Water and Life: Properties and Importance of Water in Biology

Water and Life

Introduction

Water is a fundamental molecule that supports all known forms of life on Earth. Its unique chemical and physical properties make it essential for biological processes and the maintenance of life. This chapter explores the structure of water, its emergent properties, and its critical roles in biological systems.

The Molecule That Supports All of Life

Importance of Water

• Water makes life possible on Earth by providing a medium for chemical reactions and supporting cellular structure.

• Water is the only common substance that exists naturally in all three physical states (solid, liquid, gas) in Earth's environment.

• Emergent properties of water contribute to Earth's suitability for life, including cohesion, temperature moderation, expansion upon freezing, and versatility as a solvent.

• The structure of the water molecule allows it to interact with other molecules, facilitating a wide range of biological functions.

Structure and Polarity of Water

Polar Covalent Bonds and Hydrogen Bonding

The water molecule (H2O) consists of two hydrogen atoms covalently bonded to one oxygen atom. The oxygen atom is more electronegative, pulling shared electrons closer and creating a partial negative charge (δ-) at the oxygen end and partial positive charges (δ+) at the hydrogen ends.

• Polar molecule: A molecule with regions of partial positive and negative charge due to unequal sharing of electrons.

• Hydrogen bonds: Weak attractions between the oppositely charged regions of water molecules. These bonds are responsible for many of water's unique properties.

Example: The polarity of water allows it to form hydrogen bonds with other water molecules and with other polar substances.

Emergent Properties of Water

Overview

Water exhibits four key emergent properties that are critical for life:

• Cohesive behavior

• Ability to moderate temperature

• Expansion upon freezing

• Versatility as a solvent

Cohesion and Adhesion

• Cohesion: The attraction between water molecules due to hydrogen bonding. This property helps water move upward against gravity in plants (capillary action).

• Adhesion: The attraction between water molecules and other substances, such as plant cell walls, aiding in water transport.

• Surface tension: A measure of how difficult it is to break the surface of a liquid. Water has a high surface tension due to hydrogen bonding at the air-water interface.

Example: Water droplets form beads on a surface due to high surface tension.

Moderation of Temperature

• Water can absorb or release large amounts of heat with only a slight change in its own temperature, due to its high specific heat.

• Specific heat: The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Hydrogen bonds absorb heat when they break and release heat when they form, buffering temperature changes.

Example: Coastal areas have milder climates because water moderates temperature fluctuations.

Expansion Upon Freezing

• In liquid water, molecules are close together, but in ice, hydrogen bonds stabilize and keep molecules farther apart.

• This makes ice less dense than liquid water, allowing it to float.

• Ecological significance: Floating ice insulates the water below, enabling aquatic life to survive in cold climates.

Example: Lakes do not freeze solid in winter, preserving aquatic ecosystems.

Versatility as a Solvent

• Solution: A homogeneous mixture of two or more substances.

• Solvent: The dissolving agent (water in aqueous solutions).

• Solute: The substance that is dissolved.

• Water's polarity allows it to dissolve many ionic and polar substances, making it a versatile solvent.

Example: Table salt (NaCl) dissolves in water as the polar water molecules surround and separate the ions.

Hydrophilic and Hydrophobic Substances

Definitions and Examples

• Hydrophilic substances: Have an affinity for water (e.g., ionic compounds, polar molecules, salts, acids, bases, carbohydrates).

• Hydrophobic substances: Do not have an affinity for water (e.g., nonpolar substances such as lipids and hydrocarbons).

Example: Oil does not mix with water because it is hydrophobic.

Concentration in Aqueous Solutions

Molecular Mass and Molarity

• Molecular mass: The sum of the masses of all atoms in a molecule.

• Mole (mol): A unit representing 6.02 × 1023 molecules (Avogadro's number).

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

Formula:

$\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{liters of solution}}$

Acidic and Basic Conditions Affect Living Organisms

Acids, Bases, and the pH Scale

• Acid: A substance that increases the hydrogen ion (H+) concentration of a solution.

• Base: A substance that reduces the H+ concentration, often by increasing hydroxide ions (OH-).

• In pure water, the concentrations of H+ and OH- are equal.

• The product of [H+] and [OH-] in any aqueous solution is constant:

$[H^+][OH^-] = 10^{-14}$

• pH: Defined as the negative logarithm of the hydrogen ion concentration:

$pH = -\log [H^+]$

• In a neutral solution, [H+] = 10-7 M, so pH = 7.

pH Scale and Examples

Substance	pH Value	Acidic/Basic
Gastric juice (stomach acid)	1-2	Acidic
Vinegar, wine, cola	3	Acidic
Tomatoes	4	Acidic
Pure water	7	Neutral
Human blood, tears	7.4	Slightly basic
Household ammonia	11-12	Basic
Household bleach	13	Basic

pH Calculations and Changes

• Each unit change in pH represents a tenfold change in [H+].

• For example, increasing [H+] by a factor of 1,000 lowers the pH by 3 units.

• Decreasing [H+] by a factor of 100 raises the pH by 2 units.

Example: If a solution starts at pH 8 and [H+] increases by 1,000 times, the new pH is 5.

Example: If a solution starts at pH 7 and [H+] increases by 100 times, the new pH is 5.

Additional info: Some explanations and examples were expanded for clarity and completeness, including the ecological significance of ice floating, the definition of specific heat, and the pH calculation examples.