Water and Life

Water is considered the substance that makes life possible.
Living organisms are composed predominantly of water.
Water is unique as it is the only common substance that can exist in all three physical states: solid, liquid, and gas.

A water molecule consists of three atoms: one oxygen (O) and two hydrogen (H) atoms, defining it as H2O.
The unique behavior of water arises from its molecular structure and interactions among its constituent atoms.
Water features polar covalent bonds due to the electronegativity of oxygen, which is higher than that of hydrogen.
As a result, the bonding electrons are drawn closer to oxygen, forming a V-shaped molecular structure.

Water is classified as a polar molecule, with distinct partial negative and partial positive ends:
- The oxygen end of the water molecule is more electronegative, giving it a slight negative charge.
- The hydrogen ends of the molecule exhibit a slight positive charge.
Overall, water maintains a slightly negative net charge, which is key to many of its physical properties.
The unique properties of water originate from interactions between the opposite charges of water molecules.

A hydrogen bond is defined as an intermolecular force that occurs between the charged regions of water molecules.
Specifically, the positive hydrogen end of one water molecule is attracted to the negative oxygen end of another water molecule.
Hydrogen bonds are typically weaker than covalent bonds but can be numerous, making their collective strength significant.
These bonds form, break, and reform frequently, significantly contributing to the physical properties of water.

Cohesion refers to the attraction between particles of the same substance, which in this case is water.
This property accounts for the phenomenon of surface tension, which is the measure of the strength of water’s surface.
Surface tension allows small objects to rest on top of the water surface and contributes to the formation of water droplets.

Adhesion is the attraction between water and different substances.
Water can form hydrogen bonds with various surfaces, such as glass, cotton, or soil, facilitating capillary action.
Capillary action enables water molecules to pull along in thin glass tubes, which is essential in biological systems (e.g., plants).

Water’s adhesive properties can also be observed in:
- Water droplets forming on cobwebs.
- Dew formation on surfaces.

Specific heat is defined as the amount of heat needed to raise or lower 1 gram of a substance by 1ºC.
Water has a high specific heat, meaning it resists changes in temperature when it absorbs or releases heat.
This property allows water to act as an insulator, maintaining a stable temperature in aquatic environments.

The heat of vaporization refers to the amount of energy required to convert 1 gram of a substance from a liquid to a gas.
In the case of water, significant energy must be absorbed to break hydrogen bonds before evaporation occurs, estimated at 540 calories per gram.
Water can thus absorb large amounts of energy without a significant change in temperature.

Ice is less dense than liquid water, which is why ice floats.
This property is crucial for aquatic life as it provides insulation to bodies of water during freezing conditions.
Liquid water molecules constantly form and break hydrogen bonds, allowing for fluctuating arrangements. Conversely, frozen water organizes into a crystal-like lattice arrangement, fixing the molecules at specific distances.

Homeostasis is the ability of organisms to maintain a stable internal environment despite external changes.
Water plays a critical role in this process, acting as a:
- Good insulator
- Temperature stabilizer
- Universal solvent
- Coolant in various biological processes
- Form of insulation against temperature extremes (e.g., ice covering lakes).

Water frequently exists as part of mixtures, which can be classified into two types: solutions and suspensions.

A solution consists of ionic compounds that dissolve in water, dispersing as ions, yielding an even distribution of solute throughout the solvent.
Key definitions:
- Solute: The substance being dissolved.
- Solvent: The substance into which the solute dissolves (in this case, water).

Suspensions are mixtures where substances do not dissolve in water but separate into particles that remain suspended, keeping them from settling out.

In water, about one molecule in 550 million will spontaneously dissociate into a hydrogen ion (H+) and a hydroxide ion (OH-).

The pH scale measures the concentration of hydrogen ions in a solution, ranging from 0 to 14:
- Neutral: pH 7
- Acidic: pH 0-7
- Basic: pH 7-14
Each unit change on the pH scale represents a tenfold change in ion concentration.
Example: A solution with a pH of 3 is 1,000 times more acidic than one with a pH of 6.

Acids: Strong acids typically have a pH between 1 and 3 and produce a large quantity of H+ ions.
Bases: Strong bases usually have a pH between 11 and 14, indicating a high concentration of OH- ions and a lower concentration of H+ ions.

Buffers are weak acids or bases that react with strong acids or bases, stabilizing pH and preventing drastic changes.
They are vital for maintaining homeostasis in biological systems.
Natural buffers include blood plasma, CO2 levels, and compounds like baking soda (HCO3).

An example of a buffering system in human physiology involves bicarbonate (HCO3-) and carbonic acid (H2 CO3).
Carbon dioxide reacts with water in blood plasma to form bicarbonate, which acts as a buffer to stabilize pH, modulating under various physiological conditions.
The buffering reaction is reversible and maintains equilibrium, effectively managing acid/base levels in the body.