Ch. 3Water and Life: Properties, Chemistry, and Biological Importance

Water and Life

Introduction

Water is a fundamental molecule for life on Earth, exhibiting unique physical and chemical properties that make it essential for biological systems. This chapter explores the molecular structure of water, its emergent properties, and its role in supporting life.

Polar Covalent Bonds and Hydrogen Bonding in Water

Structure of the Water Molecule

• Polar Covalent Bonds: In a water molecule (H2O), the electrons are shared unequally between oxygen and hydrogen atoms, resulting in a partial negative charge near the oxygen and a partial positive charge near the hydrogens.

• Electronegativity: Oxygen is more electronegative than hydrogen, attracting electrons more strongly.

• Polarity: Water is a polar molecule, meaning it has an uneven distribution of charge.

• Hydrogen Bonding: The polarity allows water molecules to form hydrogen bonds with each other, where the partially positive hydrogen of one molecule is attracted to the partially negative oxygen of another.

Example: Hydrogen bonds are responsible for water's high boiling point compared to other molecules of similar size.

Four Emergent Properties of Water

Overview

Water's unique properties arise from its ability to form hydrogen bonds. These properties are crucial for life on Earth.

• Cohesive Behavior

• Ability to Moderate Temperature

• Expansion Upon Freezing

• Versatility as a Solvent

Cohesion and Adhesion

• Cohesion: Hydrogen bonds hold water molecules together, resulting in high surface tension (the difficulty of breaking the surface of a liquid).

• Adhesion: Water molecules can also be attracted to other substances, such as the walls of plant cells, helping counteract gravity during water transport in plants.

Example: Water droplets forming beads on a leaf due to cohesion; water climbing up plant vessels due to adhesion.

Moderation of Temperature

• High Specific Heat: Water can absorb or release a large amount of heat with only a slight change in its own temperature. The specific heat of water is 1 cal/(g·°C).

• Thermal Energy: The kinetic energy associated with the random movement of atoms or molecules.

• Evaporative Cooling: As water evaporates, the surface cools, helping to stabilize temperatures in organisms and environments.

Example: Coastal areas have milder climates due to water's high specific heat; sweating cools the body via evaporative cooling.

Expansion Upon Freezing

• Water is less dense as a solid (ice) than as a liquid because hydrogen bonds form a crystalline lattice that spaces molecules apart.

• This property allows ice to float on liquid water, insulating aquatic life in winter.

Example: Lakes do not freeze solid from the bottom up, allowing aquatic organisms to survive beneath the ice.

Versatility as a Solvent

• Solution: A homogeneous mixture of substances.

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

• Solute: The substance that is dissolved.

• Water dissolves ionic compounds and polar molecules by surrounding them with a hydration shell.

• Nonpolar (hydrophobic) substances, such as oils, do not dissolve in water.

Example: Table salt (NaCl) dissolves in water, but oil does not.

Hydrophilic and Hydrophobic Substances

Definitions and Biological Importance

• Hydrophilic: Substances that have an affinity for water (usually polar or charged).

• Hydrophobic: Substances that repel water (usually nonpolar).

• Hydrophobic molecules, such as those in cell membranes, are crucial for forming biological barriers.

Example: Cell membranes are composed of hydrophobic lipid bilayers.

Solute Concentration and Molarity

Calculating Molarity

• Mole (mol): The amount of substance containing $6.02 \times 10^{23}$ molecules (Avogadro's number).

• Molecular Mass: The sum of atomic masses in a molecule (g/mol).

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

Formula:

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

Example: To make a 1 M NaCl solution, dissolve 58 g of NaCl (molecular mass = 58 g/mol) in 1 L of water.

Acids, Bases, and pH

Acid-Base Chemistry in Aqueous Solutions

• Water can dissociate into hydrogen ions ($H^+$) and hydroxide ions ($OH^-$).

• Acid: Increases the $H^+$ concentration of a solution (pH < 7).

• Base: Reduces the $H^+$ concentration (pH > 7).

• pH Scale: Measures the concentration of $H^+$ in a solution.

Formulas:

$[H^+][OH^-] = 10^{-14}$ (at 25°C)

$\text{pH} = -\log [H^+]$

$\text{pOH} = 14 - \text{pH}$

Example: A solution with $[H^+] = 0.0235$ M has $\text{pH} = -\log(0.0235) = 1.629$.

Buffers

• Buffer: A substance that minimizes changes in $[H^+]$ and $[OH^-]$ in a solution.

• Buffers typically consist of a weak acid and its corresponding base, which combine reversibly with $H^+$ ions.

Example: The bicarbonate buffer system in blood helps maintain a stable pH.

Ocean Acidification

Causes and Consequences

• Excess atmospheric $CO_2$ dissolves in seawater, forming carbonic acid ($H_2CO_3$), a process called ocean acidification.

• Increased $H^+$ ions combine with carbonate ions, reducing the availability of carbonate for marine organisms that build shells and skeletons.

• This threatens marine ecosystems and biodiversity.

Example: Coral reefs are at risk due to decreased carbonate ion concentration, affecting organisms that rely on calcium carbonate structures.

Summary Table: Properties of Water

Property	Description	Biological Importance
Cohesion	Hydrogen bonds hold water molecules together	Enables transport of water in plants
High Specific Heat	Water resists temperature changes	Stabilizes climate and organismal temperature
Expansion Upon Freezing	Ice is less dense than liquid water	Insulates aquatic life in winter
Versatility as a Solvent	Dissolves many substances	Facilitates chemical reactions in cells