Module 1 Notes: Foundations of Chemistry for Biology

Elements and Atoms

  • An element is a fundamental substance that cannot be chemically broken down into something else.
  • Periodic table abbreviations: H, He, Li, C, N, O, etc. Examples of elements discussed: carbon (C), nitrogen (N), hydrogen (H), oxygen (O).
  • An atom represents an element and has a specific structure consisting of subatomic particles: electrons, protons, and neutrons.
  • Subatomic particles and charges:
    • Electron: symbolized as e⁻, negative charge.
    • Proton: positive charge.
    • Neutron: neutral charge.
  • The nucleus contains protons and neutrons; electrons orbit the nucleus.
  • Hydrogen example: 1 proton in the nucleus and 1 electron in orbit; hydrogen has no neutrons in its most common form.
  • Helium example: 2 protons, 2 neutrons in the nucleus; 2 electrons orbiting around; helium is not a biological atom because its outer shell is already filled and stable.
  • The nucleus is extremely tiny compared to the overall atom.
  • As you move across the periodic table, shells (electronic energy levels) become more complex; electrons occupy shells around the nucleus.
  • Common shell capacity pattern described:
    • First shell can hold a maximum of 2 electrons.
    • Second shell can hold a maximum of 8 electrons.
    • A third shell (if present) can hold a maximum of 8 electrons.
  • The number of shells increases deeper in the table; the outermost shell is called the valence shell.
  • A brief note on hydrogen ions: a proton alone can be called a hydrogen ion (H⁺). An ion is a charged particle (positive or negative). Sometimes this is written as H⁺ or as a proton alone; terminology can switch between "proton," "hydrogen ion," and symbols.
  • The arrangement of electrons and their shells determines how atoms interact with others to form bonds and molecules.

Electron Shells and Valence

  • Outer shell (valence shell) governs chemical behavior; when not full, atoms tend to bond with others.
  • In the first shell, a maximum of 2 electrons can be accommodated; once full, bonding involves the next shell up to its capacity (commonly 8 in the second and subsequent shells in many atoms).
  • An atom’s chemical behavior is influenced by the number of electrons in its valence shell, which tends to be similar for atoms in the same column of the periodic table.
  • Covalent bonding arises from sharing of electrons between atoms to satisfy outer-shell electron requirements.
  • Types of covalent bonds (concepts emphasized):
    • Single bond: sharing of one electron pair between two atoms (one from each atom).
    • Double bond: sharing of two electron pairs between two atoms.
  • Valence and bond formation examples:
    • Hydrogen (H) has one electron and can form a single covalent bond to complete its first shell (2 electrons total).
    • Oxygen (O) has six valence electrons in its outer shell and seeks two more to reach eight; two oxygens can share two pairs (O=O) to satisfy octets.
  • The outermost shell, when not filled, drives bonding and molecule formation.

Ions and Ionic Bonding

  • Ion: a charged particle (positive or negative).
  • Hydrogen ion: H⁺ is essentially a bare proton.
  • Ionic bonding involves complete transfer of electrons from one atom to another, not sharing.
  • Example: Sodium chloride (NaCl).
    • Sodium (Na) has one electron in its outer shell and can give it up to achieve a stable configuration, becoming Na⁺ (cation).
    • Chlorine (Cl) has seven electrons in its outer shell and needs one more to complete its octet, gaining an electron to become Cl⁻ (anion).
    • In the dry state, Na⁺ and Cl⁻ attract each other to form an ionic lattice (NaCl).
    • In water, this ionic bond is weakened (solvation/heating and dielectric effects) and Na⁺ and Cl⁻ become hydrated and dispersed.
  • Notation you’ll see:
    • Sodium giving up an electron: ext{Na}
      ightarrow ext{Na}^+ + e^-
    • Chlorine gaining an electron: ext{Cl} + e^-
      ightarrow ext{Cl}^-
  • Positive ion = cation; Negative ion = anion.
  • Ionic interactions are strong electrostatic attractions between oppositely charged ions, but they differ from covalent bonds in that electrons are not shared between atoms.

