Water and Life

Introduction

Water is fundamental to all known forms of life. Its unique chemical and physical properties make it essential for biological processes and the habitability of Earth. This chapter explores the molecular structure of water, its emergent properties, and its critical roles in living organisms.

The Molecular Structure of Water

Polarity and Hydrogen Bonding

Water (H2O) is a polar molecule due to the high electronegativity of oxygen, which causes an uneven distribution of charge within the molecule. This polarity enables water molecules to form hydrogen bonds with each other.

• Polar Covalent Bonds: The oxygen atom attracts electrons more strongly than hydrogen, resulting in a partial negative charge near oxygen and partial positive charges near hydrogens.

• Hydrogen Bonds: The positive region of one water molecule is attracted to the negative region of another, forming a hydrogen bond. These are weak individually but strong collectively.

• Example: Each water molecule can form up to four hydrogen bonds with neighboring water molecules.

Emergent Properties of Water

Overview

Water exhibits several emergent properties that are critical for life on Earth. These properties arise from the collective behavior of water molecules and their ability to form hydrogen bonds.

• Cohesive Behavior

• Ability to Moderate Temperature

• Expansion Upon Freezing

• Versatility as a Solvent

Cohesive Behavior

Cohesion refers to the attraction between water molecules due to hydrogen bonding. This property enables water to move against gravity in plants (capillary action) and contributes to surface tension.

• Surface Tension: Water has a high surface tension, allowing small insects to walk on its surface.

• Example: Water transport in plants relies on cohesion and adhesion (attraction to other substances).

Ability to Moderate 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 capacity.

• Specific Heat: The amount of heat required to raise the temperature of 1 gram of water by 1°C is high, which stabilizes temperatures in organisms and environments.

• Heat of Vaporization: Water requires a large amount of energy to change from liquid to gas, which helps cool organisms through evaporation (e.g., sweating).

• Example: Coastal regions have milder climates due to the moderating effect of large bodies of water.

Expansion Upon Freezing

Unlike most substances, water expands when it freezes. This is because hydrogen bonds arrange water molecules into a crystalline lattice that is less dense than liquid water.

• Ice Floats: The lower density of ice allows it to float on liquid water, insulating aquatic life in winter.

• Example: Lakes and ponds freeze from the top down, preserving life below the ice.

Versatility as a Solvent

Water is known as the "universal solvent" because it can dissolve a wide variety of substances, especially ionic and polar compounds.

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

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

• Solute: The substance that is dissolved.

• Hydration Shell: When an ionic compound dissolves, water molecules surround each ion, separating and shielding them from one another.

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

Water in Biological Systems

Role in Cells and Organisms

Water makes up 70-95% of the mass of most cells and is the medium for most biochemical reactions.

• Hydrophilic Substances: Substances with an affinity for water (e.g., salts, sugars, proteins with polar regions).

• Hydrophobic Substances: Substances that do not interact well with water (e.g., oils, fats), often forming the basis of cell membranes.

• Example: Protein folding is influenced by the hydrophobic effect, where nonpolar regions avoid water.

Acids, Bases, and pH

Acidic and Basic Conditions

The concentration of hydrogen ions (H+) and hydroxide ions (OH-) determines the acidity or basicity of a solution. Even small changes in these concentrations can affect cellular processes.

• Acid: A substance that increases the H+ concentration of a solution.

• Base: A substance that reduces the H+ concentration, often by accepting H+ or releasing OH-.

• Strong Acids/Bases: Dissociate completely in water.

• Weak Acids/Bases: Reversibly release and accept back hydrogen ions.

The pH Scale

The pH scale measures the concentration of H+ ions in a solution, ranging from 0 (most acidic) to 14 (most basic).

• Formula:

• Neutral Solution: [H+] = 1 × 10-7 M, so pH = 7.

• Biological Fluids: Most have pH values between 6 and 8.

Buffers

Buffers are substances that minimize changes in pH by accepting or donating H+ ions. They are crucial for maintaining stable internal conditions in living organisms.

• Example: The bicarbonate buffer system in blood:

• When pH drops (more acidic), the buffer absorbs excess H+.

• When pH rises (more basic), the buffer releases H+.

Acidification and Environmental Impact

Ocean Acidification

Human activities, such as burning fossil fuels, increase atmospheric CO2, a portion of which is absorbed by oceans. This leads to the formation of carbonic acid, lowering ocean pH in a process called ocean acidification.

• Impact: Increased H+ ions combine with carbonate ions, reducing the availability of carbonate for marine organisms to build shells and skeletons (e.g., corals, mollusks).

• Example: Ocean acidification threatens marine biodiversity and disrupts chemical balances in aquatic ecosystems.

Summary Table: Properties of Water and Their Biological Importance