VM

Water and Life — Concepts 3.1–3.3

Concept 3.1: Polar covalent bonds in water molecules result in hydrogen bonding

  • Water is extremely familiar, yet its extraordinary qualities arise from the structure and interactions of its molecules.
  • Water molecule shape: a wide V shape with two hydrogen atoms joined to an oxygen atom by single covalent bonds.
    • Oxygen is more electronegative than hydrogen, so the covalent bonds are polar covalent bonds.
    • Unequal sharing of electrons gives water a polar character: the oxygen atom carries partial negative charges ((\delta^-) ), and the hydrogen atoms carry partial positive charges ((\delta^+)).
  • Hydrogen bonds between water molecules:
    • A partially positive hydrogen of one molecule is attracted to a partially negative oxygen of a nearby molecule.
    • Each water molecule can form hydrogen bonds with several others; these bonds are transient, lasting only a few trillionths of a second in liquid water.
    • In liquid water, hydrogen bonds continually break and re-form, allowing water molecules to slip closer together.
    • In ice, hydrogen bonds are stable and keep molecules farther apart, making ice less dense than liquid water.
  • Visual cue: Regions of opposite partial charges attract to form hydrogen bonds; the central water molecule can hydrogen-bond to multiple neighbors.
  • Key consequence: Hydrogen bonding is the basis for water’s emergent properties and structural organization.

Concept Check 3.1 (summary answers)

  • 1) Electronegativity: the tendency of an atom to attract electrons; in water, O is more electronegative than H, creating polar covalent bonds that enable hydrogen bonding between molecules.
  • 2) The central water molecule can hydrogen-bond to others because the partial charges on O and H create attractive interactions with the partial charges of neighboring water molecules.
  • 3) Two neighboring water molecules arranged with both H’s near O’s would be unfavorable due to electrostatic repulsion; optimal arrangement places the partial positive H near the partial negative O of another molecule.
  • 4) If O and H had equal electronegativity, water would be nonpolar, hydrogen bonds would be negligible, and many of water’s unique emergent properties would be lost.

Concept 3.2: Four emergent properties of water contribute to Earth's suitability for life

  • Water’s emergent properties arise from hydrogen bonding and the polar nature of the molecule:
    • 1) Cohesion of water molecules (and surface tension)
    • 2) Moderation of temperature
    • 3) Expansion upon freezing (ice floats)
    • 4) Versatility as a solvent (hydrophilic vs hydrophobic interactions)

Cohesion of Water Molecules

  • Cohesion: water molecules stay close due to hydrogen bonding; many bonds exist at any moment, giving water a higher level of structural order than most liquids.
  • Surface tension: at the air–water interface, water molecules form an orderly network bonded to water below and to each other, but not to air above, creating a film-like surface that resists stretching.
  • Biological significance: cohesion enables water transport in plants; hydrogen bonds pull water upward through xylem as water evaporates from leaves, transmitting the upward pull from leaves to roots.
    • Adhesion complements cohesion: water clings to cell walls (adhesion), helping resist gravity and aiding ascent in tall trees.
  • Practical note: high surface tension allows some organisms (e.g., raft spiders) to move on water.
  • Example: tallest trees can transport water upward >100 m due to cohesion and adhesion.
  • Related figures: Fig. 3.3 (walking on water), Fig. 3.4 (water transport in plants).

