Biochemistry Review Notes: Thermodynamics, Water Interactions, Hydrophobic Effect, and Acids/Bases

GSS and Course Logistics

  • There is GSS support for this class, run by Sal, who has survived the course and performed well.
  • GSS is an active-learning model with small group work to review concepts and practice problems in class.
  • Schedule: four to five students per group on Tuesdays and Thursdays in the Bio Psych Building, Room 1228.
  • Details will be sent out after class; ask questions or discuss homework as needed.
  • Sal’s tone emphasizes engagement and that this course is a good bridge to larger biochemistry topics.

Quick Recap: Thermodynamics in Biochemistry

  • Prior lecture framing: thermodynamics viewed via enthalpy (ΔH) and entropy (ΔS), and also via concentrations and distance from equilibrium.
  • Biochemists’ perspective on energy changes can differ from a purely physical view, but the principles are the same.
  • Non-spontaneous reactions in biology are common; cells use strategies to make such reactions proceed.
  • Key strategies to drive non-spontaneous reactions forward:
    • Couple with highly favorable processes, especially ATP hydrolysis (the energy release helps drive otherwise unfavorable steps).
    • Spatial and network level control: remove products or isolate them to shift the reaction balance and drive the forward direction.
  • These concepts set up later metabolism discussions (ATP homeostasis) which will be revisited in metabolism modules.

Driving Non-Spontaneous Reactions in Biology

  • ATP hydrolysis is a central driving force due to its large energy release.
  • Other reactions can also provide the favorable energetic push, not only ATP hydrolysis.
  • Mechanistic idea: remove products to alter the product/reactant ratio and push the reaction forward despite unfavorable intrinsic spontaneity.
  • Concept of networks: pathways where the product of one reaction becomes the reactant for the next allow sequential driving of otherwise unfavorable steps.
  • These principles are foundational and will be connected to exam content (often more evident after later topics).

Learning Goals for This Lecture

  • Appreciate the biochemical significance of different chemical interactions.
  • Understand why water is remarkable in biochemistry.
  • Re-familiarize with acids, bases, buffers, and Henderson–Hasselbalch concepts.
  • Prepare for deeper biochemistry by reinforcing core principles.

Chemical Interactions and Bonding: Overview

  • Common chemical interactions (not exhaustive): covalent bonds, ionic interactions, dipole–dipole interactions, hydrogen bonds, and Van der Waals interactions.
  • Covalent bonds: sharing of at least one electron pair; each atom achieves (or approaches) its octet; electrons are technically shared, not owned by either atom.
  • Ionic interactions: not literally about permanent electron transfer in the bond formation context, but involve electrostatic attraction between charged species (often after electron transfer occurs in the reaction history).
  • Dipole–dipole interactions: partial charges created by unequal sharing (e.g., carbonyl groups where oxygen is more electronegative and draws electron density toward itself).
  • Hydrogen bonds: hydrogen-bond donor (H attached to an electronegative atom) and hydrogen-bond acceptor (electronegative atom with a lone pair); hydrogen bonds are covalent-ish, not true covalent bonds.
  • Hydrogen-bond donors/acceptors are defined by where the hydrogen comes from and where the lone pair resides, respectively.
  • Van der Waals interactions: weak, non-specific attractions that become significant when many contacts occur across a surface; can be overridden by strong repulsion when atoms get too close.
  • The relative strength of these interactions shapes solubility, structure, and macromolecular assembly in biology.

Hydrogen Bonds: Donor and Acceptor Roles

  • For a hydrogen bond to form, two electronegative atoms (typically O or N in biochemistry) must be involved:
    • One atom bears a hydrogen (donor).
    • The other atom provides a lone pair (acceptor).
  • The hydrogen atom is not transferred; the bond is a noncovalent interaction where the donor gives the hydrogen for stabilization with the acceptor.
  • This concept links acid–base chemistry to hydrogen bonding (proton management is a central theme in biochemistry).
  • Hydrogen bonds are pervasive in water, proteins, nucleic acids, and many biomolecular interfaces.

Van der Waals Interactions: Distance and Energy Landscape

  • Van der Waals (vdW) interactions are weak on a per-contact basis but numerous across molecular surfaces, making them collectively significant.
  • The concept of a Van der Waals radius: an effective “sphere” around an atom where non-bonded interactions occur.
  • If two atoms approach too closely, strong repulsion occurs, pushing them apart.
  • Graphically: there is a shallow energy minimum as two atoms come into vdW contact, followed by a sharp rise in energy as they overlap beyond the favorable distance.
  • This framework helps analyze covalent bonds, hydrogen bonds, and vdW contacts within molecular assemblies.

