Bio review

Chemistry concepts

PH

  • The pH scale measures how acidic or basic a solution is, ranging from 0 to 14.

    • pH less than 7 → Acidic

    • pH = 7 → Neutral (pure water)

    • pH more then 7 → Basic/Alkaline (more OH⁻ ions)

  • Each pH unit represents a tenfold difference in hydrogen ion concentration.

    • Example: pH 3 is 10 times more acidic than pH 4, and 100 times more acidic than pH 5.

Water Molecule (H₂O)

  • Polarity: Water is a polar molecule because oxygen is more electronegative than hydrogen, creating a partial negative charge (δ⁻) on oxygen and partial positive charges (δ⁺) on hydrogens.

  • Polar bonds: The O-H bonds in water are polar covalent bonds due to the unequal sharing of electrons.

  • Hydrogen bonding: Water molecules form hydrogen bonds (weak attractions) between the δ⁺ hydrogen of one molecule and the δ⁻ oxygen of another. These bonds give water its unique properties like cohesion, high specific heat, and solvent abilities.

Buffer

  • A buffer is a solution that resists changes in pH when acids or bases are added.

  • Buffers help maintain stable pH in biological systems.

    • Example: The bicarbonate (HCO₃⁻) buffer system in blood maintains pH around 7.4.

Polar vs Non-Polar Bonds/Molecules

  • Polar bonds/molecules:

    • Unequal sharing of electrons due to differences in electronegativity.

    • Molecules have partial charges (δ⁺ and δ⁻).

    • Interact with water (hydrophilic) because "like dissolves like."

    • Example: Water (H₂O), ammonia (NH₃), glucose (C₆H₁₂O₆).

  • Non-polar bonds/molecules:

    • Equal sharing of electrons (or very small difference in electronegativity).

    • No significant charge separation.

    • Do not mix well with water (hydrophobic).

    • Example: Oxygen gas (O₂), methane (CH₄), fats, and oils.

Ionic vs Covalent Bonds

  • Ionic Bonds:

    • Formed when electrons are transferred from one atom to another.

    • Creates charged ions: cations (+) and anions (-).

    • Strong in solid form but can dissolve in water (ex: NaCl (table salt)).

  • Covalent Bonds:

    • Formed when electrons are shared between atoms.

    • Can be polar (unequal sharing) or non-polar (equal sharing).

    • Examples: H₂O (polar covalent), CH₄ (non-polar covalent).

Electronegativity

  • Definition: The ability of an atom to attract electrons in a bond.

  • Determining bond type:

    • Large difference (>1.7) → Ionic bond (ex: NaCl)

    • Moderate difference (0.5–1.7) → Polar covalent bond (ex: H₂O)

    • Small or no difference (<0.5) → Non-polar covalent bond (ex: O₂, CH₄)

Intermolecular Bonds (Simplified)

a) London Forces (Dispersion Forces)

  • What it is: Weak attraction caused by temporary shifts in electron position, creating short-lived dipoles.

  • Where it happens: In all molecules but is the only force in nonpolar molecules (like gases and oils).

  • Example: Methane (CH₄) or argon (Ar).

b) Dipole-Dipole Interactions

  • What it is: Attraction between molecules with permanent positive and negative ends.

  • Where it happens: Polar molecules (those with an uneven charge distribution).

  • Example: Hydrogen chloride (HCl)—the positive hydrogen in one HCl attracts the negative chlorine in another.

c) Hydrogen Bonding

  • What it is: A strong dipole-dipole force where hydrogen bonds with fluorine (F), oxygen (O), or nitrogen (N).

  • Where it happens: In molecules with H—F, H—O, or H—N bonds.

  • Example: Water (H₂O)—oxygen from one molecule attracts hydrogen from another, making water sticky!