Chemistry of Life: Water Properties

Why Study Water

• All life occurs in water—inside and outside the cell; water is the molecule that supports all of life.

Polarity and Hydrogen Bonding

• The polarity of water molecules makes them able to form hydrogen bonds with each other.
• Water is a polar molecule; polarity drives hydrogen bonding and contributes to water’s unique properties.
• Hydrogen bonds occur between a partially positive hydrogen atom of one water molecule and a partially negative oxygen atom of a neighboring water molecule.
• Conceptual note: Hydrogen bonds help explain why water behaves as a “sticky” molecule and why it exhibits cohesion and adhesion.

Special Properties of Water

1) Cohesion & Adhesion, surface tension, capillary action

• Cohesion: the bonding of a high percentage of water molecules to neighboring water molecules via hydrogen bonds.
• Adhesion: water’s attraction to other substances (e.g., to plant xylem walls).
• Surface tension: a measure of how hard it is to break the surface of a liquid; related to cohesion.
• Capillary action: combined effect of cohesion and adhesion that helps water move up narrow tubes (e.g., plant vessels).
  2) Good solvent
• Water dissolves many substances; polar water molecules surround and stabilize ions and polar solutes.
• Hydrophilic vs. hydrophobic:
  • Hydrophilic: substances with attraction to water (polar or ionic).
  • Hydrophobic: substances lacking attraction to water (non-polar).
    3) Lower density as a solid
• Ice is less dense than liquid water and floats; ice forms a crystalline lattice in which each water molecule is hydrogen-bonded to ~4 partners.
• This arrangement creates more open space and makes ice about ~10% less dense than liquid water.
  4) High specific heat
• Water resists temperature change; it takes a large amount of heat to raise or lower water’s temperature.
• This occurs because energy goes into breaking/forming hydrogen bonds rather than changing kinetic energy immediately.
  5) High heat of vaporization
• A large amount of energy is required to convert liquid water to water vapor; this enables evaporative cooling.

How Water Enables Plant and Biological Transport

• Cohesion and adhesion drive capillary action in plants; water climbs through xylem via these forces.
• Transpiration in plants is built on cohesion & adhesion; water drawn up through microscopic vessels due to hydrogen bonding and adhesion to walls.
• A meniscus forms due to adhesion; water climbs up surfaces like paper towels or cloth due to adhesive interactions.

Water as a Solvent: Dissolving Substances

• Polarity makes H2O a good solvent: polar H2O molecules surround positive and negative ions, dissolving solutes to create solutions.
• Hydration shells form around ions and polar molecules, enabling them to be dispersed in water.
• Hydrophilic substances dissolve readily in water;
  • Example: polar molecules or ions (e.g., salts, some proteins or amino acids).
• Hydrophobic substances do not dissolve well in water (non-polar substances like fats and hydrocarbons).
• In aqueous biochemical environments, most biochemical reactions occur in water, so concentration calculations in aqueous solutions are essential.

Ice, Water, and Environmental Impacts

• Ice floats because ice is less dense than liquid water; the hydrogen-bonded lattice keeps molecules farther apart.
• Oceans and lakes don’t freeze solid because surface ice insulates the water below, allowing life to persist in winter.
• Seasonal turnover in lakes: sinking cold, oxygen-rich water mixes with deeper layers in spring and nutrients cycle in autumn.

Water and Temperature Regulation

• Specific heat capacity: water’s ability to absorb and retain energy helps moderate Earth's climate.
• Large bodies of water store heat and release it gradually, stabilizing air temperatures.
• This moderating effect helps maintain conditions favorable for life.

Heat of Vaporization and Evaporative Cooling (Illustrated Concept)

• Evaporative cooling occurs as water absorbs heat to vaporize, removing heat from surfaces (e.g., organisms using sweating or evaporative cooling).
• Temperature vs. phase: water transitions among solid (ice), liquid, and gas (water vapor) with energy input/output reflecting hydrogen bond dynamics.

Hydrogen Bonding: Chemical and Life Implications

• Hydrogen bonding provides cohesion, adhesion, surface tension, and solvent properties that underpin many biological processes.
• Universal solvent concept: water can dissolve a wide range of substances due to polarity and hydrogen bonding.
• The combination of polarity, hydrogen bonding, and solvent properties supports the chemistry of life in aqueous environments.

Ionization of Water and pH

• Occasionally a hydrogen atom participating in a hydrogen bond shifts between molecules; the transferred proton leaves behind an electron, forming a hydrogen ion (H^+) and a hydroxide ion (OH^-):
  $\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}$
• In aqueous solution, protons associate with water to form hydronium ions (H3O^+): $\mathrm{H</em>2O} + \mathrm{H^+} \rightarrow \mathrm{H_3O^+}$
• Water ionizes to affect pH, defined as a measure of hydrogen ion concentration:
  $\mathrm{pH} = -\log [\mathrm{H^+}]$
• Neutral water corresponds to equal concentrations of H^+ and OH^- (pH = 7 at 25°C under standard conditions).

