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Lecture Notes: Water, pH, Buffers, and Amino Acids

Water and hydrogen bonding in biological contexts

  • Water is a molecule with hydrogen-bond interactions between water molecules that continuously form and break. These interactions underlie solvation, structure, and many biochemical processes.
  • For most biochemical purposes, the composition of water is treated as static (a constant solvent background) because the autoionization of water is at very low absolute concentrations compared to solutes.

Ion product of water and pH scale fundamentals

  • Ion product of water: K_w = [H^+][OH^-] = 1 \times 10^{-14}.
  • This implies that the product of the hydrogen ion concentration and hydroxide ion concentration is fixed at this value (at 25°C).
  • pH concepts:
    • \mathrm{pH} = -\log{10}[H^+],\quad \mathrm{pOH} = -\log{10}[OH^-]
    • At 25°C, \mathrm{pH} + \mathrm{pOH} = 14.
  • The pH scale ranges roughly from 0 to 14 in aqueous solutions, with pH 7 being neutral, lower values acidic, and higher values basic.
  • In biology, water is abundant, and the small autoionization ions are usually negligible relative to solute concentrations, which is part of why water’s role is treated as a constant background in many calculations.

pH homeostasis, buffers, and why pH matters in biochemistry

  • pH homeostasis is crucial: small changes in pH can disrupt enzyme activity, protein structure, and metabolic fluxes.
  • Buffers are solutions that resist changes in pH by shifting between acid and base forms.
  • General buffer concept via Henderson–Hasselbalch equation: \mathrm{pH} = pKa + \log{10}\left(\frac{[\text{A}^-]}{[\text{HA}] }\right) where HA is the weak acid and A⁻ is its conjugate base.
  • Example buffer pair in amino/biochemical contexts: the ammonium/ammonia system with NH4^+ \rightleftharpoons NH3 + H^+ and a given pK_a.
  • The pKa given in the transcript for the ammonium system is pKa = 9.6. This corresponds to the equilibrium constant for: NH4^+ \rightleftharpoons NH3 + H^+. Ka = \frac{[NH3][H^+]}{[NH_4^+]} = 10^{-9.6}.
  • Relationship of pH to the ammonium system:
    • If \mathrm{pH} < pKa\;(9.6), the protonated form NH4^+ dominates.
    • If \mathrm{pH} > pKa\,, the base form NH3 dominates.
  • At physiological pH around 7, the ratio of base to acid is: \frac{[NH3]}{[NH4^+]} = 10^{\mathrm{pH} - pKa} = 10^{7 - 9.6} \approx 10^{-2.6} \approx 2.5 \times 10^{-3}. Hence, NH4^+ dominates by a large margin (roughly 99.75% in the protonated form) at pH ~7.
  • Visualizing on a titration curve: around the pKa, buffering is most effective; below the pKa the acid form dominates; above the pK_a the conjugate base dominates.
  • In the transcript, there is mention of a question about which species dominates at a given pH relative to pK_a and how this relates to charge states and buffering.

Condensation (dehydration) reactions and the role of water

  • Condensation reactions form covalent bonds with the loss of water: general form
    • A{-}OH + HO{-}B \rightarrow AB + H_2O.
  • In aqueous environments, the reverse hydrolysis reaction is also possible: water participates to break a bond and form two molecules.
  • Equilibrium and Le Châtelier’s principle: when water is abundant, condensation reactions are pushed toward the reactants (left) because the product side includes water; increasing water concentration tends to shift the equilibrium to favor hydrolysis (the reverse process) rather than condensation.
  • The transcript notes that in proteins, amino acids can undergo covalent modifications after translation (post-translational modifications, PTMs). This expands the number of possible chemical states significantly.

Post-translational modifications (PTMs) and amino acid chemistry

  • PTMs refer to covalent modifications of amino acid residues after translation, increasing the diversity of protein function.
  • The transcript notes that the number of potential modifications increases dramatically when considering all possible PTMs.
  • A basic biochemical way to think about amino acids in solution is their zwitterionic form, especially around physiological pH:
    • The typical zwitterion form is: NH_3^+{-}CH(R){-}COO^-, with the amino group positively charged and the carboxyl group negatively charged.
  • The balance between the positively charged amino group and the negatively charged carboxyl group contributes to the overall (often near-neutral) charge of amino acids in solution at physiological pH.
  • The discussion in the transcript hints at acid–base interactions involving amino groups and protonation states, including the notion of an NH_3^+ group interacting with H^+ on the left side of a reaction and the right side involving a nucleophile in some contexts.
  • In many contexts, the protonated amino group (NH_3^+) and the deprotonated carboxylate (COO^−) coexist, giving rise to the characteristic zwitterionic form that dominates at physiological pH.

Connections to foundational principles and real-world relevance

  • Le Châtelier’s principle governs how reactions shift when water or pH is perturbed (e.g., condensation reactions vs hydrolysis and buffer systems).
  • The ion-product of water (Kw) ties directly into the pH and pOH scales and underpins acid–base chemistry in all aqueous systems.
  • Buffers play a critical role in maintaining stable intracellular and extracellular pH, enabling enzymes to function optimally and preventing denaturation.
  • PTMs are central to cellular regulation, signaling, and diverse protein functions; their combinatorial possibilities contribute to proteome complexity.
  • The zwitterionic nature of amino acids explains their behavior in solution, protein folding, ionizable side chains, and interactions with other molecules.

Practical numerical references and constants

  • Water autoprotolysis constant: K_w = [H^+][OH^-] = 1\times 10^{-14}.
  • pH and pOH relation: \mathrm{pH} + \mathrm{pOH} = 14\quad (\text{at } 25^{\circ}\mathrm{C}).
  • Henderson–Hasselbalch relation: \mathrm{pH} = pKa + \log{10}\left(\frac{[\text{base}]}{[\text{acid}] }\right).
  • Ammonium–ammonia system (acid dissociation of ammonium): NH4^+ \rightleftharpoons NH3 + H^+\,,\; pK_a = 9.6\,.
  • Acid dissociation constant for the ammonium reaction: Ka = \frac{[NH3][H^+]}{[NH_4^+]} = 10^{-9.6}.
  • Ammonia/ammonium speciation at physiological pH: at \mathrm{pH} = 7, \frac{[NH3]}{[NH4^+]} = 10^{7-9.6} \approx 2.5\times 10^{-3}.
  • Amino acid zwitterion example: NH_3^+{-}CH(R){-}COO^-.
  • Approximate water concentration as solvent: [H_2O] \approx 55.5\;\mathrm{M} (often treated as constant in biochemistry).

Summary takeaways

  • Water’s hydrogen-bond network underpins solvation and biochemical interactions, but in practice its ionization is treated as a constant background in many cellular contexts.
  • The ion product of water and the pH/pOH scales are central to understanding acid–base chemistry in biology; small pH changes can have large downstream effects on structure and function.
  • Buffers enable pH stability around physiologically relevant ranges; the pKa of the relevant acid–base pair dictates where buffering is most effective and which species dominate at a given pH.
  • Condensation reactions release water; in aqueous environments, water acts to push the equilibrium toward hydrolysis, illustrating Le Châtelier’s principle.
  • PTMs expand the functional diversity of proteins, and amino acids exist predominantly as zwitterions at physiological pH, balancing charges between NH_3^+ and COO^− groups.
  • Understanding these concepts helps explain enzyme activity, protein folding, signaling, and the design of biochemical experiments where pH, buffers, and solvent conditions are critical.