Moderation of Temperature by Water

  • Water acts as a heat bank because it absorbs/releases large amounts of heat with only small changes in its own temperature.
  • Temperature vs. thermal energy:
    • Temperature: average kinetic energy of molecules in a sample.
    • Thermal energy: total kinetic energy, depends on volume.
    • When two objects at different temperatures meet, heat flows from warmer to cooler until equilibrium.
  • Specific heat and heat capacity:
    • Specific heat (in this book): the amount of heat required to raise the temperature of 1 g of a substance by 1°C.
    • Water: specific heat = 1\;\text{cal/(g·°C)} (i.e., 1 cal raises 1 g of water by 1°C).
    • Compared with substances like ethyl alcohol: specific heat ≈ 0.6\;\text{cal/(g·°C)}.
    • Thus, water resists temperature change more than many other liquids.
  • Why high specific heat? Hydrogen bonding: heat must break bonds to raise temperature, and forming bonds releases heat.
  • Relevance to life on Earth:
    • Large bodies of water stabilize climate by absorbing heat in the day and releasing it at night.
    • Coastal temperature moderation supports life; ocean temperatures stay relatively stable, aiding marine life.
    • Organisms, mostly composed of water, resist rapid internal temperature changes.
  • Heat of vaporization: energy required to convert liquid to gas; for water, about 580\;\text{cal/g} at 25°C, higher than many other liquids, due to hydrogen-bond disruption.
  • Evaporative cooling: the most energetic water molecules leave as vapor, cooling the remaining liquid (and nearby surfaces).
    • Analogy: the fastest runners leaving a group reduce the average speed of those remaining.
  • Global and ecological implications: evaporation of tropical surface water drives heat redistribution via moist air that forms rain; evaporative cooling helps stabilize lakes, ponds, and land temperatures.
  • Practical example: sweating and plant leaf cooling rely on evaporative cooling.
  • Figure references: Fig. 3.5 (ocean temperatures), Fig. 3.6 (elephant cooling).

Ice floats on liquid water

  • Ice is less dense than liquid water because hydrogen bonds keep molecules farther apart in the solid lattice.
  • At temperatures above 4°C, water expands as it warms and contracts as it cools—behavior typical for many liquids until 4°C.
  • As temperature falls from 4°C to 0°C, molecules slow enough to form a crystalline lattice, each water molecule hydrogen-bonded to four partners.
  • Ice density: about 10% less dense than liquid water at 4°C, so ice floats.
  • Ecological significance: floating ice insulates the water below, preventing bodies of water from freezing solid and enabling aquatic life to survive under the frozen surface.
  • Climate context: loss of floating sea ice due to climate change threatens Arctic ecosystems and species dependent on ice (e.g., polar bears, seals).
  • Figures: Fig. 3.1 (Earth’s oceans with floating ice), Fig. 3.7 (Arctic ice loss effects).

Water: The Solvent of Life

  • Water as a versatile solvent arises from its polarity and ability to form hydrogen bonds with solutes.
  • Dissolution process (example: table salt NaCl):
    • Water molecules surround individual ions, forming a hydration shell that shields ions from each other and allows dissolution.
    • Na+ is attracted to the partial negative oxygens; Cl- is attracted to the partial positive hydrogens.
    • Result: Na+ and Cl- distributed in solution with water as the solvent.
  • Hydration shells: spheres of water molecules around dissolved ions or polar molecules.
  • Not all dissolving requires ionic solutes; polar molecules (e.g., sugars) can dissolve if they have polar or charged regions capable of hydrogen bonding with water.
  • Hydrophilic vs hydrophobic substances:
    • Hydrophilic: water-loving; polar or ionic, like table salt or sugar, or cotton (cellulose) which forms hydrogen bonds with water but may not dissolve.
    • Hydrophobic: water-fearing; nonpolar substances like vegetable oil that do not mix well with water.
  • Cotton towel example: cellulose fibers absorb water via hydrogen bonding but cellulose does not dissolve.
  • Hydrophobic cells: major components of cell membranes are hydrophobic (lipid-based) to form barriers.
  • Solute concentration in aqueous solutions:
    • Molecular mass and moles: molecular mass is the sum of atomic masses; example sucrose, C extsubscript{12}H extsubscript{22}O extsubscript{11}, has a molecular mass of about 342\;\text{g/mol} (in daltons: 342 Da).
    • Avogadro’s number: 6.02\times 10^{23} entities per mole.
    • 1 g contains about 6.02\times 10^{23} daltons; by definition, this links grams to moles via molecular mass.
    • 1 mole of any substance contains the same number of molecules as 1 mole of any other substance, enabling fixed-ratio experiments.
    • 1 M solution: dissolving 1 mole of solute in enough solvent to make 1 L of solution; example: 1 M sucrose would require 342 g of sucrose dissolved and diluted to 1 L total volume.
  • Practical point: molarity (M) is the unit of concentration often used by biologists for aqueous solutions.
  • Hydrophilic and hydrophobic contrasts tie into real-world biology: dissolved ions in seawater, cellular fluids, and blood composition show water’s solvent versatility.
  • Mastering Biology references: Polarity, hydration shells, and solvent concepts have practical visualization tools in the Mastering Biology eText and animations.

