Properties of Water and Their Biological Significance

Polar Covalent Bonds and Hydrogen Bonding

Water molecules are held together by polar covalent bonds, resulting in an uneven distribution of charge and the formation of a polar molecule. This polarity allows water molecules to form hydrogen bonds with each other, which are crucial for many of water's unique properties.

• Polar covalent bond: A type of chemical bond where electrons are shared unequally between atoms, resulting in partial charges.

• Hydrogen bond: A weak bond between a hydrogen atom in one molecule and an electronegative atom (such as oxygen) in another molecule.

• Example: Water (H2O) molecules form hydrogen bonds, leading to high cohesion and surface tension.

Emergent Properties of Water

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

• Cohesive and adhesive behavior

• Ability to moderate temperature

• Expansion upon freezing

• Versatility as a solvent

Cohesion and Adhesion

Cohesion refers to the attraction between water molecules due to hydrogen bonding, resulting in high surface tension. Adhesion is the attraction between water molecules and other substances, such as plant cell walls.

• Surface tension: The measure of how difficult it is to break the surface of a liquid.

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

Moderation of Temperature

Water can absorb and release large amounts of heat with only slight changes in its own temperature, due to its high specific heat.

• Specific heat: The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Equation:

• Evaporative cooling: As water evaporates, it removes heat from surfaces, helping to regulate temperature in organisms and environments.

• Example: Sweating in humans cools the body through evaporative cooling.

Expansion Upon Freezing

Water is less dense as a solid than as a liquid, which is why ice floats on liquid water. This is due to the hydrogen bonds holding water molecules farther apart in the solid state.

• Example: Ice floating on lakes insulates the water below, protecting aquatic life during winter.

Versatility as a Solvent

Water is known as the "universal solvent" due to its ability to dissolve a wide variety of substances, especially those that are polar or charged.

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

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

• Solute: The substance that is dissolved.

• Aqueous solution: A solution in which water is the solvent.

• Hydration shell: The sphere of water molecules surrounding an ion or polar molecule in solution.

• Hydrophilic substances: Have an affinity for water.

• Hydrophobic substances: Do not have an affinity for water; often nonpolar molecules like oils.

• Example: Salt (NaCl) dissolving in water forms an aqueous solution.

Solute Concentration in Aqueous Solutions

Concentration of solutes in water is important for chemical reactions in organisms. It is commonly measured in molarity (M), which is moles of solute per liter of solution.

• Molecular mass: The sum of all atomic masses in a molecule.

• Mole: molecules

• Molarity (M):

Acidic and Basic Conditions

Water can dissociate into hydrogen ions (H+) and hydroxide ions (OH-), affecting the pH of solutions. The balance of these ions is crucial for biological systems.

• Dissociation: The process by which water molecules split into H+ and OH-.

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

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

• Salt: A compound formed when an acid reacts with a base.

• Strong acids and bases: Dissociate completely in water.

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

pH Scale

The pH of a solution is defined by the negative logarithm of the H+ concentration:

• Equation:

• For neutral solutions, , so

• The pH scale ranges from 0 (very acidic) to 14 (very basic).

• Most biological fluids have pH values between 6 and 8.

Buffers

Buffers are substances that minimize changes in concentrations of H+ and OH- in a solution, helping to maintain stable pH in biological systems.

• Buffers contain a weak acid and its corresponding base, which combine reversibly with H+ ions.

• Example: The buffer system in human blood uses carbonic acid and bicarbonate to maintain pH.

• If blood becomes too basic, carbonic acid releases H+ to combine with OH-.

• If blood becomes too acidic, bicarbonate accepts H+ to make carbonic acid.

Summary Table: Four Emergent Properties of Water

Additional info: Academic context and definitions have been expanded for clarity and completeness.