Water: Properties and Polarity

• Water covers more than 70% of Earth's surface
• Water enables life as a solvent and substrate for cellular reactions
• Water transports chemicals within biological systems
• Water helps maintain constant body temperature
• Cellular components and molecules assume their shape due to water's influence

Molecular Structure of Water

• Oxygen in water has sp3 hybridized orbitals
• The water molecule is bent with an H-O-H angle of 104.5°
• The bent structure of water makes it polar

Polarity and Dipoles

• Polarity difference between oxygen and hydrogen, makes water a polar molecule
• Oxygen is more electronegative than hydrogen
• Charge separation in water creates partial negative charge on oxygen and partial positive charges on hydrogen
• Molecules with separated charges are called dipoles
• Dipoles orient opposite to an external electric field

Molecular Dipoles

• Molecules with separated charges are called dipoles
• Molecular dipoles will orient themselves in the direction opposite to the field when subjected to an electric field

Noncovalent Bonding

• Noncovalent bonds are usually electrostatic
• Noncovalent bonding occurs between the positive nucleus of one atom and the negative electron clouds of another nearby atom
• Noncovalent bonds are relatively weak and easily disrupted
• A large number of noncovalent interactions stabilize macromolecules

Types of Noncovalent Interactions in Water

• Ionic interactions occur between charged atoms or groups, such as NaCl dissociation in water
• Ionic interactions are important in protein structures, forming salt bridges
• Hydrogen bonding is a combination of electrostatic and covalent character
• Water molecules form up to 4 hydrogen bonds, leading to high boiling and melting points
  • Two hydrogen bonds occur via hydrogens
  • Two hydrogen bonds occur via nonbonding electron pairs on oxygen
• Van der Waals interactions range from 0.3–9 kJ/mol
• Hydrophobic interactions range from 3-12 kJ/mol

Van der Waals Forces

• Dipole-Dipole interactions occur between molecules with permanent dipoles
• Dipole-Induced Dipole happens when a permanent dipole induces a dipole in a nonpolar molecule
• Induced Dipole-Induced Dipole (London Dispersion Forces) describes weak interactions due to temporary dipoles, like in DNA base stacking

Thermal Properties of Water

• High heat capacity of water helps in temperature regulation
• High heat of vaporization requires significant energy for phase change, aiding cooling mechanisms
• Thermal regulation is essential in biological systems

Strength of Interactions

• Covalent bonds have a strength of >210 kJ/mol
• Ionic interactions range from 4–80 kJ/mol
• Hydrogen bonds range from 12–30 kJ/mol

Hydrophilic Molecules

• Hydrophilic molecules are ionic or polar substances with an affinity for water
• In Greek, "Hydro" means "water," and "philos" means "loving"
• Water's dipole structure and ability to form hydrogen bonds with electronegative atoms enable it to dissolve ionic and polar substances
• Hydrophilic substances are soluble in water due to ion-dipole interactions, dipole-dipole interactions, and hydrogen bonding

Solvent Properties of Water

• Water is an ideal biological solvent, easily dissolving a wide variety of biological molecules
• Some substances do not dissolve in water, demonstrating hydrophilic and hydrophobic properties
  • Salts (e.g., KCl, NaCl) are held together by ionic interactions

Dissolving in Water

• When dissolved, ions separate because polar water molecules attract them more strongly than they attract each other (ion-dipole interaction)
• Water molecules form solvation spheres around the ions

Hydrophilic and Hydrophobic Molecules

• Hydrophilic molecules are ionic or polar substances that dissolve in water via ion-dipole, dipole-dipole, and hydrogen bonding interactions
• Hydrophobic molecules are nonpolar substances that do not dissolve in water; water forms hydrogen-bonded cages around them

Dipole-Dipole Interactions

• Dipole-dipole interactions occur in organic molecules with ionizable groups
• The polar water molecule interacts with carboxyl groups of aldehydes, ketones (carbohydrates), and hydroxyl groups of alcohols
• Water forms hydrogen-bonded cage-like structures around hydrophobic molecules, forcing them out of solution (droplet formation or phase separation)

Amphipathic Molecules

• Amphipathic molecules contain both hydrophilic and hydrophobic regions, forming micelles in water
• Amphipathic molecules contain both polar (hydrophilic) and nonpolar (hydrophobic) regions
• Ionized fatty acids are amphipathic, with a water-soluble carboxylate group and a hydrophobic long carbon chain
• Amphipathic molecules form micelles in water with polar heads outward and nonpolar tails inward

Hydrogen Bonding

• A hydrogen atom attached to an oxygen or nitrogen atom becomes highly polarized with a partial positive charge
• This charge interacts with nonbonding electrons on another oxygen or nitrogen atom, forming strong hydrogen bonds
• Polar heads of molecules orient toward water, while nonpolar tails aggregate away from water

Osmotic Pressure

• Osmosis is the movement of water through a semipermeable membrane from low to high solute concentration
• Osmotic pressure (π) is defined by the equation π = iMRT
  • i = Van't Hoff factor
  • M = Molarity (mol/L)
  • R = Gas constant (0.082 L·atm/K·mol)
  • T = Temperature in Kelvin
• Osmolarity = iM (osmol/Liter)

Osmotic Pressure and Its Importance in Cells

• Osmotic pressure is a crucial factor affecting cellular function
• Cells contain a high concentration of solutes, including small organic molecules, ionic salts, and macromolecules

