Water, Hydrogen Bonding, and Noncovalent Interactions - Practice Flashcards
Water and Its Role in Biology: Comprehensive Notes
Overview
Today's session covers why water is central to biochemistry and physiology: solvent properties, hydrogen bonding, noncovalent interactions, hydrophobic effects, osmosis, and acid-base chemistry related to pH.
Emphasis on how water determines structure and function of proteins, nucleic acids, and membranes; and how hydration influences metabolism, temperature regulation, and waste elimination.
The Importance of Water in Biology
Life evolved in water; water offered UV protection and a harsh but adaptable aqueous environment.
Organisms typically contain 70–90% water; many reactions occur in aqueous environments.
Water content in humans varies with metabolic status, age, and gender: 55–80% body water.
Water intake comes from food and drinks; water is eliminated via breathing, urine, sweat, and bowel movements; maintenance of water balance is continuous.
Where Water Is Found in the Body and Its Functions
Cells: cytoplasm is primarily water.
Interstitial spaces (between cells) and blood are water-rich environments.
Physiological roles of water:
Brain uses water to manufacture hormones and neurotransmitters.
Saliva production and moisture for mucous membranes; saliva acts as a lubricant and aids digestion.
Regulation of body temperature via sweating and respiration.
Water acts as a shock absorber for brain and spinal cord.
Amniotic fluid in pregnancy is largely water.
Water converts ingested food into usable components; helps deliver oxygen; aids elimination of wastes.
Water as a Biochemical Solvent
Water is an ideal biological solvent due to its dipolar nature and hydrogen-bonding network.
Water can act as both hydrogen bond donor and acceptor.
Structure of water drives solvent properties and interactions with biomolecules.
Water Molecule: Structure and Bonding Basics
Chemical formula: ext{H}_2 ext{O}; two hydrogens covalently bonded to one oxygen.
Oxygen has atomic number 8 and an electronic configuration that yields two unpaired electrons in the outer shell, enabling covalent bonding with hydrogen.
Electronegativity: Oxygen is more electronegative than hydrogen, creating a polar covalent bond.
Resulting dipole: Hydrogen bears a partial positive charge ($\delta^+$) and oxygen bears a partial negative charge ($\delta^-$).
Geometry: Water has a bent geometry with an angle slightly less than the ideal tetrahedral angle: approximately 104.5^{\circ}; due to electronegativity, the bond angle is around \sim 105^{\circ}, not the perfect 109.5^{\circ}.
Bonding concepts:
Covalent O–H bonds: electrons are shared between O and H.
Lone pairs: Oxygen has two lone pairs of electrons left after bonding.
Dipole moment: water is a polar molecule due to unequal sharing of electrons.
Hydrogen Bonding: Donor, Acceptor, and Strength
Definition: a hydrogen bond is an electrostatic interaction between a hydrogen attached to a highly electronegative atom (donor) and another electronegative atom with lone pairs (acceptor).
Water can act as both hydrogen bond donor and acceptor.
Common hydrogen-bonding partners: electronegative centers such as oxygen and nitrogen.
Orientation matters: hydrogen bonds are strongest when donor and acceptor atoms are aligned in a straight line (colinear) to maximize electrostatic interaction; bent arrangements yield weaker bonds.
Bond energies:
Hydrogen bond: typically 4-6\ \text{kJ/mol} for neutral atoms; 6-10\ \text{kJ/mol} when one partner is charged.
Lifetime and dynamics:
Individual hydrogen bonds are short-lived, with lifetimes around \sim 20\ \text{ps} (picoseconds).
At any instant, liquid water participates in roughly 3.4 hydrogen bonds per molecule on average; water can transiently form up to four hydrogen bonds, but the network is constantly reorganizing.
Dipole-dipole interactions form the basis for water’s high boiling point, high melting point, surface tension, cohesion, and adhesion.
Hydrogen bonding also explains why ice floats and why water has unique solvent properties.
Water's Phases and Hydrogen-Bond Network
Ice (hexagonal form) vs liquid water:
Ice features a lattice in which each water can form up to four hydrogen bonds, leading to a highly ordered, low-entropy structure.
