Water and Life

Define polarity and how it relates to hydrogen bonds.
Describe the four properties of water, what is responsible for them, and their importance for life.
Explain why water makes a good solvent.
Define molarity and procedures for making solutions.
Understand the pH scale and how to interpret pH values of water and other substances.
Define buffers, their functionality, and their importance in biology.

Definition: Atoms are the basic building blocks of all matter in the universe.
Everything around us, including air and devices, consists of atoms.
Structure of an Atom:
- Nucleus: Contains positively charged protons and neutral neutrons.
- Electron Cloud: Surrounds the nucleus and contains negatively charged electrons.
Role of Electrons: Responsible for bonding between atoms.

Formation of Ionic Compounds:
- Metals lose electrons while nonmetals gain electrons to achieve an octet (eight electrons in their outer shell).
Ionic Bond Definition: A bond formed when one atom gives up one or more electrons to another atom.
Typically occurs between:
- Metal (loses electrons, becomes a cation)
- Nonmetal (gains electrons, becomes an anion)
Attraction Mechanism: Opposite charges attract, resulting in a strong electrostatic force that holds the ions together.

Definition: A covalent bond is formed when two atoms share one or more pairs of electrons.
Example: In a water molecule (H₂O), two hydrogen atoms share electrons with one oxygen atom via covalent bonds.
Stability: Sharing of electrons allows atoms to fill their outer electron shells and become more stable.

Approximately three-quarters of Earth’s surface is covered in water, predominantly in liquid form.
Water also exists as a solid (ice) and a gas (water vapor).
Unique properties of water arise from polar covalent bonds and hydrogen bonds, making it essential for life.

Structure of Water Molecule:
- Comprised of one oxygen atom and two hydrogen atoms, connected by two single covalent bonds.
Electronegativity: Oxygen is more electronegative than hydrogen, which causes electrons to spend more time closer to the oxygen atom.
Resulting in:
- Polar covalent bonds
- Partial charges (polarity)

Definition: A hydrogen bond forms between partially positive hydrogen atoms and partially negative oxygen atoms.

The four key properties of water that facilitate an environment suitable for life are:
- Cohesive Behavior: Hydrogen bonds hold water molecules together, facilitating the movement of water against gravity in plants.
- Ability to Moderate Temperature: Water absorbs and releases heat, helping to stabilize temperatures.
- Expansion upon Freezing: Ice is less dense than liquid water, allowing it to float.
- Versatility as a Solvent: Water can dissolve many substances due to its polarity.

Mechanism: Hydrogen bonds contribute to the cohesion of water molecules.
Biological Implication:
- Cohesion helps transport water from roots to leaves via evaporation.
- Adhesion, or clinging to other substances, assists water movement upward in plants.
Surface Tension: The cohesive force results in surface tension, where water molecules stick together at the surface, forming a film.

Heat Absorption and Release: Water absorbs heat from warmer air and releases it to cooler air.
Specific Heat:
- Water has a high specific heat of 1 calorie/(g · °C), indicating it requires significant heat to change its temperature.
- Comparison: Ethanol has a lower specific heat of 0.6 calorie/(g · °C).
Heat of Vaporization: Water requires a large amount of heat to convert from liquid to gas, which contributes to evaporative cooling (e.g., sweating leads to cooling).

Density: Ice is less dense than liquid water, enabling it to float.
Structure: In ice, each water molecule forms stable hydrogen bonds with four other molecules, creating a crystalline structure that keeps molecules apart compared to liquid water.
Ecological Impact:
- If ice sunk, aquatic life would be severely impacted as ponds and lakes would freeze solid.
- Floating ice insulates liquid water below, allowing aquatic ecosystems to thrive under ice.

Solution Definition: A solution is a homogeneous mixture of substances.
Components of a Solution:
- Solvent: The dissolving agent (water in aqueous solutions).
- Solute: The substance being dissolved.
Polarity Related to Solvation: Polar covalent bonds allow water to dissolve various polar and ionic compounds.

Example: When sodium chloride (NaCl) dissolves in water, spheres of hydration form around ions, demonstrating water's interaction with ionic molecules.
Hydrophilic vs. Hydrophobic: Water can dissolve hydrophilic (water-attracting) substances, but not hydrophobic (water-repelling) substances, which is significant for cell membrane structure.

Molarity Definition: The concentration unit most commonly used in science, defined as moles of solute per liter of solution.
Mole Definition: One mole (mol) equals $6.02 imes 10^{23}$ molecules of any substance (Avogadro’s number).
Molarity Example:
- 1M solution of sodium chloride means 1 mole of NaCl in 1 liter of water.
Molecular Mass: The mass of one mole of a substance, calculated as the sum of the mass numbers of all atoms in the molecule.
- Example: For sodium chloride (NaCl), molecular mass = 58 grams.

pH Definition: pH scale measures acidity or basicity; water is neutral at pH 7.
Hydronium Ion: Often represented by a hydrogen ion (H+); the concentration of H+ in pure water is $10^{-7} ext{ M}$ .
Acids and Bases:
- Acid: Increases H+ concentration (e.g., $ext{HCl} <br>ightarrow ext{H}^+ + ext{Cl}^-$ )
- Base: Increases OH- concentration (e.g., $ext{NaOH} <br>ightarrow ext{Na}^+ + ext{OH}^-$ )
pH Calculation:
- $ext{pH} = - ext{log} [ ext{H}^+]$
- Example: Water has a pH of 7, while urine (with $[H^+] = 10^{-6}$ ) is 10 times more acidic than water, and coffee is 100 times more acidic than water.

Importance: Maintain pH within a narrow range for biological systems (pH 6-8).
Functionality: Buffers can accept excess H+ ions or donate H+ ions when necessary to minimize pH changes.
Common Buffer Example: Carbonic acid (H2CO3) found in human blood acts as a buffer:
- Forms when CO2 reacts with water and can dissociate into bicarbonate ion (HCO3-) and H+.
- Lowers pH by donating H+, raises pH by accepting H+.