Water: Properties, Chemistry, and Biological Importance

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Water is a polar molecule, meaning it has regions of partial positive and partial negative charge due to its molecular structure.

Polarity: The oxygen atom in a water molecule is more electronegative than hydrogen, causing electrons to be shared unequally. This results in a partial negative charge near the oxygen and partial positive charges near the hydrogens.
Polar Covalent Bonds: The bonds between oxygen and hydrogen in water are polar covalent, contributing to the molecule's overall polarity.
Hydrogen Bonding: The polarity of water allows it to form hydrogen bonds with other water molecules or polar substances.
Example: Water's polarity enables it to dissolve many ionic and polar substances, making it an excellent solvent.

Water exhibits several unique properties due to its molecular structure and hydrogen bonding.

Cohesion: Water molecules stick to each other due to hydrogen bonding.
Adhesion: Water molecules stick to other surfaces, aiding processes like capillary action.
Surface Tension: The cohesive forces at the surface of water create a 'skin' that resists external force.
Moderation of Temperature: Water can absorb or release large amounts of heat with only slight changes in its own temperature.
Expansion Upon Freezing: Water becomes less dense as it freezes, causing ice to float.

These properties are essential for many biological processes, such as water transport in plants.

Cohesion: Attraction between water molecules due to hydrogen bonding.
Adhesion: Attraction between water molecules and other substances.
Surface Tension: The measure of how difficult it is to stretch or break the surface of a liquid.
Example: Water droplets forming on a leaf due to cohesion and adhesion.

Water helps regulate temperature in organisms and environments due to its high specific heat and heat of vaporization.

Heat: Total kinetic energy due to molecular motion.
Temperature: Average kinetic energy of molecules.
Calorie: Amount of heat required to raise the temperature of 1 g of water by 1°C.
Specific Heat: The amount of heat needed to change the temperature of a substance. For water, this is relatively high ().
Heat of Vaporization: The amount of heat required to convert water from liquid to gas.
Evaporative Cooling: As water evaporates, it removes heat, cooling the surface.
Example: Sweating cools the human body through evaporative cooling.

Water expands as it freezes, making ice less dense than liquid water.

Hydrogen Bonds: In ice, water molecules form a crystalline structure stabilized by hydrogen bonds, spacing molecules farther apart.
Importance: Ice floats on water, insulating aquatic life in cold climates.
Temperature of Freezing: Water freezes at 0°C (32°F).
Example: Lakes freeze from the top down, allowing organisms to survive below the ice.

Water is known as the "universal solvent" due to its ability to dissolve many substances.

Solute: The substance dissolved in a solution.
Solvent: The substance that dissolves the solute (water in aqueous solutions).
Aqueous Solution: A solution in which water is the solvent.
Hydrophilic: Substances that interact well with water (e.g., salts, sugars).
Hydrophobic: Substances that do not interact well with water (e.g., oils, fats).
Colloid: A stable suspension of fine particles in a liquid.
Hydration Shell: Water molecules surround and isolate ions or polar molecules in solution.
Molar Mass: The mass of one mole of a substance (g/mol).
Molarity: Concentration of solute in solution ().
Example: Table salt (NaCl) dissolving in water forms Na+ and Cl- ions surrounded by hydration shells.

Acids and bases affect the concentration of hydrogen ions in aqueous solutions, which is measured by pH.

Acid: Substance that increases the hydrogen ion concentration in a solution.
Base: Substance that reduces the hydrogen ion concentration.
pH Scale: Measures the concentration of hydrogen ions ().
Neutral Solution: pH = 7; [H+] = [OH-]
Acidic Solution: pH < 7; [H+] > [OH-]
Basic Solution: pH > 7; [H+] < [OH-]
Example: Lemon juice is acidic (pH ≈ 2), while household ammonia is basic (pH ≈ 11).

Buffers help maintain stable pH in biological systems by neutralizing excess acids or bases.

Buffer: A substance that minimizes changes in pH by accepting or donating H+ ions.
Bicarbonate Buffer System: Important in blood and cellular fluids. Involves the equilibrium:

Function: Helps maintain blood pH around 7.4.
Relationship: CO2 and bicarbonate are interconverted to buffer changes in pH.
Example: During exercise, increased CO2 production is buffered by the bicarbonate system.

Property	Description	Biological Importance
Cohesion	Water molecules stick together	Enables transport in plants
Adhesion	Water molecules stick to other surfaces	Assists capillary action
Surface Tension	Resistance to breaking the surface	Allows small organisms to walk on water
High Specific Heat	Absorbs/release heat with little temperature change	Stabilizes climate and body temperature
Expansion Upon Freezing	Ice is less dense than liquid water	Insulates aquatic life in winter
Versatile Solvent	Dissolves many substances	Facilitates biochemical reactions