뒤로Water and Life: Properties, Acids, Bases, and pH
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Water and Life
Polar Covalent Bonds and Hydrogen Bonding
Water's unique properties arise from its molecular structure and the nature of its chemical bonds. The electrons in water's polar covalent bonds spend more time near the oxygen atom, making water a polar molecule with an uneven charge distribution. This polarity enables water molecules to form hydrogen bonds with each other, which are critical for many of water's biological functions.
Polar covalent bond: A bond where electrons are shared unequally, resulting in partial charges.
Hydrogen bond: A weak attraction between the partial positive charge of hydrogen and the partial negative charge of oxygen in adjacent water molecules.
Example: Water molecules interact through hydrogen bonds, leading to high cohesion and surface tension.
Four Emergent Properties of Water
Water exhibits four key properties that make it essential for life: cohesive behavior, ability to moderate temperature, expansion upon freezing, and versatility as a solvent. These properties are all consequences of hydrogen bonding.
Cohesion: Water molecules stick together, resulting in high surface tension.
Moderation of temperature: Water absorbs and releases heat slowly, stabilizing environmental and organismal temperatures.
Expansion upon freezing: Ice is less dense than liquid water, allowing it to float.
Versatility as a solvent: Water dissolves a wide range of substances due to its polarity.
Cohesion and Surface Tension
Cohesion refers to the attraction between water molecules, which is responsible for water's high surface tension. This property allows certain organisms, such as insects, to walk on water and enables the upward transport of water in plants.
Surface tension: The difficulty of breaking the surface of a liquid due to cohesive forces.
Adhesion: Attraction between water and other substances, such as plant cell walls.
Example: Water transport in plants relies on both cohesion and adhesion.
Moderation of Temperature
Water's high specific heat allows it to absorb or release large amounts of heat with minimal temperature change. This property is vital for maintaining stable temperatures in organisms and environments.
Kinetic energy: Energy of motion; thermal energy is the kinetic energy of molecules.
Specific heat: The amount of heat required to change the temperature of 1 g of a substance by 1ºC.
Water's specific heat: 1 cal/(g·ºC), much higher than most substances.
Equation: (where q = heat, m = mass, c = specific heat, ΔT = temperature change)
Example: Coastal areas have moderate climates due to water's high specific heat.
Evaporative Cooling
Evaporation is the process by which molecules transition from liquid to gas. As water evaporates, the surface cools, helping regulate temperature in organisms and environments.
Heat of vaporization: The heat required to convert 1 g of liquid to gas.
Evaporative cooling: The cooling effect as the hottest molecules leave the surface.
Example: Sweating cools the body through evaporative cooling.
Expansion Upon Freezing
Water is less dense as a solid than as a liquid, which is unusual among substances. At 0ºC, water molecules form a crystalline lattice, making ice about 10% less dense than liquid water. This property allows ice to float, insulating aquatic life below.
Density: Water reaches its maximum density at 4ºC.
Example: Floating ice prevents lakes and oceans from freezing solid.
Water as a Solvent
Water's polarity makes it an excellent solvent, capable of dissolving ionic and polar substances. In aqueous solutions, water surrounds ions and polar molecules, forming hydration shells.
Solution: Homogeneous mixture of substances.
Solvent: The dissolving agent (water in aqueous solutions).
Solute: The substance dissolved.
Hydration shell: Sphere of water molecules surrounding dissolved ions.
Example: Table salt (NaCl) dissolves in water, forming Na+ and Cl– ions surrounded by water.
Hydrophilic and Hydrophobic Substances
Substances that interact well with water are termed hydrophilic, while those that do not are hydrophobic. Hydrophobic molecules, such as oils, are major components of cell membranes.
Hydrophilic: Affinity for water; often polar or ionic.
Hydrophobic: Repelled by water; often nonpolar.
Example: Oil is hydrophobic and does not mix with water.
Solute Concentration in Aqueous Solutions
Most biological reactions occur in aqueous solutions. The concentration of solutes is often measured in moles and molarity.
Molecular mass: Sum of atomic masses in a molecule (e.g., H2O = 18 daltons).
Mole: 6.02 × 1023 molecules (Avogadro's number).
Molarity (M): Moles of solute per liter of solution.
Example: Making a 1 M solution of NaOH involves dissolving 40 g NaOH in 1 L water.
Acids, Bases, and pH
Acidic and Basic Conditions
Water molecules can dissociate into hydrogen ions (H+) and hydroxide ions (OH–). The concentrations of these ions are crucial for biological processes.
Acid: Increases H+ concentration; donates protons or accepts electron pairs.
Base: Reduces H+ concentration; accepts protons or donates electron pairs.
Strong acids/bases: Dissociate completely in water.
Weak acids/bases: Partially dissociate, allowing reversible reactions.
The pH Scale
The pH scale measures the acidity or basicity of a solution, defined as the negative logarithm of the H+ concentration.
Equation:
Neutral solution: [H+] = 10–7 M, pH = 7.
Acidic: pH < 7; Basic: pH > 7.
Logarithmic scale: Each unit change represents a tenfold difference in H+ concentration.
Buffers and Biological Importance
Buffers are substances that minimize changes in pH by accepting or donating H+ ions. They are essential for maintaining stable internal conditions in living organisms.
Buffer: Contains a weak acid and its conjugate base; resists changes in pH.
Example: The bicarbonate buffer system in human blood maintains pH near 7.4.
Additional info: The bicarbonate buffer system is crucial for preventing acidosis and alkalosis in humans. The reactions are:
