Physical and Chemical Properties of Water and pH Study Notes

Chemical Bonds and Molecular Polarity

• Chemical Bond Strength and Types:

  • Covalent Bonds (Polar and Non-polar)
  • Ionic Bonds
  • Hydrogen Bonds
  • Van der Waals Forces

• Electronegativity:

  • Defined as the measure of an atom's attraction for the electrons it shares in a chemical bond with another atom.
  • A higher electronegativity indicates a stronger attraction for shared electrons.
  • Covalent bonds range in their degree of valence electron sharing based on the electronegativity difference between the bonded atoms.

• Types of Covalent Bonds:

  • Nonpolar Covalent Bond: Electrons are shared equally between atoms.
  • Polar Covalent Bond: Electrons are shared unequally.
  • The more electronegative atom carries a partial negative charge, denoted as $\delta-$.
  • The atom deprived of electrons carries a partial positive charge, denoted as $\delta+$.
  • A molecule is considered polar if one end is partially positive and the other partially negative, though the net charge of the entire molecule remains zero.

• Molecular Polarity and Symmetry:

  • Polar Groups: Oxygen, nitrogen, and sulfur share electrons unequally with hydrogen. Functional groups like $-OH$, $-NH$, and $-SH$ tend to be located asymmetrically in biological molecules, creating polar regions.
  • Nonpolar Groups: Carbon-hydrogen bonds (e.g., in methane) are typically arranged symmetrically. The partial charges cancel out, resulting in a nonpolar molecule.

• Polar and Nonpolar Interactions:

  • Hydrophilic ("water loving"): Polar molecules that associate readily with water. Polar molecules attract other polar molecules and charged ions/molecules, forming polar associations.
  • Hydrophobic ("water fearing"): Nonpolar substances excluded by water and other polar molecules. These molecules tend to clump together in nonpolar associations to reduce the surface area exposed to the polar environment.

Hydrogen Bonds

• Definition: Attractions between a partially positive hydrogen atom (which is sharing electrons unequally with oxygen, nitrogen, or sulfur) and a partially negative atom involved in a different covalent bond.
• Bond Types:
  • Intramolecular: Occur between atoms within the same molecule.
  • Intermolecular: Occur between atoms in different molecules.
• Characteristics and Significance:
  • While individual hydrogen bonds are weak, they are collectively strong when numerous.
  • They stabilize the three-dimensional structures of large biological molecules, such as proteins.
  • They are responsible for the life-sustaining properties of water.
  • Hydrogen bonds begin to break extensively at temperatures exceeding $45\,^{\circ}\text{C}$.
• Visualization: Non-contact atomic force microscopy (NC-AFM) has been used to visualize hydrogen bonds (e.g., Zhang et al., Science, 2013).

Physical Properties of Water

• The Water Lattice:

  • Liquid Water: Forms a dynamic water lattice where each molecule constantly breaks and reforms hydrogen bonds with neighbors, averaging $3.4$ bonds per molecule.
  • Ice Lattice: A rigid, crystalline structure where each molecule forms $4$ hydrogen bonds. This structure spaces molecules further apart than in the liquid lattice.

• Density:

  • Ice is approximately $10\%$ less dense than liquid water, allowing it to float. This property is vital for aquatic life, as ice forms a top layer and provides insulation.
  • Water reaches its maximum density at $4\,^{\circ}\text{C}$.
  • Biological Impact in Winter: In a frozen pond, the ice layer allows river otters to visit, aquatic insects to survive in air pockets, fish to take oxygen from the water, and protists to provide food, while common frogs and pond turtles hibernate.

• Temperature Regulation:

  • The hydrogen-bond lattice retards the escape of water molecules during heating.
  • Water remains liquid over a wide range ($0\,^{\circ}\text{C}$ to $100\,^{\circ}\text{C}$).
  • Specific Heat: The amount of heat energy required to increase the temperature of a specific quantity of water. Water has a high specific heat, meaning it can absorb or release large amounts of heat with minimal temperature change.

• Measurements of Heat:

  • calorie (small calorie): The heat energy required to raise the temperature of $1\,\text{g}$ of water by $1\,^{\circ}\text{C}$.
  • Calorie (capital C): A kilocalorie (kcal) or $1,000$ calories.

• Heat of Vaporization:

  • A large amount of heat ($540$ to $586\,\text{calories/g}$) must be added to break hydrogen bonds and transform liquid water into gas.
  • This allows for evaporative cooling in humans and other organisms.