Water: A Key Biochemical Player

  • Water has a strong permanent dipole (partial negative on the oxygen, partial positive on the hydrogens) and can form both dipole–dipole interactions and hydrogen bonds.
  • Water molecules can act as both hydrogen-bond donors and acceptors (each water molecule can donate twice and accept twice due to two hydrogens and two lone pairs).
  • Hydrophilic vs hydrophobic distinctions:
    • Hydrophilic substances interact favorably with water via dipole interactions and hydrogen bonding.
    • Hydrophobic substances interact poorly with water and are stabilized by vdW interactions; they tend to cluster in water to minimize disruption of the hydrogen-bond network.
    • Amphiphilic (amphipathic) molecules have both hydrophilic and hydrophobic regions.
  • The hydrophobic effect: a major driving force in biochemistry.
    • When hydrophobic molecules are in water, water forms highly structured hydration shells around them, which is entropically unfavorable (negative ΔS) for water but can be offset by favorable interactions among hydrophobic molecules (primarily vdW) that cluster together.
    • The net result is often a negative ΔG, favoring aggregation of hydrophobic parts, which drives protein folding, membrane formation, and biomolecular assembly.
    • This hydrophobic-driven organization reduces the exposed hydrophobic surface area to water, thereby reducing the disruption to water’s hydrogen-bond network overall.
  • Visual metaphors used in class emphasize how water “interacts with the party” of hydrophobic molecules and how clustering around the hydrophobic cores can be enthalpically favorable and entropically unfavorable for water, but the net effect promotes assembly.

Hydrophobic Effect: Implications for Biochemistry

  • Drives protein folding: hydrophobic residues tend to bury themselves inside the protein, away from water.
  • Drives membrane formation and cellular compartmentalization: lipid bilayers sequester hydrophobic tails away from water.
  • Critical for biomolecular interactions: hydrophobic interfaces influence how proteins and ligands interact.
  • Overall message: the hydrophobic effect is arguably the main driving force behind many biochemical processes and structures.

Hydrophiles, Hydrophobes, and Amphipathic Molecules: Summary

  • Hydrophilic: interacts well with water; often charged or highly polar; includes many polar functional groups and hydrogen-bond-capable moieties.
  • Hydrophobic: lacks favorable interactions with water; tends to be nonpolar and noncharged; primarily stabilized by van der Waals interactions with other hydrophobic groups.
  • Amphipathic (amphiphilic): molecules with both hydrophilic and hydrophobic regions, enabling the formation of micelles, bilayers, and other assemblies.
  • These properties influence solubility, localization, and interactions in biological systems.

Acids, Bases, and Buffers: Bronsted–Lowry View (and Beyond)

  • Bronsted–Lowry definitions (emphasized in this course):
    • Acid: a substance that donates a proton (H+).
    • Base: a substance that accepts a proton (H+).
  • In aqueous solutions, water participates in acid–base chemistry by acting as a base (accepting H+) and forming hydronium (H3O+); the conjugate acid–base pairs are interconvertible.
  • Notion of conjugate pairs:
    • An acid conjugate base is the species left after donating a proton.
    • A base conjugate acid is the species formed after accepting a proton.
  • The strength of acids is described by the equilibrium constant Ka for the acid dissociation in water.
  • Water’s role in acid–base reactions: large excess of water means its activity is essentially 1, so water is not included in Ka expressions for many weak acids.

