Water and Life: Properties, Structure, and Biological Importance

Water and Life

Overview: The Molecule That Supports All of Life

Water is essential for all known forms of life. Its unique chemical and physical properties make it indispensable for biological processes and the maintenance of life on Earth.

Polar Covalent Bonds and Hydrogen Bonding in Water

Structure and Polarity of Water Molecules

• Polar covalent bonds in water molecules cause electrons to spend more time near the oxygen atom than the hydrogen atoms.

• This results in a polar molecule with an uneven distribution of charge (partial negative on oxygen, partial positive on hydrogens).

• The polarity enables water molecules to form hydrogen bonds with each other.

Example: The partial charges on water molecules allow them to attract each other, forming a network of hydrogen bonds.

Emergent Properties of Water

Four Properties That Facilitate Life

• Cohesive behavior

• Ability to moderate temperature

• Expansion upon freezing

• Versatility as a solvent

Cohesion and Surface Tension

• Cohesion is the attraction between water molecules due to hydrogen bonding, resulting in high surface tension.

• Surface tension makes water behave as if coated with an invisible film, allowing small objects or insects to rest on its surface.

• Adhesion is the attraction between water molecules and other substances, such as plant cell walls, aiding in water transport against gravity.

Example: Water moves upward in plants due to cohesion and adhesion, enabling the transport of nutrients from roots to leaves.

Moderation of Temperature by Water

• Water absorbs heat from warmer air and releases it to cooler air, moderating temperature changes.

• It can absorb or release large amounts of heat with only slight changes in its own temperature.

Temperature and Heat

• Kinetic energy: Energy of motion.

• Thermal energy: Kinetic energy associated with random motion of atoms or molecules.

• Temperature: Average kinetic energy of molecules in a body of matter.

• Heat: Thermal energy transferred from one body to another.

• Calorie (cal): Heat required to raise 1 g of water by 1°C.

• Joule (J): SI unit of energy; $1\,\text{J} = 0.239\,\text{cal}$, $1\,\text{cal} = 4.184\,\text{J}$.

Water’s High Specific Heat

• Specific heat: Amount of heat needed to change 1 g of a substance by 1°C.

• Water’s specific heat: $1\,\text{cal}/(\text{g} \cdot ^\circ\text{C})$.

• Hydrogen bonding causes water to resist temperature changes, minimizing temperature fluctuations and supporting life.

Example: Coastal areas experience milder climates due to water’s high specific heat moderating air temperatures.

Evaporative Cooling

• Evaporation: Transformation from liquid to gas.

• Heat of vaporization: Heat required for 1 g of liquid to become gas.

• Evaporative cooling occurs as the hottest molecules leave as gas, cooling the remaining liquid.

Example: Sweating in mammals cools the body through evaporative cooling.

Expansion Upon Freezing

• Water is less dense as a solid (ice) than as a liquid due to hydrogen bonds forming a crystalline lattice at 0°C.

• Ice floats on liquid water, insulating aquatic life below and preventing bodies of water from freezing solid.

Example: Arctic sea ice provides a habitat for organisms like ringed seals (Phoca hispida).

Versatility as a Solvent

• Solution: Homogeneous mixture of substances.

• Solvent: Dissolving agent; solute: Substance dissolved.

• Aqueous solution: Water is the solvent.

• Water’s polarity allows it to dissolve ionic compounds (forming hydration shells) and many polar molecules, including proteins.

Example: Table salt (NaCl) dissolves in water as Na+ and Cl− ions become surrounded by water molecules.

Hydrophilic and Hydrophobic Substances

• Hydrophilic: Affinity for water (e.g., salts, sugars).

• Hydrophobic: Repels water (e.g., oils, major components of cell membranes).

Acidic and Basic Conditions Affect Living Organisms

Dissociation of Water

• Hydrogen atoms in water can shift between molecules, forming hydronium ions ($\mathrm{H_3O^+}$) and hydroxide ions ($\mathrm{OH^-}$).

• Pure water is in dynamic equilibrium, with equal concentrations of $\mathrm{H^+}$ and $\mathrm{OH^-}$.

Acids, Bases, and the pH Scale

• Acid: Increases $\mathrm{H^+}$ concentration in solution.

• Base: Reduces $\mathrm{H^+}$ concentration.

• Strong acids/bases dissociate completely; weak acids/bases dissociate partially and reversibly.

• pH scale: $\mathrm{pH} = -\log [\mathrm{H^+}]$

• Neutral solution: $[\mathrm{H^+}] = 10^{-7}\,\text{M}$, $\mathrm{pH} = 7$

• Acidic: $\mathrm{pH} < 7$; Basic: $\mathrm{pH} > 7$

Table: Properties of Acids and Bases

Property	Acids	Bases
Effect on $[\mathrm{H^+}]$	Increase	Decrease
pH	< 7	> 7
Common Example	HCl	NaOH

Buffers

• Buffers are substances that minimize changes in $[\mathrm{H^+}]$ and $[\mathrm{OH^-}]$ in a solution.

• They typically consist of a weak acid and its conjugate base, which reversibly bind $\mathrm{H^+}$ ions.

Example: The bicarbonate buffer system in blood helps maintain pH near 7.4.

Acidification and Environmental Impact

• Human activities, such as burning fossil fuels, increase atmospheric CO2, a portion of which dissolves in oceans, forming carbonic acid and lowering pH (ocean acidification).

• Ocean acidification reduces carbonate ion concentration, affecting marine organisms that rely on calcium carbonate for shells and skeletons (e.g., corals).

Example: Coral reefs are threatened by ocean acidification, which impairs their ability to build skeletons and maintain ecosystem health.