Water 2

Relevant Compounds:
  - Water (H2O, MW 18 g/mol)
  - Ammonia (NH3, MW 17 g/mol)
  - Methane (CH4, MW 16 g/mol)
Structural Characteristics: All these molecules belong to the tetrahedral family but exhibit different thermal properties due to their bonding and molecular configuration.
Heat of Fusion:
  - Water: 6.01 kJ/mol
  - Ammonia: 5.66 kJ/mol
  - Methane: 0.94 kJ/mol
Analysis:
- The heats of fusion decrease significantly from water to methane.
- This trend indicates the strength of intermolecular forces. Draw the structural formula to visualize differences.
Hydrogen Bonding:
  - Water forms extensive hydrogen bonds due to two unshared electron pairs on oxygen, contributing to its unique properties in solid state.
  - Ammonia also forms hydrogen bonds but with a single unshared electron pair on nitrogen. Methane lacks hydrogen bonding capability entirely.
  - If NH3 were to form ice (melting point -97.8 °C), it would be expected to be less dense than liquid ammonia due to the structural framework of hydrogen bonds in the solid-state facilitating a more open lattice structure.

General Principle: Water can ionize, resulting in the formation of hydrogen ions (H+) and hydroxide ions (OH-).
  - The process can be represented as:
    - $H2O <br>ightleftharpoons H^+ + OH^-$
  - This is a reversible reaction which defines the water equilibrium constant.
Hydronium Ion Formation:
- In an aqueous solution, a proton can combine with a water molecule to form a hydronium ion:
- $H_2O + H^+ <br>ightarrow H_3O^+$
Ion Product of Water (Kw):
  - The ion product is defined as:
    - $K_w = [H^+][OH^-]$
  - At 25°C and 1 atm pressure, $K_w = 1.0 imes 10^{-14}$

Definitions:
  - Acid: A proton donor
  - Base: A proton acceptor
  - Many organic molecules function as weak acids or weak bases.
Conjugate Bases:
- The deprotonated product of an acid-base dissociation reaction is termed a conjugate base.
- Many biomolecules display acidic and/or basic characteristics.

Weak acids and bases do not fully dissociate in solution.
Dissociation Constant for Weak Acid (HA):
  - $HA <br>ightleftharpoons H^+ + A^-$
  - Dissociation constant (Ka) defined as:
    - $K_a = rac{[H^+][A^-]}{[HA]}$
pKa Definition:
  - The pKa is the negative logarithm of the acid dissociation constant:
    - $pK_a = - ext{log}(K_a)$
  - A lower pKa indicates a stronger acid.
For bases, a similar expression can be used:
  - $K_b = rac{[BH^+][OH^-]}{[B]}$
  - Similarly, for base dissociation:
    - $pK_b = - ext{log}(K_b)$
    - The relationship:
      - $K_a imes K_b = K_w$
  - An increase in K means a decrease in pK, indicating stronger acids or bases.

pH Definition:
  - The pH scale measures hydrogen ion concentration:
    - $pH = - ext{log}[H^+]$
    - Alternatively,
    - $pH = - ext{log}[H_3O^+]$
pKa Relationship: pKa is a measure of a weak acid's strength and calculated as
- $pK_a = - ext{log}(K_a)$

Example Table:

Acid	Anion	Ka	pKa	Strength
Acetic Acid	CH3COO	1.76 x 10^-5	4.76	Monobasic
Carbonic Acid	H2CO3	4.5 x 10^-7	6.35	Monobasic
Bicarbonate	HCO3^-	5.61 x 10^-11	10.33	Weakest
Lactic Acid	CH3CHCOOH	1.38 x 10^-4	3.862	Weak acid
Phosphoric Acid	H3PO4	7.25 x 10^-3	2.14	Strongest

Activity vs Concentration: Equilibrium constants should be determined with activities wherever possible, but concentrations might be substituted for reasonable accuracy in dilute solutions.

Importance:
- Regulation of pH is essential in biological systems (e.g., Le Chatelier’s principle).
- Changes in pH can have pathological consequences (e.g., acidosis or alkalosis).
Buffer Components: Buffers typically consist of a weak acid and its conjugate base.
Equilibrium:
- Buffers work by establishing equilibrium between their components and follow Le Chatelier’s principle — equilibrium shifts in a direction that helps relieve stress.

Definition: The ability of a buffer to maintain a specific pH is based on:
1. The molar concentration of the acid-conjugate base pair.
2. The ratio of their concentrations.
Concept: Buffers have a buffering capacity that is directly proportional to their concentration.

Equation:
- The equation defines the relationship between pH, pKa, and concentrations of conjugate base ([A-]) and undissociated weak acid ([HA]):
- $pH = pK_a + ext{log} rac{[A^-]}{[HA]}$
Optimal Buffering: Buffers are most effective when composed of equal parts weak acid and conjugate base, providing best buffering usually within one pH unit above and below pKa.

Example: Bicarbonate Buffer is crucial in blood plasma for maintaining pH:
  - The reaction:
    - $CO2 + H2O <br>ightleftharpoons H^+ + HCO3^-$ (HCO3^- is bicarbonate).
    - This reaction is reversible and managed by the enzyme carbonic anhydrase.

Calculate the pH of a mixture of 0.25 M acetic acid and 0.20 M sodium acetate. The pKa of acetic acid is 4.76.
Calculate the ratio of lactic acid and lactate required in a buffer system at pH 4.9. The pKa of lactic acid is 3.86.
Assignments: Complete the problems and upload the document on Moodle.