Water Movement Across Membranes

• Cells may gain water (swelling) or lose water (shrinking) depending on the concentration of solutes in their environment
  • If the external environment is hypotonic, water enters the cell, potentially causing rupture
  • If the external environment is hypertonic, water exits the cell, leading to shrinkage
  • In an isotonic environment, water movement is balanced, and the cell remains stable

Solution Types

• Isotonic: No net water movement (cell remains the same)
• Hypotonic: Water moves into the cell (cell swells and may rupture)
• Hypertonic: Water moves out of the cell (cell shrinks)

Self-Ionization of Water

• The self-ionization of water is represented as: H₂O ⇌ H₃O⁺ + OH⁻

Ion Product Constant of Water (Kw)

• Kw = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴ at 25°C
  • Pure water: [H₃O⁺] = [OH⁻] = 1 × 10⁻⁷ M (neutral pH = 7)

pH Calculation

• pH = -log[H₃O⁺]
• pOH = -log[OH⁻]
• pH + pOH = 14
• Acidic: pH < 7
• Neutral: pH = 7
• Basic: pH > 7

Water Ionization

• Water ionization is the chemical reaction in which two water molecules react to produce a hydronium (H₃O⁺) and a hydroxide (OH⁻) ion
• Water ionization occurs endothermically due to electric field fluctuations between molecules caused by nearby dipole librations from thermal effects and favorable localized hydrogen bonding
• Reaction: H₂O + H₂O ⇌ H₃O⁺ + OH⁻
• Ions may separate but normally recombine within a few minutes to seconds
• At 25°C, water ionizes into equal amounts
  • [H₃O⁺] = [OH⁻] = 1 × 10⁻⁷ M
  • Ion product constant: Kw = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴ M²
• When external acids or bases are added to water, the ion product remains constant at Kw = 1 × 10⁻¹⁴

Bronsted-Lowry Theory

• Acid: Proton (H⁺) donor
• Base: Proton (H⁺) acceptor

Example Reactions

• HCl + H₂O ⇌ H₃O⁺ + Cl⁻ (HCl is the acid, Cl⁻ is the conjugate base)
• NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (NH₃ is the base, NH₄⁺ is the conjugate acid)

pH Scale

• pH = -log[H₃O⁺]
• pH + pOH = 14
• Acidic: pH < 7
• Neutral: pH = 7
• Basic: pH > 7

Examples of pH Values

• Gastric juice: pH 1.5-3
• Blood: pH 7.35-7.45
• Household ammonia: pH 11-12

Strength of Acids and Bases

• Strong acids and bases ionize nearly 100% in water (e.g., HCl, H₂SO₄, NaOH, KOH)
• Weak acids and bases partially ionize (e.g., acetic acid, lactic acid, ammonia)

Dissociation Constant (Ka) and pKa

• Ka = [H₃O⁺][A⁻] / [HA]
• pKa = -log Ka (lower pKa = stronger acid)

Titration

• Titration is a process where measured volumes of a base are added to an acid to determine its concentration. The Equivalence point occurs when moles of acid = moles of base

Henderson-Hasselbalch Equation

• pH = pKa + log ([A⁻]/[HA])
  • Helps predict pH changes in buffer solutions

Buffer Solutions

• A buffer resists pH change when small amounts of strong acid or base are added. Buffers work by:
  • Accepting hydrogen ions when in excess
  • Donating hydrogen ions when depleted.

Physiological Buffers

• Phosphate buffer system: H₂PO₄⁻ / HPO₄²⁻ (intracellular regulation)
• Bicarbonate buffer system: HCO₃⁻ / H₂CO₃ (blood pH regulation)
• Hemoglobin, serum albumins stabilize pH

Example Calculation

• During wine fermentation, a buffer system consisting of tartaric acid and potassium hydrogen tartrate forms
• If [hydrogen tartrate] = 2 × [tartaric acid], and pKa = 2.96: pH = pKa + log(2) = 2.96 + 0.30 = 3.26

Role of polarity in water

• Water has an uneven distribution of charge due to the higher electronegativity of oxygen, causing a partial negative charge near oxygen and partial positive charges near hydrogen

Types of interactions

• Ionic interactions occur between charged atoms (e.g., NaCl)
• Hydrogen interactions are weak electrostatic attraction between H and electronegative atoms
• Van der Waals interactions are weak electrostatic attraction between H and electronegative atoms

Properties of water contributing to dissolving in polar substances

• Polarity and hydrogen bonding allow water to interact with charged and polar molecules, breaking ionic bonds and forming hydration shells

pH scale

• pH measures [H₃O⁺]; pure water at 25°C has pH 7, with [H₃O⁺] = [OH⁻] = 1 × 10⁻⁷ M
• pH = -log(2.6 × 10⁻³) ≈ 2.58

Tonicity

• Hypertonic: Higher solute concentration outside (cells shrink)
• Isotonic: Equal solute concentration (no change)
• Hypotonic: Lower solute concentration outside (cells swell)

Role of buffers

• Buffers resist pH changes by neutralizing added acids/bases
• Bicarbonate buffer (HCO₃⁻/H₂CO₃) maintains blood pH'

Henderson-Hasselbalch equation

• Relates pH to pKa and the ratio of conjugate base to acid
  • It allows precise pH calculations in buffer systems