Covalent Bonding: Polar and Nonpolar

  • Covalent bonding involves sharing electrons between atoms.
  • A single covalent bond corresponds to one shared pair of electrons; a double bond corresponds to two shared pairs, etc.
  • Nonpolar covalent bond:
    • Electrons are shared relatively equally between the two atoms.
    • Example: O=O (oxygen-oxygen double bond) and nonpolar H–H (single bonds) where there is no significant charge separation.
  • Polar covalent bond:
    • Electrons are not shared equally due to differences in electronegativity, leading to partial charges on atoms.
    • Example: Water (H₂O) has an electronegative oxygen pulling electron density toward itself, creating a partial negative charge on O and partial positive charges on the hydrogens.
    • In water, the bond is often described as polar covalent, with electrons spending more time near the oxygen nucleus and less near the hydrogens.
  • Polar covalent bonds introduce directionality and partial charges (δ⁺, δ⁻ or σ as a partial charge indicator) that influence interactions with other molecules.
  • Important illustrative examples:
    • Hydrogen molecule (H–H): both hydrogens share one electron each, resulting in a nonpolar covalent bond.
    • Oxygen molecule (O=O): shares two electrons (double bond) in a nonpolar covalent manner.
    • Water molecule (H–O–H): two polar covalent bonds with a bent geometry; oxygen holds partial negative charge, hydrogens partial positive charges.
  • Methane (CH₄) as a simple hydrocarbon:
    • Carbon has four valence electrons and forms four single covalent bonds with four hydrogens, sharing one electron with each hydrogen to fill carbon’s valence shell.
    • Methane is a nonpolar molecule with a tetrahedral arrangement around the carbon.
  • Molecular representations:
    • A bond is depicted as a line between atoms (single bond = one line; double bond = two lines).
    • Molecules can have three-dimensional geometry; methane is shown with carbon in the center and four hydrogens around it in a three-dimensional arrangement.
  • In covalent bonding, equal sharing yields nonpolar covalent bonds; unequal sharing yields polar covalent bonds.
  • A note on molecular polarity:
    • If a molecule has regions with stronger electron pull toward certain atoms (due to electronegativity), the molecule is polar.
    • Lipids (fats, oils, waxes) are largely nonpolar and do not mix with water (nonpolar–polar immiscibility).
  • Water as a key polar solvent: polarity drives many interactions in biology and chemistry.

Hydrogen Bonding and DNA

  • Hydrogen bonds: not covalent bonds; they do not involve electron sharing or electron transfer.
  • They are attractions between a positively charged region of one molecule (typically a partially positive hydrogen, δ⁺) and a negatively charged region of another molecule (often a lone pair on O, N, or similar, with δ⁻).
  • Hydrogen bonds are weaker than covalent bonds but are critical for the structure of large biomolecules.
  • DNA example: double helix held together by hydrogen bonds between complementary bases on the two strands.
    • Adenine (A) pairs with Thymine (T); Guanine (G) pairs with Cytosine (C).
    • Base pairing is driven by hydrogen-bonding patterns (A–T two H-bonds; G–C three H-bonds in standard models), stabilizing the double helix.
  • Hydrogen bonding also explains intermolecular interactions like hydrogen bonds between water and ammonia (NH₃) or water and itself, contributing to water’s high cohesion.
  • In summary:
    • Covalent bonds involve electron sharing (and sometimes electronegativity-driven polarity).
    • Hydrogen bonds involve attractions between partial charges without electron sharing.

Water, Ammonia, and Polarity in Molecules

  • Oxygen and nitrogen are highly electronegative and pull electron density toward themselves in compounds like water (H₂O) and ammonia (NH₃).
  • Water is a polar molecule: the oxygen end is partially negative (δ⁻) and the hydrogen ends are partially positive (δ⁺).
  • Ammonia can also form hydrogen bonds with water and exhibits polar covalent bonds between N–H.
  • Polar vs nonpolar interactions:
    • Polar molecules tend to interact with other polar molecules (soluble in each other).
    • Nonpolar molecules tend to interact with nonpolar molecules and resist mixing with water (oil and water example).
  • Oil and vinegar demonstration: oil (nonpolar) does not dissolve in water (polar) and tends to separate, though shaking can temporarily mix them into droplets due to dispersion.
  • The concept of electronegativity describes an atom’s attraction for electrons within bonds; higher electronegativity means stronger pull on shared electrons.

