Water & pH: Comprehensive Study Notes (Bullet-point format)

Water properties, structure, and why it’s special

• Water polarity and hydrogen bonding drive most of its unusual properties.
  • In ice, hydrogen bonds are stable; in liquid water, hydrogen bonds constantly break and re-form.
  • Water is polar, which underpins its solvent abilities and many interactions with biological molecules.
  • Visual cue: H2O with partial charges on H and O; hydrogen bonds link water molecules in a network.
• Special properties stem from polarity and hydrogen bonding; these properties enable water to act as a solvent, influence temperature regulation, and affect biological structures.

Water as a solvent

• When substances dissolve in water, their molecules separate and become surrounded by water molecules.
• Key terms:
  • Solute: the dissolved substance (e.g., sugar).
  • Solvent: the dissolving medium (water).
  • Solution: the resulting mixture.
• Water is an excellent solvent for many substances due to its polar bonds.
• Hydrophilic vs. hydrophobic:
  • Water-soluble hydrophilic molecules are attracted to water and dissolve readily.
  • Hydrophobic molecules resist water and tend to separate from water.
• Examples:
  • Sugar dissolves in water; sugar is hydrophilic (polar) and can dissolve because of water’s polarity.
  • Lipid bilayers are hydrophobic and resist dissolution in water; this hydrophobicity is essential for membrane integrity.

Hydrophilic vs. hydrophobic and cell membranes

• Hydrophilic molecules: ions or polar molecules that interact favorably with water via charges or hydrogen bonding.
• Ionic substances (e.g., NaCl) dissolve because water is attracted to ions.
• Polar molecules (e.g., urea) dissolve because they form hydrogen bonds with water.
• Hydrophobic molecules do not dissolve well in water; lipid bilayers are hydrophobic in their interior, helping maintain cell membrane integrity.

Water’s surface phenomena: cohesion, adhesion, and surface tension

• Cohesion: water–water attraction; explains high surface tension.
  • Cohesion contributes to phenomena like surface tension and water transport in capillaries.
• Adhesion: water–polar surface attraction; enables capillary action.
• Surface tension relevance:
  • In alveolar air sacs, surface tension would collapse the thin water layer unless mitigated.
  • Pulmonary surfactant (a phospholipid) reduces surface tension, preventing alveolar collapse; premature infants may have surfactant deficiencies.

Capillary action and plant water transport

• Adhesion of water to polar surfaces (e.g., cell walls) plus cohesion among water molecules drives capillary action.
• Capillary action helps water move up narrow tubes due to increased surface area for adhesion.
• In plants, this underpins the movement of water from roots to shoots via xylem.
• Real-world note: water movement direction in xylem is driven by adhesion to the walls and cohesion between water molecules, aided by root pressure and transpiration pull.

Water’s high specific heat and evaporative cooling

• Specific heat: the amount of heat needed to change the temperature of 1 g of a substance by 1°C.
  • Water has a high specific heat, so large bodies of water moderate climate and help organisms thermoregulate.
• Evaporative cooling: water requires significant energy to evaporate, which helps cool organisms as water on surfaces or skin evaporates.

Ice, density, and implications for biology

• Ice floats on liquid water because ice is less dense than liquid water (water expands when it freezes).
• As water freezes, the molecules are farther apart, creating a lattice that is less dense.
• Floating ice insulates the water below, protecting aquatic ecosystems during cold periods.
• Question raised: what can expansion and ice crystal formation do to the integrity of cell membranes and cell walls? They can disrupt them.

Strategies to reduce freeze damage

• Anhydrobiosis: near-total loss of water in a dormant state; observed in many prokaryotes, some protists, and some multicellular organisms; often in spores or seeds.
• Antifreeze molecules: disrupt the hydrogen-bond lattice of ice crystals.
  • Examples: ethylene glycol, glycerol, urea; some organisms (e.g., wood frog) naturally accumulate these compounds.
• Antifreeze proteins: small proteins that interfere with crystal formation; their action is linked to specific structural features (e.g., interaction with water, disruption of lattice).
• GMO approaches are being explored (e.g., frost-proof strawberries) but practical crops with animal antifreeze traits are not yet commercially available.

Antifreeze proteins and related concepts

• Antifreeze proteins have regions (e.g., a blue-interacting portion and a threonine-rich region) that disrupt ice-crystal formation, contributing to freeze tolerance.
• There is ongoing discussion about applying these concepts to crops; currently, many discussions involve non-GMO approaches or misinterpretations in popular media.

pH and acid–base balance: fundamentals

• pH is a measure of the hydrogen ion concentration [H⁺] in water; the more H⁺, the more acidic.
• Definition:
  • $\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$
• Pure water has [H⁺] = 10⁻⁷ M, so pH = 7 (neutral).
• For acids and bases, remember the idea of deprotonation (loss of H⁺) and protonation (gain of H⁺).
• Hydrogen ion exchange (autoionization) in water:
  • One molecular event: $2\,\mathrm{H<em>2O} \rightleftharpoons \mathrm{H</em>3O^+} + \mathrm{OH^-}$
  • Pure water maintains steady-state concentrations of H⁺ and OH⁻, each around 10⁻⁷ M at 25°C.
• The pH scale is logarithmic: small changes in pH reflect large changes in [H⁺].
  • Example: a solution with pH = 3 has 10³ times more H⁺ than a solution with pH = 6, i.e., 1000× more.
• Example values from common substances (approximate):
  • Battery acid: pH ≈ 0
  • Gastric juice: pH ≈ 1
  • Lemon juice: pH ≈ 2
  • Vinegar, wine, soda: pH ≈ 2–3
  • Black coffee: pH ≈ 5
  • Pure water: pH ≈ 7
  • Seawater: pH ≈ 8
  • Milk of magnesia: pH ≈ 10
  • Ammonia: pH ≈ 11
  • Bleach: pH ≈ 12–14