Acids and Bases

• Acids: donate protons (H^+) to water, forming hydronium (H_3O^+) ions; have pH < 7.
• Bases: donate hydroxide (OH^-) or accept protons; have pH > 7.
• Strong acids (e.g., H2SO4, HCl) completely dissociate in water to yield ions.
• Strong bases (e.g., NaOH) completely dissociate to yield ions.
• Weak acids (e.g., acetic acid, CH_3COOH) do not completely dissociate; partial ionization occurs.
• Weak bases also do not completely dissociate in water.

Examples of Acids and Bases

• Strong acids: HCl, H2SO4 show 100% ionization in water.
• Weak acids: acetic acid (HA) partially ionizes to H^+ and A^- (e.g., Ac^- in solution).
• Strong bases: NaOH, KOH fully dissociate to Na^+/K^+ and OH^-.
• In practice, pH depends on the balance between H^+ and OH^- concentrations.

pH Scale and Typical Values

• pH ranges from 0 (highly acidic) to 14 (highly basic); 7 is neutral.
• Examples of pH values (approximate):
  • Stomach acid: ~1
  • Lemon juice: ~2
  • Vinegar: ~2–3
  • Tomatoes: ~4
  • Black coffee: ~5
  • Pure water: ~7
  • Seawater: ~8
  • Baking soda solution: ~9–10
  • Household ammonia: ~11–12
  • Oven cleaner: ~13–14
• A tenfold change in H^+ concentration corresponds to a change of 1 pH unit.

Buffers and Maintenance of pH

• Buffers are reservoirs of H^+ that donate H^+ when [H^+] falls and absorb H^+ when [H^+] rises, helping maintain pH.
• Typical biological buffers exist to keep cellular pH within narrow ranges essential for structure and function.
• In many biological systems, pH must be tightly controlled because small pH changes can alter molecule shapes and functions.

Physiological Buffer System: Blood pH and Carbonate Buffer

• Blood pH must be maintained between ~7.38 and 7.42.
• Major components:
  • HCO_3^- (bicarbonate): a weak base
  • H2CO3 (carbonic acid): a weak acid
  • They exist in equilibrium:
    $\mathrm{CO<em>2 + H</em>2O \rightleftharpoons H<em>2CO</em>3 \rightleftharpoons HCO_3^- + H^+ }$
• When [H^+] increases (e.g., during strenuous exercise or acidic drug overdose), the equilibrium shifts to the left, consuming H^+ and forming more H2CO3 and CO_2 to be exhaled, helping restore pH.
• Buffers in blood help stabilize pH during metabolic and respiratory changes.

Key Connections and Implications

• pH affects the shape of molecules and, therefore, their function; cellular processes depend on maintaining proper pH.
• The ability of water to act as both solvent and participant in chemical equilibria enables complex biochemistry in living systems.
• Hydration shells around ions and molecules enable many metabolic reactions to occur in solution.
• The buffering system in blood illustrates how organisms regulate internal environments (homeostasis) in the face of external and internal perturbations.

Quick Reference: Core Equations and Concepts

• Water autoprotolysis (ionization):
  $\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}$
• Hydronium formation in solution:
  $\mathrm{H<em>2O} + \mathrm{H^+} \rightarrow \mathrm{H</em>3O^+}$
• pH definition:
  $\mathrm{pH} = -\log [\mathrm{H^+}]$
• Carbonic-bicarbonate buffering system (blood):
  $\mathrm{CO<em>2 + H</em>2O \rightleftharpoons H<em>2CO</em>3 \rightleftharpoons HCO_3^- + H^+ }$
• Ice vs liquid water density and hydrogen-bonding context: ice forms a crystalline lattice with each molecule hydrogen-bonded to ~4 partners; ice ~10% less dense than liquid water, enabling floating ice and insulation of deeper liquid water.
• Hydrogen bonding basis for cohesion, adhesion, surface tension, capillary action, and solvent abilities.

Conceptual Takeaways

• Water’s polarity and hydrogen bonding underpin its unique physical properties that support life.
• Water as solvent enables chemistry of life by dissolving substances and forming hydrated ions/molecules.
• Water’s heat capacity and heat of vaporization contribute to climate regulation and evaporative cooling in organisms.
• pH and buffers are central to maintaining stable cellular environments necessary for enzymes and metabolic pathways.
• Environmental and physiological processes (ice formation, lake turnover, blood buffering) illustrate the real-world importance of water chemistry.