Study Notes on Water and Life

Species-Specific Ecology: Ringed seals (Phoca hispida) are obligate ice-dwelling mammals. They rely on stable Arctic sea ice for critical life stages, including pupping, nursing, and molting.
Subnivean Lairs: These seals create caves under the snow (subnivean lairs) on top of the ice to protect their young from predators and extreme cold.
Threat from Climate Change: Accelerated melting reduces the availability of stable ice platforms and decreases snow depth. If snow depth is less than $20\text{ cm}$ , lairs may collapse or fail to provide adequate insulation, leading to increased pup mortality from freezing or predation.

Polar Covalent Bonds: Within a single water molecule ( $H_2O$ ), the oxygen atom is more electronegative than the hydrogen atoms, meaning it has a stronger pull on shared electrons. This results in polar covalent bonds.
Molecular Polarity: The molecule is bent at an angle of approximately $104.5^{\circ}$ due to two lone pairs of electrons on the oxygen atom (VSEPR theory). This uneven geometry creates a dipole where the oxygen end is partially negative ( $\delta^-$ ) and the hydrogen ends are partially positive ( $\delta^+$ ).

Intermolecular Attraction: Hydrogen bonds form when the $\delta^+$ hydrogen of one water molecule is attracted to the $\delta^-$ oxygen of a neighboring molecule.
Strength: While individual hydrogen bonds are much weaker (about $1/20^{th}$ the strength) than covalent bonds, their collective effect is powerful.
States of Matter:
- Liquid: Hydrogen bonds behave like "flickering clusters," forming and breaking in trillionths of a second.
- Solid (Ice): As temperature drops, molecules move too slowly to break bonds. They lock into a crystalline lattice where each molecule is hydrogen-bonded to four partners at a fixed distance.

Cohesion: The tendency of water molecules to stick to each other. This allows for the transport of water against gravity in plants (xylem transport).
Adhesion: The clinging of water to other substances, such as the cell walls of plant vessels, which helps counter the downward pull of gravity.
Surface Tension: Resulting from the collective strength of hydrogen bonds at the air-water interface, water exhibits an unusually high surface tension ( $72.8\text{ dynes/cm}$ at $20^{\circ}C$ ).

High Specific Heat: Water's specific heat is $1\text{ cal/g}\cdot{}^{\circ}C$ (or $4.184\text{ J/g}\cdot{}^{\circ}C$ ).
- Heat must be absorbed to break hydrogen bonds before the molecules can move faster.
- Heat is released when hydrogen bonds form as water cools.
Environmental Buffer: This property allows large bodies of water to absorb massive amounts of solar heat during the day with only a slight change in temperature, stabilizing local climates.

Density Anomaly: Most substances contract and become denser when they solidify. Water reaches its maximum density at $4^{\circ}C$ . Below this temperature, it begins to expand as the crystalline lattice forms.
Ecological Importance: Ice floats because it is roughly $10\%$ less dense than liquid water. This floating layer insulates the liquid water below, preventing entire lakes from freezing solid and allowing life to survive beneath the surface.

Mechanism: Water is an effective solvent because its polar molecules are attracted to ions and polar solutes.
Hydration Shells: When an ionic compound (like $NaCl$ ) is placed in water, the oxygen ends of water molecules surround the sodium cations ( $Na^+$ ), and the hydrogen ends surround the chloride anions ( $Cl^-$ ), effectively shielding and dissolving the ions.
Hydrophobic Exclusion: Nonpolar substances (like oils) do not form hydrogen bonds and are pushed aside by water, a phenomenon critical for the formation of cell membranes (lipid bilayers).

Dissociation of Water: Occasionally, a hydrogen atom leaves its electron behind and joins another water molecule: $2H2O \rightleftharpoons H3O^+ + OH^-$ .
Mathematical Definition: pH is calculated as the negative logarithm (base $10$ ) of the hydrogen ion concentration: $pH = -\log[H^+]$ .
Neutrality: In pure water at $25^{\circ}C$ , the concentration of both $H^+$ and $OH^-$ is $10^{-7}\text{ M}$ , resulting in a pH of $7$ .

Homeostasis: Most biological fluids must maintain a pH between $6$ and $8$ .
Bicarbonate System: A key buffer in human blood is the carbonic acid ( $H2CO3$ ) and bicarbonate ( $HCO_3^-$ ) pair. If pH rises, the acid dissociates to release $H^+$ ; if pH drops, the base absorbs $H^+$ .

Chemical Process: As humans burn fossil fuels, more $CO2$ is absorbed by the oceans. The reaction is: $CO2 + H2O \rightarrow H2CO_3$ (carbonic acid).
Loss of Carbonate: The carbonic acid dissociates into $H^+$ and bicarbonate. These extra hydrogen ions then react with carbonate ions ( $CO_3^{2-}$ ) to form more bicarbonate.
Biological Conflict: By "scavenging" carbonate ions, the process reduces the concentration of $CO3^{2-}$ available for calcifying organisms (e.g., corals, mollusks) to build their calcium carbonate shells ( $CaCO3$ ).