Hydrophilic vs Hydrophobic in Biological Contexts

  • Hydrophilic substances: have an affinity for water, can form hydrogen bonds with water, and often dissolve; examples include many ions and polar molecules.
  • Hydrophobic substances: lack affinity for water, tend to aggregate away from water (e.g., lipids in membranes).
  • Cell structure implication: membranes are built from hydrophobic interiors and hydrophilic exteriors to create barriers while allowing selective transport.

Solute Concentration, Molarity, and Buffers

  • Molarity: ext{Molarity} = \frac{ ext{moles of solute}}{\text{liters of solution}}; standard unit used for solution concentration in biology.
  • Buffers: solutions that resist changes in pH by absorbing or releasing H+ ions when acids or bases are added.
    • Common buffer system in blood: carbonic acid/bicarbonate (H extsubscript{2}CO extsubscript{3} / HCO extsuperscript{−}_{3}).
    • Reaction: \mathrm{H2CO3 \rightleftharpoons H^+ + HCO_3^-}
    • When pH rises, more carbonic acid dissociates to release H+; when pH drops, bicarbonate removes H+ by forming carbonic acid.
  • Ocean acidification (context for buffering and pH): increasing atmospheric CO2 dissolves in seawater, forms carbonic acid, lowers ocean pH, and reduces carbonate ion concentration essential for calcification.
  • Equation of the carbonate system in seawater (simplified):
    • CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3−
    • H+ + CO3^{2−} ⇌ HCO3−
    • CO3^{2−} + Ca^{2+} → CaCO3 (calcification)
  • Practical implication: as oceans acidify, carbonate ion availability drops, hindering calcification in corals and other marine organisms.
  • Visuals: Fig. 3.12 shows atmospheric CO2 and its fate in the ocean; coral calcification responses studied in 2018 CO2 enhancement experiments.

Concept 3.3: Acidic and basic conditions affect living organisms

  • Water autoionization (self-ionization):

    • 2 H2O ⇌ H3O+ + OH− (often simplified as H+ + OH− interplay)
    • In aqueous solution, H+ is commonly represented as H3O+ (hydronium) for practicality: \mathrm{H^+ \equiv H_3O^+}
    • Equilibrium is dynamic; most water remains undissociated: at 25°C, [H+] = [OH−] = 10^{-7}\,\text{M} in pure water.
    • The key point: the product of the concentrations is constant: [\mathrm{H^+}][\mathrm{OH^-}] = 10^{-14} at 25°C.

  • Acids and bases:

    • Acid: increases H+ concentration in solution (e.g., HCl → H+ + Cl−).
    • Base: decreases H+ concentration by accepting H+ or by providing OH− (e.g., NH3 + H+ → NH4+, NaOH → Na+ + OH−).
    • Strong acids/bases dissociate completely (single arrows): HCl, NaOH.
    • Weak acids/bases dissociate reversibly (double arrows): carbonic acid (H2CO3 ⇌ H+ + HCO3−) and ammonia (NH3 + H+ ⇌ NH4+).

  • pH scale:

    • pH = -\log [\mathrm{H^+}]; at 25°C neutral solution has pH 7 with [H+] = [OH−] = 10^{-7}\,\text{M}.
    • Each pH unit represents a tenfold difference in H+ concentration: a solution of pH 3 is 10^3 times more acidic than pH 6.
    • Values outside neutral (0–14) indicate acidic (
    • Many biological fluids have pH around 6–8 (blood ~7.4). Gastric juice is highly acidic (pH ~2).

  • Buffers (in more detail):

    • Buffers resist pH changes by providing a weak acid–base pair that can absorb excess H+ or OH−.
    • Example in blood: carbonic acid–bicarbonate system; maintains stable blood pH despite metabolic processes.

  • Acids, bases, and the carbonic acid–bicarbonate system in life:

    • Carbon dioxide from respiration dissolves in blood to form carbonic acid, contributing to pH regulation.

