Water — Chapter 2.2 Notes (Campbell)

Water is fundamental to life, comprising approximately $60\%-70\%$ of the human body by weight and covering roughly $70\%$ of the Earth's surface. It is believed that life originated in ancient aquatic environments, and water remains the primary medium for most cellular chemistry and metabolic reactions within the cytoplasm.
Key properties that make water essential for life include its high heat capacity, high heat of vaporization, excellent solvent capabilities for polar molecules, strong cohesive and adhesive properties, and its ability to dissociate into ions, which is crucial for regulating pH.

Water molecules are inherently polar, consisting of two hydrogen atoms covalently bonded to a single, more electronegative oxygen atom. This uneven distribution of electrons creates distinct partial charges within the molecule.
The oxygen atom, being more electronegative, pulls shared electrons closer to itself, resulting in a partial negative charge ( $\delta^{-}$ ) on the oxygen. Conversely, the hydrogen atoms acquire partial positive charges ( $\delta^{+}$ ).
This polarity enables the formation of hydrogen bonds, weak electrostatic attractions that occur between the partially positive hydrogen of one water molecule and the partially negative oxygen of an adjacent water molecule. These bonds are responsible for many of water's unique properties.

Water exhibits an unusual property: it is less dense as a solid (ice) than as a liquid. This anomaly is vital for life on Earth.
This reduced density in the solid state is a direct consequence of how hydrogen bonds orient themselves during freezing. As water cools below $4^\circ \text{C}$ , the hydrogen bonds become more stable and arrange water molecules into a rigid, open, crystalline lattice structure.
In this lattice, water molecules are held farther apart than they are in the more disordered liquid state. This expansion leads to ice having a lower density, allowing it to float on liquid water.
The insulating layer of ice that forms on the surface of bodies of water in cold temperatures protects aquatic life below from freezing solid, thereby sustaining diverse ecosystems in winter.

Water possesses the highest specific heat capacity of any common liquid, defined as the amount of heat energy one gram of a substance must absorb or lose to change its temperature by one degree Celsius (approx. $4.184 \ \text{J/(g} \cdot ^\circ \text{C})$ for water).
This high capacity is mainly due to the extensive network of hydrogen bonds between water molecules. A significant amount of absorbed energy is required to break these bonds before kinetic energy (and thus temperature) of the molecules can increase substantially.
Consequently, water takes a considerable amount of energy and time to heat up and, conversely, a long time to cool down. For instance, water's heat capacity is about 5 times that of sand, explaining why coastal land masses cool faster than the adjacent sea.
This property is profoundly important for living organisms, as it helps to buffer internal body temperatures against external fluctuations, allowing for more stable physiological conditions. It also moderates global climates.

The heat of vaporization is the substantial amount of energy required to change one gram of a liquid substance into a gas (for water, it's approx. $2260 \ \text{J/g}$ or $540 \ \text{cal/g}$ at $100^\circ \text{C}$ ).
Water has an exceptionally HIGH heat of vaporization. This is again attributed to the strong hydrogen bonds, which must be overcome for water molecules to escape the liquid phase and enter the gaseous state.
Because of the large energy input needed for evaporation, water acts effectively as a heat reservoir and does not boil as quickly as many other liquids.
The process of evaporation, requiring so much energy absorption from the surroundings, produces a significant cooling effect. This is crucial for regulating temperature in both organisms (e.g., evaporative cooling through sweating in mammals, or transpiration in plants) and the environment.

Water is often called the "universal solvent" because of its ability to dissolve more substances than any other liquid. This remarkable property stems directly from its polarity and its capacity to form hydrogen bonds.
How it works:
- For ionic compounds (like NaCl), the partial negative charge of oxygen in water molecules (δ−δ−) is attracted to the positive ions (e.g., Na+Na+), while the partial positive charges of hydrogen (δ+δ+) are attracted to the negative ions (e.g., Cl−Cl−). This interaction, called hydration, surrounds the ions and pulls them away from the crystal lattice, dissolving them.
- For polar covalent compounds (like sugar or proteins), water molecules form hydrogen bonds with the polar functional groups on these molecules, effectively surrounding them and keeping them in solution.
This solvent capability is vital for life, enabling the transport of nutrients, minerals, and waste products both within organisms and in ecosystems. It facilitates countless biochemical reactions by allowing reactants to mix and interact readily.