Ice has lower density than liquid water, so ice floats.
As temperature rises from ice to water to vapor, hydrogen bonds progressively break, increasing molecular disorder (entropy) and changing structural arrangements.
In liquid water, bonds are continually breaking and reforming, contributing to high surface tension and cohesive/adhesive properties.
Water as a Solvent: Solubility and Hydration
Water dissolves charged and polar substances well; nonpolar substances are poorly soluble (the classic polar-polar vs nonpolar rule).
Hydration and solvation shells form around solutes, stabilizing ions and polar molecules.
Salt dissolution example: water hydrates ions (e.g., Na⁺, Cl⁻) by orienting water’s dipoles around ions, reducing lattice energy and increasing entropy.
Sugars and carbohydrates dissolve due to hydrogen-bonding capacity from their hydroxyl groups (OH) with water.
Amphipathic molecules (containing both polar and nonpolar parts) form micelles in water: hydrophobic tails cluster away from water, while polar heads interact with water, releasing ordered water molecules and increasing overall entropy.
Lipid bilayers and micelles are stabilized by the hydrophobic effect; polar head groups interface with water, while nonpolar tails are sequestered inside.
Common solubility rule: “like dissolves like” — polar substances dissolve well in water; nonpolar substances tend to be insoluble unless they have amphipathic regions.
Hydrophobic Effect: Entropy-Driven Assembly
Hydrophobic effect arises because water reorganizes around nonpolar solutes to maximize hydrogen bonding among its molecules, which reduces entropy.
To regain entropy, water expelled from around nonpolar regions leads to aggregation of nonpolar moieties (e.g., lipid tails clustering in micelles or lipid bilayers forming structures that reduce surface area exposed to water).
Metaphor used in lecture: a party where the teacher (hydrophobic force) reduces freedom of students (water molecules) around nonpolar guests; clustering of nonpolar tails reduces the ‘ordered water’ around them, increasing overall entropy.
Hydrophobic regions drive protein folding, protein-protein interactions, and micelle formation; they are also critical for binding hydrophobic ligands and steroid hormones.
Noncovalent Interactions: Summary and Roles
Ionic (electrostatic) interactions: between permanently charged species or between ions and permanent dipoles.
Hydrogen bonds: noncovalent interactions involving dipoles, donor and acceptor atoms (often N and O); no full electron transfer, but a strong directional attraction.
Van der Waals (dispersion/attractive) interactions: universal, weak individually but cumulatively strong; dependent on distance and orientation; contribute to steric complementarity and DNA base stacking.
Hydrophobic interactions: aggregation of nonpolar regions in aqueous environments due to water’s drive to maximize hydrogen bonding and entropy.
Orientation and proximity govern whether interactions form and stabilize macromolecular structures (proteins, nucleic acids, enzyme-substrate complexes).
Example applications:
DNA base pairing: A–T (2 H-bonds) and G–C (3 H-bonds) stabilize the double helix.
Protein folding and receptor-ligand binding: multiple weak interactions collectively stabilize complexes.
Protein and Nucleic Acid Structure: The Role of Hydrogen Bonding and Van der Waals
Hydrogen bonds can occur within polypeptides, between peptide backbones, or between side chains that are positioned to interact.
Covalent bonds (e.g., disulfide bonds between cysteine residues) can stabilize secondary and tertiary structures; noncovalent interactions (H-bonds, ionic interactions, Van der Waals, hydrophobic effects) also stabilize folds and interactions.
Orientation is critical: proper alignment of residues and functional groups enables noncovalent interactions essential for structure and function.
Water as a Solvent: Practical Examples and Implications
Water dissolves salts by hydrating ions, reducing lattice energy, and increasing entropy.
Simple polar molecules (e.g., ethanol, glycerol) dissolve well due to their ability to hydrogen bond with water.
Nonpolar gases like O₂ and N₂ are poorly soluble in water because they cannot form hydrogen bonds with water; CO₂ is nonpolar and has limited solubility but can dissolve and form carbonic acid in water.