Key Equations: Acids, Bases, and Buffers

  • Acid dissociation constant (Ka) for a weak acid HA donating a proton to water:
    Ka=[H+][A−][HA]K_a = \frac{[H^+][A^-]}{[HA]}
  • pKa as a measure of an acid’s propensity to donate a proton:
    extpK<em>a=−log⁡</em>10(Ka)ext{p}K<em>a = -\log</em>{10}(K_a)
  • Henderson–Hasselbalch equation (linking pH, pKa, and concentrations of conjugate base/acid):
    extpH=extpK<em>a+log⁡</em>10frac[A−][HA]ext{pH} = ext{p}K<em>a + \log</em>{10} \\frac{[A^-]}{[HA]}
  • Important interpretation for pH vs pKa:
    • pH is the actual proton activity in solution (measure of hydronium ion concentration).
    • pKa is a property of the acid that indicates how strongly it tends to donate a proton (the group’s intrinsic desire to be protonated/deprotonated).
    • They are not the same number; pKa is a property of the acid, pH is the condition of the solution.
  • Practical rule of thumb for pH relative to pKa (as explained in lecture):
    • When pH < pKa, the protonated form predominates; when pH > pKa, the deprotonated form predominates.
    • The buffering region is within about ±1 pH unit of the pKa where buffering capacity is highest.
  • Example: acetic acid (HOAc) with pKa ≈ 4.8:
    • At pH = 4.8, about 50% of HOAc is protonated (HA) and 50% is deprotonated (A−).
    • The buffering range for HOAc/AcO− is approximately from pH ≈ 3.8 to pH ≈ 5.8 (±1 pH unit from pKa).
  • Polyprotic acids:
    • Have more than one dissociable proton and thus more than one pKa value.
    • Each pKa is treated separately in Henderson–Hasselbalch analyses.
    • At physiological pH (~7), most acids are already deprotonated; common acid–base behavior in biochemistry often involves conjugate acids/bases rather than the fully protonated/neutral forms.
  • Practical biochemistry framing: biochemistry often concerns reaction steps where species are already in the conjugate forms present at pH ~7; thus, the roles of acids and bases are viewed in terms of their ability to donate or accept protons in the context of the reaction, rather than as isolated acid/base species.

Connecting pH, pKa, and Biological Context

  • The vast majority of biochemical reactions occur around physiological pH (~7).
  • In this context, acids are generally already deprotonated (conjugate bases) and bases are generally already protonated (conjugate acids) relative to the formal definitions, but actual behavior depends on the specific pKa values of the functional groups involved.
  • The language often used: a species can act as an acid or a base depending on the reaction context (act as an acid vs act as a base), rather than being inherently one or the other in every situation.
  • The example of amines and carboxylates:
    • A protonated amine (NH3+) can donate a proton in a reaction (acting as an acid) if it loses the proton.
    • A carboxylate (COO−) can accept a proton (acting as a base) to become carboxylic acid (COOH) in some contexts.
  • This contextual “act as an acid” or “act as a base” framing is emphasized to align with real biochemistry at neutral pH.

Practical Takeaways and Study Cues

  • Understand the roles of ΔG°, ΔG, K_eq, and Q in determining spontaneity and how biology manipulates them:
    • \Delta G = \Delta G^b0 + RT \ln Q
    • ΔG=−RTln⁡Keq+RTln⁡Q\Delta G = -RT \ln K_{eq} + RT \ln Q
    • If Q < Keq, the reaction tends to proceed forward; if Q > Keq, it tends to proceed backward.
  • ATP hydrolysis is a major cellular driver of otherwise unfavorable reactions due to its large negative free energy change.
  • Cellular strategies to push unfavorable reactions forward include product removal and pathway organization to couple reactions effectively.
  • Hydrophobic interactions, water structure, and the hydrophobic effect are central to macromolecular folding and membrane assembly.
  • Distinctions between hydrophilic, hydrophobic, and amphipathic molecules explain solubility, localization, and assembly in cells.
  • Hydrogen bonds and van der Waals interactions play complementary roles in stabilizing biomolecules and directing interactions.
  • Buffering capacity is highest within ±1 pH unit of a given pKa; this is a key concept for maintaining cellular pH homeostasis and designing experiments.
  • Polyprotic acids require careful, stepwise analysis of each proton dissociation event; no single pKa explains all protonation events.

Quick Numeric References (from lecture examples)

  • Covalent O–H bond distance (approximate): rO−H≈0.96 A˚r_{O-H} \approx 0.96\,\text{Å}
  • Van der Waals radii (approximate):
    • Oxygen: rvdW,O≈1.4 A˚r_{vdW,O} \approx 1.4\,\text{Å}
    • Hydrogen: rvdW,H≈1.2 A˚r_{vdW,H} \approx 1.2\,\text{Å}
  • Hydrogen bond distance (typical): dHB≈1.8 A˚d_{HB} \approx 1.8\,\text{Å}
  • Water’s dipole and hydrogen bonding capacity enable multiple donor/acceptor interactions in biological settings.
  • Acetic acid example:
    • pKa ≈ 4.84.8;
    • When pH = pKa, about half of the molecules are protonated (HA) and half are deprotonated (A−).

Final Thoughts

  • The lecture ties together how physical chemistry underpins biological phenomena, focusing on water as a key solvent, the spectrum of chemical interactions, and the central role of pH, pKa, and buffering in biological systems.
  • The hydrophobic effect stands out as a unifying driving force behind protein folding, membrane formation, and macromolecular association.
  • Mastery of these concepts lays the groundwork for later metabolism, enzymology, and structural biology topics that build on these fundamentals.