Water and Life: Detailed Notes

Water's Central Role in Life

• This chapter centers on one compound: water (H₂O). Water is so integral to life on Earth that scientists use its presence as a key sign when searching for life elsewhere. The reference to the CD Institute (the search for extraterrestrial intelligence) is a playful introduction to this idea.

• The molecule shown with unequal sharing of electrons (polar covalent bond) is water. Unequal sharing of electron pairs is called a polar covalent bond, leading to partial charges on the molecule.

  • Polar covalent bond: unequal sharing of electrons.
  • Partial charges are indicated as $\delta^{+}$ (partial positive) and $\delta^{-}$ (partial negative).
  • In water, hydrogen atoms carry partial positive charges ($\delta^{+}$) and the oxygen carries a partial negative charge ($\delta^{-}$).

• Oxygen is highly electronegative, making water a polar molecule and enabling hydrogen bonding between water molecules. Oxygen is the most electronegative element among common biological elements, with fluorine being more electronegative overall, though not biologically central in this context.

• Hydrogen bonds: attractions between water molecules (not as strong as covalent or ionic bonds). They are weaker individually but collectively drive water's remarkable properties.

• Water’s polarity allows it to interact with charged or polar solutes via hydrogen bonding, which leads to four emergent properties that support life.

Emergent Properties of Water

• Four emergent properties arise from water's polarity and hydrogen bonding, enabling life:

  • Cohesion: water molecules stick to each other.
  • Adhesion: water sticks to other surfaces (e.g., cell walls, xylem).
  • Surface tension: cohesive forces at the surface create a skin-like layer.
  • Temperature moderation: hydrogen bonding makes water resist rapid temperature changes (high specific heat).

• Cohesion and plant water transport: water can move from roots to leaves with no mechanical pumping. Air currents around leaves pull water up, effectively acting like a million tiny straws. Water molecules cohere to each other and adhere to the walls of the conducting channels, enabling upward transport against gravity.

• Surface tension: cohesion among water molecules generates surface tension, allowing organisms like water striders to rest on water surfaces.

• Temperature as a form of energy: heat vs. temperature

  • Heat is a measure of total kinetic energy of all molecules in a sample.
  • Temperature is a measure of the average kinetic energy of molecules in a sample.
  • Example: a hot cup of coffee vs. a swimming pool. The coffee may have a higher temperature, but the pool has more total heat due to a larger number of water molecules.
  • Typical unit for heat: calories; 1 calorie is the energy required to raise the temperature of 1 g of water by 1 °C.
  • Food energy on labels uses kilocalories (kcal), often called Calories (with a capital C): 1 Cal = 1 kcal = 1000 cal.
  • SI unit for energy is the joule (J). The conversion between calories and joules is interconvertible; 1 cal ≈ 4.184 J.
  • Important relations:
  • $1\ \text{cal} = 4.184\ \text{J}$
  • $1\ \text{Cal} = 1\ \text{kcal} = 1000\ \text{cal} = 4184\ \text{J}$

• Specific heat: a physical property that defines how much heat is required to change the temperature of a material. Water has a high specific heat due to hydrogen bonding, so it resists temperature change; metals typically have a lower specific heat.

• Phase changes and density of water:

  • In most substances, density increases from gas to liquid to solid. Ice floats on liquid water because ice is less dense.
  • The floating ice is due to hydrogen bonds forming a network that stabilizes at a greater distance on average in the solid phase, creating a less dense structure.
  • This density anomaly is crucial for life: as water freezes, a layer of ice forms on the surface while liquid water remains beneath, allowing life to persist below during cold periods.

• Liquid water as a solvent:

  • Water is a great polar solvent as long as solutes are polar or charged.
  • A solution is a uniform mixture of solvent and solute (the substance dissolved).
  • Hydration shells form around dissolved ions and polar molecules, stabilizing them in solution.

Hydration and Dissolution

• Hydration shell: the clathrate-like arrangement in which water molecules surround a dissolved ion or polar molecule.