  • Ocean acidification (revisited):

    • Increased atmospheric CO2 dissolves in seawater, forms carbonic acid, releases H+, lowers pH, and reduces carbonate ions necessary for CaCO3 formation in marine organisms.

  • Concept Check 3.3 (summary):

    • 1) A basic solution at pH 9 has far fewer H+ than an acidic solution at pH 4; there are 10^5 times more H+ in the acidic solution.
    • 2) For 0.01 M HCl, pH = 2.0 (strong acid fully dissociates).
    • 3) For acetic acid (CH3COOH) as a buffer: acid HA ⇌ H+ + A−; identify acid, base, H+ donor, and H+ acceptor.
    • 4) If you add a strong acid to equal volumes of water and acetic acid buffer, pH will shift depending on buffer capacity; buffers resist large pH changes.

  • Scientific Skills Exercise (interpretation of data):

    • An example exercise analyzes how carbonate ion concentration ([CO3^{2-}]) in seawater affects coral reef calcification.
    • Graph interpretation involves identifying axes, independent vs dependent variables, and describing the relationship (negative correlation: higher [CO3^{2-}] corresponds to higher calcification rate).
    • Practical calculations: estimate rates at given [CO2], convert to days for a fixed CaCO3 accumulation (e.g., 30 mmol CaCO3 per m^2).

  • Summary of key connections across concepts:

    • Water’s emergent properties (cohesion, temperature regulation, ice density, solvent power) are all rooted in its polar structure and hydrogen bonding.
    • These properties underpin biological processes (plant transport, protein stability in solution, blood buffering) and ecological phenomena (ocean acidification, coral calcification).
    • The pH system and buffers are essential for maintaining cellular conditions; disruption can alter protein structure, enzyme activity, and metabolic pathways.

  • Quick recap of formulas and constants to remember:

    • Hydrogen ion concentration in neutral water at 25°C: [\mathrm{H^+}] = [\mathrm{OH^-}] = 10^{-7}\,\text{M}
    • Water dissociation product: [\mathrm{H^+}][\mathrm{OH^-}] = 10^{-14}
    • pH definition: \mathrm{pH} = -\log [\mathrm{H^+}]
    • Specific heat of water: c = 1\;\text{cal/(g·°C)}
    • Heat of vaporization of water (approx.): \Delta H_{vap} \approx 580\;\text{cal/g}
    • Avogadro’s number: N_A = 6.02\times 10^{23} molecules/mol
    • Molarity: \text{M} = \frac{\text{moles of solute}}{\text{liters of solution}}
    • Carbonate system (simplified): CO2 + H2O ⇌ H2CO3 ⇌ H^+ + HCO3^- and H^+ + CO3^{2-} ⇌ HCO3^-

  • Cat’s drinking behavior and surface phenomena (extension):

    • Some discussions illustrate how water’s cohesion and adhesion enable the cat to lap water via a column of water that is drawn upward by cohesive forces and the meniscus effect before gravity collapses it. This example highlights the role of cohesion, adhesion, and surface tension in real-world fluid dynamics.

  • Connections to current events and broader implications:

    • Climate change alters Earth’s water-related properties (e.g., sea ice extent, ocean temperatures) due to changing energy balance and greenhouse gas concentrations.
    • Ocean acidification, driven by rising atmospheric CO2, threatens calcifying marine organisms and reef ecosystems by reducing carbonate ion availability for CaCO3 formation.
    • Understanding water’s chemistry helps explain patterns in ecology, physiology, and environmental science, and informs policy and conservation efforts.

Final quick-reference: Key terms

  • Polar covalent bonds, hydrogen bonds, cohesion, adhesion, surface tension, capillary action, specific heat, heat of vaporization, evaporative cooling, density anomaly of water, solvent of life, hydrophilic, hydrophobic, solution, solvent, solute, hydration shell, molarity, mole, Avogadro's number, buffering, pH, acids, bases, carbonic acid, bicarbonate, ocean acidification, calcification, CaCO3, carbonate ions, hydronium (H3O+), hydroxide (OH−)