Ammonia is highly soluble in water due to its own polarity and hydrogen bonding capability; the interaction with water relates to metabolic pathways (e.g., urea cycle) and nitrogen metabolism.
Osmosis and Osmotic Pressure
Cytoplasm is a highly concentrated aqueous solution; osmotic pressure drives water movement across membranes.
Isotonic: water influx equals efflux; cells maintain shape and size.
Hypertonic environment: higher external solute concentration draws water out; cells shrink (crenation in red blood cells is an analogue).
Hypotonic environment: lower external solute concentration draws water in; cells swell and may lyse.
Terms:
Endosmosis: water moving into a cell (hypotonic environment).
Exosmosis: water moving out of a cell (hypertonic environment).
Raisins vs grapes analogy: raisins in water swell due to osmosis; grapes in sugar solution shrink due to water movement out of cells.
Acids, Bases, and pH: Water’s Autoprotonation and pH Scale
Definitions:
Acid: substance that donates a proton (H⁺); conjugate base is what remains after donation.
Base: substance that accepts a proton; conjugate acid is the species formed after accepting a proton.
Conjugate pairs (acid/base) are central to buffer systems.
Water is amphoteric: it can act as both an acid (proton donor) and a base (proton acceptor).
Autoionization of water:
ext{H}_2 ext{O}
ightleftharpoons ext{H}^+ + ext{OH}^-In reality, the process forms hydronium ions, often represented as ext{H}_3 ext{O}^+ and ext{OH}^-, with solvent stabilization via hydrogen bonding.
Ionic product of water (Kw):
General expression: K_w = [ ext{H}^+][ ext{OH}^-]
Transcript states: Kw \approx 1.8 \times 10^{-16} at 25°C (note: the standard accepted value is Kw = 1.0 \times 10^{-14} at 25°C; the transcript’s number appears to be a misquote or an error). The key concept is that at a given temperature, the product of the two ion concentrations is constant.
pH and pOH:
ext{pH} = -\log_{10} [ ext{H}^+]
ext{pOH} = -\log_{10} [\text{OH}^-]
At 25°C: ext{pH} + \text{pOH} = 14
Neutral pH is 7, where [ ext{H}^+] = [\text{OH}^-] = 1.0 \times 10^{-7} \text{ M}
Calculation example (from lecture): If [ ext{H}^+] = 10^{-6} \text{ M}, then ext{pH} = 6 and, with Kw ≈ 1.0 × 10^{-14}, [ ext{OH}^-] = \frac{K_w}{[ ext{H}^+]} = 1.0 \times 10^{-8} \text{ M}, giving ext{pOH} = 8.
Educational note: students practiced calculating [ ext{OH}^-] from a given [ ext{H}^+] using Kw; the instructor emphasized memorizing Kw and the pH/pOH relationships, and provided practice problems.
Real-world contexts:
Blood is slightly alkaline; beverages and foods span acidic to basic ranges.
Neutral pH (7) is a reference point used in biology and medicine.
Practical Calculations and Concepts Discussed
Molarity of water:
1 liter of water has approximately 1000 g; molecular weight of water is 18 g/mol.
Molarity of pure water ≈ \frac{1000\text{ g}}{18\text{ g/mol}} ≈ 55.5\text{ M}
Ionic product of water (Kw) and the pH scale form the basis for buffer calculations and pH adjustments in biochemical systems.
The equation Kw = [ ext{H}^+][ ext{OH}^-] leads to the relation [ ext{H}^+] = \frac{Kw}{[ ext{OH}^-]} and vice versa.
The log relationships: ext{pH} = -\log{10} [\text{H}^+] and ext{pOH} = -\log{10} [\text{OH}^-], with ext{pH} + \text{pOH} = 14 at standard conditions.
Exercise prompts: compute the hydroxide concentration from a given hydrogen ion concentration; interpret pH values in practical terms.
Key Takeaways and Connections to Foundational Principles
Hydrogen bonding underpins water’s unique properties and explains water’s high boiling and melting points, surface tension, and solvent behavior.
Noncovalent interactions—ionic, hydrogen bonding, Van der Waals, and hydrophobic effects—collectively stabilize biomolecular structures and drive molecular recognition, enzyme-substrate binding, and receptor-ligand interactions.