Chemical Reactions: Bonds Broken and Bonds Made

  • Chemical reactions involve making and breaking bonds, not creating or destroying atoms (atoms are conserved).
  • A simple example: formation of water from hydrogen and oxygen.
    • Reactants: 2 \, ext{H}2 + ext{O}2
    • Products: 2 \, ext{H}_2 ext{O}
    • Balanced equation: 2\ \mathrm{H2} + \mathrm{O2} \rightarrow 2\ \mathrm{H_2O}
  • Explanation of the process:
    • Two H–H single bonds are broken and O=O double bonds are reorganized into two O–H bonds per water molecule.
    • No atoms are gained or lost; bonds are rearranged.
  • The line notation in drawings represents bonds; a single bond is a single line, a double bond is two lines, etc.
  • The concept of reactants and products helps in understanding chemical changes and the pathways of metabolism in biology.

Nonpolar vs Polar Molecules and Solubility in Biology

  • Lipids (fats, oils, waxes) are largely nonpolar due to C–H and C–C bonds; they do not mix with polar solvents like water.
  • This polarity difference is essential for cell membranes:
    • Lipid bilayers form barriers that create distinct intracellular and extracellular environments.
    • The hydrophobic tails face inward, away from water, while polar (or charged) head groups face the aqueous surroundings.
  • The “like dissolves like” principle helps explain solubility and membrane structure in biology.

About Molecules and Three-Dimensional Structure

  • A molecule is a stable assembly of at least two atoms bonded together.
  • Even simple molecules like methane (CH₄) demonstrate three-dimensional geometry (tetrahedral arrangement around carbon).
  • The three-dimensional arrangement matters for interactions in biomolecules (proteins, nucleic acids, membranes) and for understanding how molecular interactions occur inside cells.

Module 1: Living Systems – Connections to Food and Metabolism

  • Theme: food is central to life because it provides energy and nutrients essential for cellular processes.
  • Metabolism: the set of life-sustaining chemical reactions in organisms, including two broad processes:
    • Catabolism: breaking down substances into smaller units to release energy.
    • Anabolism: building larger molecules from smaller subunits, using energy.
  • Monomers and Macromolecules:
    • Small subunits (monomers) join to form larger molecules (macromolecules).
    • Macromolecules include carbohydrates, proteins, lipids, and nucleic acids (DNA, RNA).
  • Hierarchical organization in biology:
    • Atoms combine to form molecules.
    • Molecules combine to form macromolecules (polymers).
    • Macromolecules function within organelles (e.g., mitochondria).
    • Cells are organized into tissues and organs, forming organ systems in an organism.
  • Rough cellular biology gist mentioned for later chapters:
    • Cells (~30 trillion in the human body) arise from atoms forming molecules, which assemble into macromolecules, organize into organelles, then cells, tissues, organs, and organ systems.
  • Practical course structure mentioned:
    • Each module begins with learning objectives (a checklist of what to know for tests).
    • Homework is designed to reinforce concepts and prepare for exams; early modules have combined due dates (Module 1 and 2 due on the same day, Aug 31 at 11:59 PM).
    • Repetition helps; attempting the same homework multiple times exposes you to different questions from a larger test bank.
  • Recap of essential takeaways for Module 1:
    • Basic chemistry foundations (elements, atoms, electrons, nuclei, shells, valence).
    • Bonding types (covalent, polar covalent, nonpolar covalent, ionic, hydrogen bonds).
    • Polarity and its biological relevance (solubility, membranes, DNA structure).
    • Core examples: H₂, O₂, H₂O, CH₄, NaCl, the role of water in biology, and how metabolic processes connect chemistry to life.
  • Next steps mentioned:
    • Continue into Module 2 with deeper biological topics while continuing to reference these chemistry basics.

Quick Reference: Key Terms and Concepts

  • Element, Atom, Nucleus, Proton, Neutron, Electron
  • Shells, Valence Shell, Octet Rule, Covalent Bond, Single Bond, Double Bond
  • Nonpolar Covalent Bond, Polar Covalent Bond, Electronegativity, Partial Charge (σ or δ⁺/δ⁻)
  • Ionic Bond, Ion, Cation, Anion
  • Hydrogen Bond, DNA Base Pairing (A–T, G–C)
  • Molecule, Methane (CH₄), Water (H₂O), Ammonia (NH₃)
  • Solubility, Polarity vs Nonpolarity, Hydrophobic vs Hydrophilic
  • Reactants, Products, Chemical Reactions, Conservation of Atoms
  • Metabolism, Catabolism, Anabolism, Macromolecules, Organelles, Cells, Tissues, Organ Systems
  • Practical implications: cell membranes, dissolution in water, and the role of polarity in biology