• Cohesion, Adhesion, and Surface Tension:

  • Cohesion: The tendency of water molecules to stick to each other due to the hydrogen-bond lattice.
  • Adhesion: The tendency of water molecules to stick to the walls of tubes (like plant vessels) by forming hydrogen bonds with charged or polar groups.
  • Surface Tension: Unbalanced bonding at the surface (where water faces air) puts surface molecules under tension. This makes the surface resistant to separation, causing water to form droplets and supporting the weight of small insects.

Water as a Biological Solvent

• Organization of Cells:

  • The water lattice resists nonpolar molecules, forcing them into nonpolar associations. This creates distinct polar and nonpolar environments critical for cell organization.
  • Biological Membranes: Consist of lipid molecules with dual polarity. One end is polar; the other is nonpolar. Because they are surrounded by water, the molecules associate into a bilayer, exposing only the polar ends to water.

• Solvent Properties:

  • Water is a solvent for charged or polar molecules.
  • Hydration Layer: Water molecules coat the surface of other polar/charged molecules or ions, reducing their attraction to one another and promoting entry into the solution.
  • Water (solvent) surrounds the dissolved substance (solute), preventing reassociation (e.g., sodium and chloride ions).

Chemical Reactions and Concentrations

• Concentration: The number of molecules or ions in a unit volume (e.g., $mL$ or $L$).
• Calculations:
  • Mass Number: The number of protons and neutrons (e.g., Carbon has a mass number of $12$).
  • Avogadro’s Number: $6.022 \times 10^{23}$ atoms or molecules per mole.
  • Molarity (M): The number of moles of a substance dissolved in $1\,L$ of solution. Two solutions with the same volume and molarity contain the same number of molecules.
• Formulas:
  • $n = \frac{m}{M}$ (Number of moles = mass / molecular mass)
  • $c_1v_1 = c_2v_2$ (Standard dilution equation)

Water Ionization and pH

• Dissociation of Water:

  • Water dissociates into positively charged hydrogen ions ($H^+$ or protons) and negative hydroxide ions ($OH^-$).
  • Reversible reaction: $2H_2O \rightleftharpoons H_3O^+ + OH^-$.
  • In pure water, concentrations are equal: $1 \times 10^{-7}\,\text{moles/L}$ (or $0.0000001\,\text{mol/L}$).

• Acids and Bases:

  • Acids: Proton donors. They release $H^+$ when dissolved in water. Example: $HCl \rightarrow H^+ + Cl^-$.
  • Bases: Proton acceptors. Most release a hydroxide ion ($OH^-$) and a cation. Example: $NaOH \rightarrow Na^+ + OH^-$. Some, like Ammonia ($NH_3$), accept protons directly to produce $NH_4^+$ and release $OH^-$.

• The pH Scale:

  • pH is the measure of hydrogen ion concentration ("potential of Hydrogen").
  • Measured from $0$ to $14$ based on logarithms.
  • Neutral: pH $7$ (e.g., pure water).
  • Acidic: pH $< 7$. pH $0$ represents $1\,M\,HCl$.
  • Basic (Alkaline): pH $> 7$. pH $14$ represents $1\,M\,NaOH$.
  • Each whole unit represents a ten-fold ($10\times$) difference in $H^+$ concentration. For example, pH $4$ is $10\times$ more acidic than pH $5$, and pH $11$ is $1000\times$ more basic than pH $8$.

Buffers and Environmental Impact

• Buffers: Substances that compensate for pH changes by absorbing or releasing $H^+$. Most are weak acids or bases that dissociate reversibly.

• Carbonic Acid–Bicarbonate Buffer System:

  • Buffers blood pH. Carbonic acid ($H_2CO_3$) dissociates into bicarbonate (${HCO_3}^-$) and $H^+$.
  • Equation: $H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$.
  • Under acidic conditions: $H^+ + HCO_3^- \rightarrow H_2CO_3$.
  • Under alkaline conditions: $OH^- + H_2CO_3 \rightarrow HCO_3^- + H_2O$.
  • Normal blood pH is $7.4$.

• Ocean Acidification:

  • Approximately $33\%$ of human-generated $CO_2$ is absorbed by oceans.
  • $CO_2 + H_2O \rightarrow H_2CO_3$.
  • Increased $H^+$ combines with carbonate ions ($CO_3^{2-}$) to produce bicarbonate ($HCO_3^-$), depleting the carbonate required for calcification by corals and other marine organisms.

• Acid Precipitation:

  • Burned fossil fuels release sulfur and nitrogen oxides.
  • These react with atmospheric water to form acid rain, fog, or snow with a pH lower than $5.2$.
  • It can sicken wildlife, kill plants/trees, and cause human respiratory diseases like bronchitis and asthma.