Water and Its Properties

Water & the Fitness of the Environment

Water's Polarity

• A molecule is considered polar when there is an uneven distribution of charges.
• This uneven distribution is due to differences in electronegativity among the elements in the compound.
• Oxygen atom carries a partial negative charge ($\delta^-$), hydrogen atoms carry a partial positive charge ($\delta^+$.

Structure of Water

• Water molecules are held together by covalent bonds.
• Covalent bonds involve the sharing of electrons between oxygen and hydrogen atoms WITHIN each molecule.
• Hydrogen bonds form due to water's polarity.

Hydrogen Bonds

• Hydrogen bonds are intermolecular forces, not true bonds involving electron sharing.
• They occur between atoms of DIFFERENT molecules.
• Specifically, they arise when an electronegative atom (oxygen or nitrogen) is attracted to a hydrogen atom on another molecule that is already bonded to a nitrogen or oxygen atom.
• Hydrogen bonds are weak and easily broken/reformed.
• They are crucial for many life processes, including:
  • Water properties
  • DNA replication
  • Protein structure

Water's Hydrogen Bonding Capacity

• Due to its polarity, a single water molecule can form hydrogen bonds with up to four other water molecules.

Water's Special Properties

• Cohesion & Adhesion
• Good solvent
• Lower density as a solid
• High specific heat
• High heat of vaporization

Cohesion and Adhesion

• Cohesion: Water molecules are attracted to each other through hydrogen bonding.
  • This is responsible for surface tension.
  • Example: Drinking from a straw.
• Adhesion: Water molecules are attracted to other polar molecules through hydrogen bonding.
  • This leads to capillary action.
  • Examples: Meniscus formation, water climbing up a paper towel or cloth.

Surface Tension

• Surface tension allows some insects to walk on water.

Capillary Action

• Capillary action involves both cohesion and adhesion.
• Important in plants, facilitating water transport up xylem cells.

Transpiration

• The process of water movement through a plant and its evaporation from aerial parts, such as leaves.

Water as a Solvent

• Water's polarity makes it an excellent solvent for polar and ionic compounds.
• Water molecules surround positive and negative ions, dissolving them.
• Solvents dissolve solutes, creating solutions.

Dissolving Process

• Cations (e.g., $Na^+$) are attracted to the negative pole of water.
• Anions (e.g., $Cl^-$) are attracted to the positive pole of water.
• The dissolved ions are unable to reassociate into a solid.

Hydrophilic vs. Hydrophobic Substances

• Hydrophilic: Water-loving substances.
  • Includes ionic compounds and polar covalent compounds
  • Examples: salt, ions
• Hydrophobic: Water-fearing substances.
  • Nonpolar covalent compounds.
  • Examples: oil, fat

Density of Water

• Water is less dense as a solid than as a liquid.
• Water molecules spread out upon freezing, forming a crystal lattice.
• The gaps in the lattice cause solid water to occupy more volume, thus decreasing its density.
• $D = \frac{M}{V}$

Importance of Ice Floating

• If ice sank, oceans and lakes would freeze solid.
• Floating ice insulates the water below, preventing it from freezing.
• Seasonal turnover of lakes facilitate oxygen cycling.

High Specific Heat

• Water has a high specific heat.
• Specific heat: the amount of heat required to raise 1 gram of a substance by 1 degree Celsius.
• Water resists changes in temperature due to its high specific heat.

Importance of High Specific Heat

• Oceans maintain relatively constant temperatures.
• This leads to:
  • Climate moderation
  • Stable ocean temperatures

High Heat of Vaporization

• Water has a high heat of vaporization.
• Evaporative cooling: As liquid evaporates, the surface cools because the hottest molecules leave as gas.

Dissociation of Water

• Water molecules can dissociate into:
  • Hydronium ions ($H_3O^+$, an acid)
  • Hydroxide ions ($OH$, a base)

Ionization of Water and pH

• When water ionizes, $H^+$ splits off from $H_2O$, leaving $OH^-$.
• If $[H^+] = [OH^-]$, water is neutral.
• If [H^+] > [OH^-], water is acidic.
• If [H^+] < [OH^-], water is basic.
• pH scale: Measures how acidic or basic a solution is.
  • Ranges from 1 to 14.
  • pH decreases as $H^+$ increases.
  • Each pH unit represents a tenfold difference in $H^+$ and $OH^-$ concentrations.

Buffers

• A buffer minimizes changes in the concentrations of $H^+$ and $OH^-$ in a solution.
• Buffers accept or donate $H^+$.
• Cellular pH must be kept around 7.4 for proper function.
• Example: Bicarbonate and carbonic acid form an acid-base equilibrium.

Acid Precipitation

• Caused by pollutants like $SO<em>2$ and $NO</em>x$.