The hydrophobic effect arises from water’s tendency to maximize hydrogen-bonding opportunities, driving nonpolar regions to aggregate; this mechanism is critical in protein folding and membrane organization.
The concept of hydrophobicity integrates thermodynamics (entropy changes) with structural biology to explain macromolecular assembly.
Understanding Kw, pH, and pH-dependent biology is essential for biochemistry, physiology, and pharmacology.
Clarifications and Common Questions from the Session
Difference between covalent and hydrogen bonds:
Covalent bond: actual sharing of electrons (solid line in diagrams).
Hydrogen bond: noncovalent interaction with no electron sharing, shown as dashed/dotted lines; donor and acceptor atoms involved.
Water as an acid and a base (amphoterism): water can donate a proton to act as an acid, or accept a proton to act as a base.
The role of water in gas, liquid, and solid phases: hydrogen bonding is strongest in ice (four H-bonding opportunities per molecule in an ideal lattice) and weaker but highly dynamic in liquid water, leading to phase-dependent properties.
Quick Reference Formulas and Values (LaTeX)
Water molecule: ext{H}_2 ext{O}
Dipole and partial charges: \delta^+ on H, \delta^- on O
Hydrogen bond energy (range): E_{HB} \approx 4-10\ \text{kJ/mol} depending on environment
Covalent O–H bond energy (rough): E_{HO\text{ covalent}} \approx 420\ \text{kJ/mol}
Hydrogen bond length: approximately 1.7\ \text{Å} \approx 0.17\ \text{nm}
Hydrogen bond lifetime in liquid water: ≈ \sim 20\ \text{ps}
Water’s molarity (approximate): [\text{H}2\text{O}]{\text{l}} \approx 55.5\ \text{M}
Ionic product of water (Kw): K_w = [\text{H}^+][\text{OH}^-]
Preferred historically cited value (experimental): K_w \approx 1.0\times 10^{-14} at 25°C; lecture notes mention a value around 1.8\times 10^{-16} (note the discrepancy with the standard value)
Neutral pH: \text{pH} = 7 when [\text{H}^+] = [\text{OH}^-] = 1.0\times 10^{-7}\ \text{M}
pH and pOH relation: \text{pH} + \text{pOH} = 14 at 25°C
Study Tips and Relevance for Exams
Memorize the basic water properties derived from hydrogen bonding: high boiling point, high surface tension, cohesion/adhesion, and the ability to form extensive hydrogen-bond networks.
Be able to distinguish covalent bonds from hydrogen bonds and identify donor/acceptor roles in diagrams or questions.
Understand the hydrophobic effect as an entropy-driven process and its role in micelle formation and protein folding.
Know the essential equations for Kw, pH, and pOH, and be able to perform simple calculations given either [H⁺] or [OH⁻].
Remember the real-world contexts: osmosis in cells, isotonic/ hypotonic/hypertonic solutions, and the importance of solvent properties in drug design and biochemistry.
Note on Breaks and Class Logistics (as mentioned in transcript)
A brief five-minute break was conducted during the session to manage rest and attention.
A quick reminder about calculator usage in the assessment: exams/tools may provide calculators; verify with the instructor.
Summary of Key Pointers to Remember
Water is essential not only as a solvent but as an active participant in biomolecular structure and function through hydrogen bonding and noncovalent interactions.
The hydrophobic effect is central to biomolecular organization, protein folding, membrane formation, and the behavior of amphipathic molecules.
Osmosis and osmotic pressure play critical roles in cellular homeostasis; isotonic solutions prevent cellular edema or shrinkage.
Acid-base chemistry in water defines pH, pOH, and Kw; proton hopping and the hydronium/hydroxide dynamics explain water’s conductivity and buffering capacity.
Quick Final Check: Sample Question Reflection
Which two weak interactions are most important in forming a cluster of lipid molecules in aqueous solution? Answer: Hydrophobic effect plus hydrogen bonding (the hydrophobic effect relies on water hydrogen-bond reorganizations; thus option E in the lecture examples).
End of Notes
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