chapter3

Chapter 3: Water and Life

The Molecule That Supports All of Life

  • Water enables life on Earth.

  • It is the only common substance in the natural environment to exist in all three physical states (solid, liquid, gas).

  • Unique emergent properties of water contribute to its suitability for life.

  • Structure of the water molecule promotes interaction with other molecules.

Concept 3.1: Polar Covalent Bonds in Water Molecules Result in Hydrogen Bonding

  • In the water molecule (H₂O), electrons in polar covalent bonds favor oxygen over hydrogen, as oxygen is more electronegative than hydrogen.

  • This leads to uneven distribution of charge, making water a polar molecule.

  • The polarity allows water molecules to form hydrogen bonds with one another.

Figure 3.2: Hydrogen Bonds Between Water Molecules

  • Illustrated depiction demonstrating the polarity within water molecules and the resulting hydrogen bonds.

  • Hydrogen atoms ($H^{+}$) show partial positive charges and oxygen atoms ($O^{2-}$) exhibit partial negative charges, leading to the formation of hydrogen bonds (illustrated as dashes in the figure).

Concept 3.2: Four Emergent Properties of Water Contribute to Earth’s Suitability for Life

  1. Cohesive Behavior: Water molecules are attracted to each other due to hydrogen bonding.

  2. Ability to Moderate Temperature: Water has a high specific heat capacity.

  3. Expansion Upon Freezing: Ice is less dense than liquid water, allowing it to float.

  4. Versatility as a Solvent: Water can dissolve a wide variety of substances.

Cohesion of Water Molecules

  • Cohesion refers to the attraction between like molecules, facilitated by hydrogen bonds.

  • It assists in the upward transport of water against gravity in plants through xylem and leads to phenomena such as surface tension.

  • Adhesion: Attraction between different substances, e.g., water and plant cell walls.

Figure 3.3: Water Transport in Plants

  • Shows how evaporation pulls water upward in plants through cohesion and adhesion, detailing water-conducting cells with a scale of 300 μm.

Temperature and Heat

  • Kinetic Energy: Energy due to motion.

  • Thermal Energy: The kinetic energy associated with atoms/molecules in motion.

  • Temperature: Average kinetic energy of molecules in a substance.

  • Heat: Thermal energy in transit between bodies of matter.

Caloric Measurement

  • A calorie (cal) is defined as the heat needed to raise 1 g of water by 1°C, or the heat released when it cools by the same amount.

  • "Calories" on food packages are actually kilocalories (kcal); where 1 kcal = 1,000 cal.

  • Joule (J): Another unit of energy defined as 1 J = 0.239 cal or 1 cal = 4.184 J.

Moderation of Temperature by Water

  • Water can absorb heat from warmer environments and release it to cooler environments.

  • Water can absorb/release significant amounts of heat with minimal change in its own temperature.

Water’s High Specific Heat

  • Specific heat is defined as the heat required to change the temperature of 1 g of a substance by 1°C.

  • Water's specific heat is $1 ext{ cal}/( ext{g} imes ext{°C})$.

  • Water resists temperature changes due to its high specific heat, traced to hydrogen bonding.

  • Heat is absorbed when hydrogen bonds break and released when they form, minimizing temperature fluctuations that permit life.

Evaporative Cooling

  • Evaporation: Transformation from liquid to gas.

  • Heat of Vaporization: The amount of heat required to convert 1 g of a liquid to gas.

  • Evaporation of a liquid cools the remaining surface, termed evaporative cooling, which is crucial for temperature regulation in organisms and aquatic environments.

Floating of Ice on Liquid Water

  • Ice floats due to the stable hydrogen bonds in its crystalline structure, making it less dense than liquid water.

  • Water achieves maximum density at 4°C; if ice sank, aquatic life would be endangered as bodies of water would freeze solid.

Figure 3.6: Ice Properties

  • Displays the differences in molecular structures and densities between ice and liquid water.

Climate Change Concerns

  • Scientists express concern over global warming's impact on polar environments, affecting the habitat of numerous species that depend on ice.

Water: The Solvent of Life

  • A solution is a homogeneous mixture where the solvent dissolves the solute; in most biological contexts, this solvent is water, leading to aqueous solutions.

Water’s Solvent Properties

  • Water's polarity renders it a versatile solvent.

  • When ionic substances (e.g. table salt) dissolve, ions are surrounded by water molecules forming a hydration shell.

Dissolution of Nonionic Polar Molecules

  • Water can also dissolve polar molecules and large polar proteins if they possess ionic/polar regions.

Hydrophilic and Hydrophobic Substances

  • Hydrophilic: Substances that are attracted to water (e.g., salts and sugars).

  • Hydrophobic: Substances that do not mix with water, like oil molecules, primarily due to their non-polar nature; these are essential components of cellular membranes.

Solute Concentration in Aqueous Solutions

  • Chemical reactions in biological systems typically involve solutes in aqueous solutions.

  • Mass is used to quantify solute concentrations.

Molecular Mass and Molarity

  • Molecular Mass: The collective mass of all atoms in a molecule.

  • Moles: A quantitative measure based on Avogadro's number (6.022 x 10²³ molecules/mole).

  • Molarity (M): Expressed as number of moles of solute per liter of solution.

Possible Evolution of Life on Other Planets

  • Search for extraterrestrial life focuses on planets where water exists.

  • Over 800 exoplanets have been discovered, some showing evidence of water vapor; Mars has shown similar findings.

Concept 3.3: Acidic and Basic Conditions Affect Living Organisms

  • A hydrogen atom can shift between water molecules, leading to proton transfer.

  • The molecule losing a proton becomes a hydroxide ion ($OH^-$) while the one gaining becomes a hydronium ion ($H₃O^+$), often simplified as $H^+$.

Dynamic Equilibrium of Water

  • Water reaches a dynamic equilibrium where dissociation and recombination of water molecules occur at equal rates, maintaining homeostasis.

Effects of Ion Concentration Changes

  • Although dissociation events are statistically infrequent, their implications on cellular chemistry can be significant due to alterations in $H^+$ and $OH^-$ concentrations.

pH Scale and Biological Relevance

  • Alterations in $H^+$ and $OH^-$ levels, induced by solute introduction (acids and bases), can shift chemical balances.

  • The pH scale describes acidity and basicity, with acidic solutions having $[H^+] > [OH^-]$ and basic solutions having $[H^+] < [OH^-]$.

Defining Acids and Bases

  • Acid: Increases $[H^+]$.

  • Base: Decreases $[H^+]$.

  • Strong acids/bases dissociate completely in water, while weak acids/bases can reversibly release/accept $H^+$, affecting pH similarly.

Figure 3.UN04: pH Scale Overview

  • Demonstrates the pH scale, identifying acidic (0-6), neutral (7), and basic (8-14) regions.

Constants in Aqueous Solutions

  • The product of $[H^+]$ and $[OH^-]$ remains constant at 25°C, described mathematically.

  • pH is calculated as the negative logarithm of hydrogen ion concentration: $pH = - ext{log}([H^+])$; for neutral solutions, $pH = 7$.