Water and Life: Properties, Chemistry, and Biological Importance

Chapter 3: Water and Life

Introduction

Water is a fundamental molecule for all known forms of life. Its unique chemical and physical properties make it essential for biological processes and the maintenance of life on Earth. This chapter explores the molecular structure of water, its emergent properties, and its role as a solvent, as well as the concepts of acids, bases, pH, and the environmental issue of ocean acidification.

Polar Covalent Bonds and Hydrogen Bonding in Water

Structure and Polarity of Water

• Water Molecule: Composed of two hydrogen atoms covalently bonded to one oxygen atom (H2O).

• Polar Covalent Bonds: The electrons in the O-H bonds spend more time near the oxygen atom, making it partially negative and the hydrogens partially positive.

• 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, which are weak individually but strong collectively.

Definition: Electronegativity is the tendency of an atom to attract electrons in a covalent bond.

Four Emergent Properties of Water

1. Cohesive Behavior

• 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 plant cell walls, helping counteract gravity in plants.

• Example: Water moving up plant stems (capillary action).

2. Ability to Moderate Temperature

• High Specific Heat: Water can absorb or release a large amount of heat with only a slight change in its own temperature.

• Specific Heat: The amount of heat required to raise 1 gram of a substance by 1°C. For water, it is 1 cal/(g·°C).

• Heat Absorption and Release: Heat is absorbed when hydrogen bonds break and released when they form.

• Biological Importance: Stabilizes ocean and organism temperatures, moderating climate and internal body temperatures.

Definition: Kinetic energy is the energy of motion; thermal energy is the kinetic energy associated with the random movement of atoms or molecules.

3. Expansion Upon Freezing

• Ice Floats: Water is less dense as a solid than as a liquid because hydrogen bonds form a crystalline lattice that spaces molecules apart.

• Biological Importance: If ice sank, bodies of water would freeze solid, making life impossible in aquatic environments.

4. Versatility as a Solvent

• Solution: A homogeneous mixture of substances.

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

• Solute: The substance that is dissolved.

• Hydration Shell: When ionic compounds dissolve, each ion is surrounded by water molecules.

• Water dissolves: Ionic and polar (hydrophilic) substances, and large molecules with ionic/polar regions.

• Water does not dissolve: Non-polar (hydrophobic) substances, such as oils.

Hydrophilic vs. Hydrophobic Substances

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

• Hydrophobic: Substances that repel water (usually non-polar), such as oils and many cell membrane components.

Solute Concentration and Molarity

Calculating Molarity

• Mole (mol): 6.02 × 1023 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 Water

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

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

• pH Scale: Measures the concentration of H+ ions; ranges from 0 (most acidic) to 14 (most basic).

• Neutral Solution: [H+] = [OH-], pH = 7.

• Strong acids/bases: Dissociate completely in water.

• Weak acids/bases: Reversibly release and accept H+ ions.

Formulas:

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

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

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

• Example: A solution with [H+] = 0.0235 M has pH = -log(0.0235) ≈ 1.63.

Buffers

• Definition: Buffers are substances that minimize changes in concentrations of H+ and OH- in a solution.

• Composition: Usually consist of a weak acid and its corresponding base.

• Function: Maintain stable pH in biological systems, crucial for cellular processes.

• Example: The bicarbonate buffer system in blood: $\mathrm{HCO_3^- + H^+ \leftrightarrow H_2CO_3}$

Ocean Acidification

• Definition: The process by which excess atmospheric CO2 dissolves in oceans, forming carbonic acid and lowering pH.

• Impact: Increased H+ ions combine with carbonate ions, reducing carbonate available for marine organisms to build shells and skeletons.

• Environmental Concern: Threatens marine biodiversity and ecosystem stability.

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 change	Stabilizes climate and organism temperature
Expansion upon Freezing	Ice is less dense than liquid water	Prevents bodies of water from freezing solid
Versatility as a Solvent	Dissolves many substances	Facilitates chemical reactions in cells

Key Terms

• Polar molecule

• Hydrogen bond

• Cohesion

• Adhesion

• Specific heat

• Solvent

• Solute

• Hydrophilic

• Hydrophobic

• Molarity

• Acid

• Base

• pH

• Buffer

• Ocean acidification