Acids, bases, and buffering terminology

• Release of a proton (H⁺) is called deprotonation; acceptance of a proton is protonation.
• An acid is typically a substance that has already donated a proton (deprotonated after donation) and carries a negative charge; a base is typically one that has accepted a proton and carries a positive charge.
• Related naming:
  • Deprotonation (noun); deprotonate (verb) = release of H⁺
  • Protonation (noun); protonate (verb) = acceptance of H⁺

pH indicators and buffers

• Red cabbage indicator uses anthocyanin pigments that shift color with pH (acidic, neutral, basic ranges) as shown in structural schematic with R groups (R1, R2) and glycosylation (OGly).
• Buffers resist changes in pH by providing a weak conjugate acid–base pair.
  • Common buffers and their approximate pKa values observed in this course materials:
  • NH4⁺/NH3 buffer: pKa ≈ 9.25
  • H2CO3/HCO3⁻ buffer: pKa ≈ 6.37
  • HCO3⁻/CO2 buffer relation shown as a buffer equilibrium with a pKa around 10.3 (context-dependent; often shown in educational materials as related to carbonate system equilibria)
  • H3PO4/H2PO4⁻ buffer: pKa ≈ 2.12
  • H2PO4⁻/HPO4²⁻ buffer: pKa ≈ 7.21
  • HPO4²⁻/PO4³⁻ buffer: pKa ≈ 12.32
• Ocean/physiological buffering relies on carbonic acid-bicarbonate system and phosphate buffering among others.

Ocean acidification and atmospheric CO2

• Hydration of CO2 in water (ocean):
  • CO2 + H2O ⇌ H2CO3 ⇌ H⁺ + HCO3⁻
• If oceanic CO2 increases (due to higher atmospheric CO2), ocean pH decreases (ocean becomes more acidic).
• This is a central problem in ocean chemistry and climate science.

The hydrosphere and elemental composition of seawater

• Essential elements in seawater originate from weathering and erosion of the lithosphere, carried to the oceans and concentrated as salts.
• Seawater salinity is about 35 g of salts per kg of seawater (35‰).
• Major ions and their approximate fractional composition by weight in seawater (as provided in course materials):
  • Na⁺: ~30.6%
  • Cl⁻: ~55.0%
  • K⁺: ~1.1%
  • Ca²⁺: ~1.2%
  • Mg²⁺: ~3.7%
  • Others (Sr, Br, C, etc.): ~7.7%
• Common osmolarity reference: salinity ~3.5% in marine water; estuaries/brackish water ~0.9%; freshwater contains only trace amounts of salts.
• “Normal saline” used in medical and lab applications is about 0.9% (osmolarity similar to body fluids).

Quick reference: pH implications for biology

• Life can adapt to pH extremes (acidophiles in highly acidic environments like the Danakil Depression, pH ~ 0.2; alkaliphiles in alkaline environments like Mono Lake, pH ~ 9.8).
• However, life cannot tolerate large pH swings outside the adapted range; proteins and enzymes are highly pH-sensitive.
• In clinical scenarios, a drop in blood pH indicates acidosis, while an increase indicates alkalosis, with pH 7.4 often cited as the normal human blood value.
• Indicators of pH changes in biology rely on buffers and specific molecules that shift color or reactivity with pH changes.

Connections to broader principles and real-world relevance

• Water’s solvent properties enable biochemistry: solubility of ions and polar molecules is central to metabolism, signaling, and transport.
• Cohesion and adhesion underpin plant transpiration, capillary action in tissues, and blood flow on micro scales.
• Heat capacity and evaporative cooling explain climate moderation by oceans and thermoregulation in organisms.
• Ice physics informs ecology (ice insulation) and medical/science challenges (freeze damage to cells, cryopreservation strategies).
• pH control via buffers is essential in physiology, environmental science (ocean acidification, acid rain), medicine, and laboratory work.
• Understanding salinity and ion composition of seawater informs marine biology, oceanography, and desalination technologies.

Key formulas and numerical references (LaTeX)

• pH definition:
  $\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$
• Water autoprotolysis (hydrogen ion exchange):
  $2\,\mathrm{H<em>2O} \rightleftharpoons \mathrm{H</em>3O^+} + \mathrm{OH^-}$
• Ocean CO2 hydration and carbonate system:
  $\mathrm{CO<em>2} + \mathrm{H</em>2O} \rightleftharpoons \mathrm{H<em>2CO</em>3} \rightleftharpoons \mathrm{H^+} + \mathrm{HCO_3^-}$
• Scale of pKa values from buffers (approximate):
  • NH4^+/NH3: $\mathrm{p}K_a \approx 9.25$
  • H2CO3/HCO3^-: $\mathrm{p}K_a \approx 6.37$
  • HCO3^-/CO2 (carbonate system context): $\mathrm{p}K_a \approx 10.3$
  • H3PO4/H2PO4^-: $\mathrm{p}K_a \approx 2.12$
  • H2PO4^-/HPO4^{2-}: $\mathrm{p}K_a \approx 7.21$
  • HPO4^{2-}/PO4^{3-}: $\mathrm{p}K_a \approx 12.32$
• Temperature and biological buffering references can shift slightly with conditions; values above are educational anchors from the slides.