Water and Life (Chapter 3): Properties, Structure, and Biological Importance

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Water is essential for all known forms of life. Its unique chemical and physical properties make Earth habitable and support biological processes.

Water is the only common substance that exists naturally in all three physical states: solid, liquid, and gas.
Water's emergent properties (arising from its molecular structure) are crucial for life on Earth.
The structure of the water molecule enables it to interact with other molecules, facilitating various biological functions.

A water molecule (H2O) consists of two hydrogen atoms covalently bonded to one oxygen atom.
Electrons spend more time near the oxygen atom, making the bond polar covalent.
This results in a polar molecule with partial negative charge (δ−) near oxygen and partial positive charges (δ+) near hydrogens.
Polarity allows water molecules to form hydrogen bonds with each other.

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

Water's structure gives rise to four key properties that support life:

Cohesion: Hydrogen bonds hold water molecules together, allowing for the transport of water against gravity in plants.
Adhesion: Attraction between different substances, such as water and plant cell walls, aids in water movement.
Surface tension: Water has high surface tension due to hydrogen bonding, making it difficult to break the surface (e.g., water striders walking on water).

Example: Evaporation pulls water upward in plants, with cohesion and adhesion facilitating this movement.

Water absorbs heat from warmer air and releases stored heat to cooler air, stabilizing temperatures in organisms and environments.
Water can absorb or release large amounts of heat with only slight changes in its own temperature due to its high specific heat.
Hydrogen bonding is responsible: heat is absorbed when bonds break and released when bonds form.

Example: Coastal climates are milder due to water's ability to moderate temperature.

Ice floats because hydrogen bonds in ice are more "ordered," making ice less dense than liquid water.
Water reaches its greatest density at 4°C.
If ice sank, bodies of water would freeze solid, making life impossible.

Example: Floating ice insulates water below, enabling aquatic life to survive in winter.

Solution: Homogeneous mixture of substances.
Solvent: Dissolving agent (e.g., water).
Solute: Substance dissolved (e.g., salt, sugar).
Aqueous solution: Solution where water is the solvent.
Water's polarity makes it a versatile solvent, able to dissolve ionic and polar substances.

Example: Table salt (NaCl) dissolves in water as Na+ and Cl− ions are surrounded by water molecules.

Hydrophilic: Substances with an affinity for water (e.g., salts, acids, bases, carbohydrates).
Hydrophobic: Substances that do not interact with water (e.g., lipids, hydrocarbons).

Molecular mass: Sum of all atomic masses in a molecule.
Number of molecules is measured in moles (mol), where 1 mol = molecules (Avogadro's number).
Molarity (M): Number of moles of solute per liter of solution.

A hydrogen atom in a hydrogen bond between two water molecules can shift, leaving its electron behind and becoming a hydrogen ion (H+).
The molecule that lost the proton becomes a hydroxide ion (OH−).
The molecule with the extra proton becomes a hydronium ion (H3O+), often represented as H+.

Example: HCl is a strong acid; NaOH is a strong base.

pH Value	Example
1-2	Battery acid, gastric juice
3-4	Vinegar, cola
7	Pure water, blood
12-13	Household bleach, oven cleaner

Starting with a solution of pH 8, increasing [H+] by a factor of 1,000 results in a new pH of 5.
Starting with a solution of pH 7, decreasing [H+] by a factor of 100 results in a new pH of 9.

Additional info: Each pH unit represents a tenfold difference in H+ concentration.