Water: Its Physical Properties and Biological Importance

Introduction to Water in Biology

Water is fundamental to all known forms of life. Its unique physical and chemical properties make it indispensable for biological processes and the habitability of Earth.

• Biological reactions occur in water, serving as the medium for cellular processes.

• All living organisms require water more than any other substance.

• Most cells are surrounded by water, and cells themselves are composed of approximately 70–95% water.

• The abundance of water is a primary reason Earth is habitable.

Structure and Bonding in Water

Polar Covalent Bonds and Hydrogen Bonding

The water molecule (H2O) is polar, with oxygen being more electronegative than hydrogen. This polarity leads to hydrogen bonding, which is responsible for many of water's unique properties.

• Polarity: Oxygen atom has a partial negative charge, while hydrogen atoms have partial positive charges.

• Hydrogen bonds: Weak attractions between the partially positive hydrogen of one water molecule and the partially negative oxygen of another.

Properties of Water Conducive to Life

Temperature Moderation

Water has a high specific heat, meaning it resists changes in temperature. This property helps stabilize climates and internal temperatures of organisms.

• Specific heat of water:

• Hydrogen bonding causes water to absorb heat when bonds break and release heat when bonds form.

• Minimizes temperature fluctuations within limits that permit life.

Evaporative Cooling

Evaporation of water removes heat from surfaces, helping organisms and environments maintain stable temperatures.

• Heat of vaporization: The amount of energy required for water to change from liquid to gas.

• As water evaporates, the remaining surface cools (evaporative cooling).

• Helps stabilize temperatures in organisms and bodies of water.

Expansion Upon Freezing

Water expands as it freezes, making ice less dense than liquid water. This property is crucial for aquatic life.

• Hydrogen bonds in ice are more ordered, causing ice to float.

• Water reaches its greatest density at 4°C.

• If ice sank, bodies of water would freeze solid, making life impossible.

Cohesive Behavior

Water molecules stick together due to hydrogen bonding, resulting in high surface tension and capillary action.

• Cohesion: Attraction between water molecules.

• Adhesion: Attraction between water molecules and other substances.

Versatile Solvent

Water is known as the "universal solvent" because it readily dissolves many substances, facilitating chemical reactions in cells.

• Solution: Homogeneous mixture of two or more substances.

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

• Solute: The substance dissolved.

• Water forms hydrogen bonds with charged and polar molecules.

Table: Types of Substances in Water

Solute Concentration in Aqueous Solutions

Moles and Molarity

Chemical reactions in cells depend on the concentration of solutes, which is commonly expressed in moles and molarity.

• Mole (mol): molecules (Avogadro's number).

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

Dissociation of Water and pH

Water Dissociation

Water molecules can dissociate into ions, affecting the pH of solutions.

• Hydrogen ion (H+): Proton transferred from one water molecule to another.

• Hydroxide ion (OH-): Water molecule that lost a proton.

• Hydronium ion (H3O+): Water molecule with an extra proton.

• At equilibrium, only a small fraction of water molecules dissociate.

Acids and Bases

Acids and bases alter the concentration of hydrogen ions in a solution.

• Acid: Substance that increases H+ concentration.

• Base: Substance that reduces H+ concentration, either directly or indirectly.

• Strong acids/bases: Dissociate completely in water (e.g., HCl, NaOH).

• Weak acids/bases: Dissociate partially (e.g., NH3).

pH Scale

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

• Formula:

• Pure water: M, pH = 7 (neutral)

• Acidic solutions: pH < 7

• Basic solutions: pH > 7

Table: pH Values of Common Substances

Buffers and Regulation of pH

Buffers

Buffers help maintain stable pH in biological systems by accepting or donating hydrogen ions.

• Buffer: Weak acid and its corresponding base.

• Carbonic acid-bicarbonate buffer system: Regulates pH in blood and other biological fluids.

• Chemical equilibrium shifts to counteract changes in pH.

Carbonic Acid-Bicarbonate Buffer System

• Blood pH is maintained between 7.3 and 7.5.

• Reaction:

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

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

Environmental Impact: Ocean Acidification

Ocean Acidification

Human activities, such as burning fossil fuels, increase atmospheric CO2, which dissolves in oceans and forms carbonic acid, lowering ocean pH.

• CO2 + H2O → H2CO3 (carbonic acid)

• Carbonic acid dissociates:

• Increased H+ combines with carbonate ions, reducing carbonate availability for marine organisms.

• Carbonate is required for calcification (formation of calcium carbonate, CaCO3) by corals and other marine life.

Table: Effects of Ocean Acidification

Example: Coral reefs are threatened by ocean acidification, which impairs their ability to build calcium carbonate skeletons.

Additional info: The notes have been expanded to include definitions, examples, and tables for clarity and completeness.