_b ]

Measurement of pH

• pH measures H+ concentration in a solution:

  • High H+ concentration = acidic solution (low pH).

  • Neutral solution has minimal H+ concentration (pH around 7).

  • Basic solution has very low H+ concentration (high pH).

Buffer Solutions

• A buffer solution consists of a weak acid and its conjugate base, capable of resisting pH changes upon the addition of acids or bases.

• Example: Sodium acetate in acetic acid serves as a buffer, crucial for maintaining blood pH near neutrality.

Types of Buffer Solutions

• Acidic Buffers: Mix of a weak acid with its salt, maintaining an acidic environment (e.g., acetic acid and sodium acetate, pH 4.74).

• Alkaline Buffers: Mix of weak bases with their salts, preserving basic conditions (e.g., ammonia and ammonium chloride, pH 9.25).

Titration of Weak Acids

• Adding NaOH to a weak acid neutralizes H+ ions, shifting the dissociation equilibrium and increasing pH.

• The weak acid resists pH change near its pKa, allowing for controlled pH adjustments.

Mechanism of Buffer Action

• Upon acid addition, protons are consumed by acetate ions to form acetic acid.

• Base addition sees hydroxide ions react with hydrogen ions to form water, maintaining pH stability.

Henderson-Hasselbalch Equation

• This equation relates pH, pKa, and concentrations of acidic and basic forms:[ pH = pKa + \log \left( \frac{[A^-]}{[HA]} \right) ]

• It facilitates pH calculations for buffer solutions but cannot be applied to strong acids or bases.

Practical Applications and Limitations

• The Henderson equation helps in:

  • Calculating buffer pH.

  • Finding pKa values.

  • Preparing buffer solutions of desired pH.

• Limitations include its inapplicability to strong acids and bases.

Numerical Problems in Buffer Solutions

• Problem-solving scenarios are provided to create buffers of specific pH values using known concentrations of acids and bases, demonstrating practical application of theory to laboratory situations.

VSEPR Theory Introduction

• VSEPR (Valence Shell Electron Pair Repulsion) theory, established by Nyholm and Gillespie, predicts molecular geometry based on electron pair repulsions.

• The arrangement minimizes electron pair repulsion, determining the shape of molecules.

Molecular Geometry Based on Electron Pairs

• The geometry of molecules is derived from the number of electron pairs:

  • 2 pairs: Linear

  • 3 pairs: Trigonal planar or bent

  • 4 pairs: Tetrahedral (with lone pairs affecting the shape)

  • 5 pairs: Trigonal bipyramidal or distorted shapes, etc.

Postulates of VSEPR Theory

• Identifies a central atom bonded to surrounding atoms, with electron pairs orienting to minimize repulsion.

• Lone pairs cause distortions in molecular shapes, influencing geometry.

Example: Ammonia and Water

• Ammonia exhibits trigonal pyramidal geometry due to one lone pair, while water has a bent shape owing to its two lone pairs.

Limitations of VSEPR Theory

• Cannot account for isoelectronic species with the same electron number yet different shapes.

• Unexplained geometries of certain transition metal compounds.

Crystal Field Theory (CFT)

• Developed by H. Bethe and Van Vleck, CFT considers ligands as point charges leading to d-orbital splitting in metal complexes.

• Characteristics of splitting include variations in energy levels in octahedral and tetrahedral complexes.

Summary of Octahedral and Tetrahedral Splitting

• IN octahedral complexes, the eg orbitals have higher energy than t2g orbitals. Complexes can be classified into high or low spin depending on the strength of ligands and energy considerations.

Square Planar Complexes

• In square planar complexes, four ligands interact primarily in the xy plane, resulting in higher splitting energy and often low spin configurations.

Conclusion

• Buffer solutions, pH measurements, and molecular geometries are essential concepts in chemistry that illustrate the interplay of acid-base reactions, molecular structure, and molecular behavior under various conditions.