  • Example: table salt (sodium chloride, NaCl) dissolves well in water because Na⁺ and Cl⁻ ions separate and become surrounded by water's partial charges.
  • Na⁺ ions are surrounded by water’s partial negative charges (on the oxygen atoms, $\delta^{-}$), while Cl⁻ ions are surrounded by water’s partial positive charges (on the hydrogen atoms, $\delta^{+}$).

• Hydration shells prevent ions from recombining while water remains present; removal of water leads to dissolution reversal.

• Big molecules: large molecules (hundreds to millions of Daltons) with charges on their surfaces can also form hydration shells and dissolve if they are polar or charged; nonpolar big molecules are not well dissolved by water.

• Nonpolar molecules and hydrophobicity:

  • Water interacts poorly with nonpolar molecules.
  • Oil and water do not mix because oil is nonpolar; hydrophobic substances are water-fearing.
  • In cells, the balance between hydrophilic (water-loving) and hydrophobic (water-fearing) substances influences structure and function.

Measuring Matter: From Daltons to Moles

• Atomic masses and the periodic table:
  • Atomic mass units (amu) appear under atomic symbols as approximate masses (no units shown on some tables).
  • 12.011 is the average mass of a carbon atom, i.e., about 12.011 Daltons per atom.
  • This means carbon has an average atomic mass of 12.011 Daltons (amu).
  • These numbers help estimate how much of a substance is present in macroscopic samples.
• From Daltons to grams per mole:
  • The two ideas are related: the mass of one mole of a substance (its molar mass) in grams equals the mass of one atom or molecule in Daltons (amu).
  • In other words, 1 amu = 1 g/mol, linking microscopic mass to macroscopic measurement.
• The mole and Avogadro's number:
  • A mole is a counting unit for amount of substance, analogous to dozen or pair, but for particles.
  • Avogadro's number: $N_A = 6.022\times 10^{23}$ particles per mole.
  • This number converts between the number of molecules and the amount in moles.
  • Conceptually: 1 mole of any substance contains $N_A$ entities and weighs its molar mass in grams.
• Molar mass and formula weight:
  • To determine molar mass, sum the atomic masses of all atoms in the formula.
  • Example: Sucrose has formula $\mathrm{C<em>{12}H</em>{22}O_{11}}$.
  • Molar mass calculation:
  • $M_{\text{sucrose}} = 12\times(12.011) + 22\times(1.008) + 11\times(15.999) \approx 342.296\ \text{g/mol}$
  • This is often rounded to $\approx 342\ \text{g/mol}$ or equivalently about $342\ \text{Da}$ per molecule.
  • Therefore, 1 mole of sucrose weighs about 342 g and contains $N_A$ molecules.
• Quick recap on moles:
  • A mole is a counting unit that links macroscopic mass to number of particles via the molar mass and Avogadro's number.
  • Example conclusion: a sample of 342 g of sucrose contains $N_A$ molecules of sucrose.
• Visualization of scale:
  • A short video illustrates just how large a mole is; the key idea is that a mole contains an enormous number of particles, enabling macroscopic measurements to reflect microscopic reality.

Key Notation Recap (Useful References)

• Polar covalent bonds: unequal sharing of electrons; $\delta^{+}$ on hydrogen, $\delta^{-}$ on oxygen.
• Hydrogen bond: an attraction between polar molecules (weaker than covalent/ionic bonds).
• Calorie definitions:
  • $1\ ext{cal} = 4.184\ \text{J}$
  • $1\ \text{Cal} = 1\ \text{kcal} = 1000\ \text{cal} = 4184\ \text{J}$
• Energy units and conversions:
  • Energy is measured in joules (J) or calories (cal).
  • Temperature vs. heat: temperature is average kinetic energy; heat is total kinetic energy across all molecules.
• Hydration and dissolution:
  • Hydration shell around ions and polar solutes in water.
• Molar mass and the mole:
  • 1 amu equals 1 g/mol; $N_A = 6.022\times 10^{23}$.
  • Sucrose molar mass: $M_{\text{sucrose}} \approx 342\ \text{g/mol}$.
• Important distinctions:
  • Polar and ionic substances are typically hydrophilic (water-loving).
  • Nonpolar substances are typically hydrophobic (water-fearing).