Chemistry for Biology: Atoms, Molecules, Bonding, and Water (Lecture Notes)

Session setup, approach, and big picture

• Today and upcoming sessions: the instructor will answer questions and, starting next Wednesday, will close the session to Q&A at that time; no changes after that point. He emphasizes keeping overall understanding in view and staying engaged during the session.
• Teaching format for each lecture:
  • Start with a document camera (doc cam) on one screen, possibly on both screens at the start.
  • When discussing the lecture and the webcast questions, he will switch to the document view.
  • The purpose is to give a big-picture view of the topic for that day.
  • Each lecture will try to be brief, with opportunities to ask questions.
• Topic of today: Chemistry for biology (Chemistry applied to biology).
• Overall trajectory: From atoms and molecules to macromolecules (proteins, DNA, RNA) to the first living organism (unicellular).
• Goal of the session: Build a foundational understanding of how chemistry underpins biology.

Biological hierarchy recap (from last lecture)

• Atoms compose everything in the universe.
• Atoms combine to form molecules.
• Molecules combine to form macromolecules such as proteins, DNA, RNA.
• Macromolecules organize to form cells, leading to the first living unicellular organism.

The Atom: atomic structure and models

• The instructor emphasizes using the electron orbital model (electronic structure) rather than a simple static model.
• Example used: Helium (He):
  • 2 electrons in the first (and only) electron shell.
  • 2 protons (positive charge) and 2 neutrons (no charge) in the nucleus.
• Key terms and units:
  • Protons define the element (the identity of the element).
  • Neutrons define isotopes (same protons, different neutrons).
  • Electrons determine ions and overall reactivity of the atom (how it will interact with others).
  • Atomic Mass Unit (AMU) and Dalton (Da): common mass unit in biology; AMU is the mass scale used for subatomic particles; Dalton is another name for AMU.
• Important conceptual points:
  • Protons = element identity; Neutrons = isotopes; Electrons = ions and reactivity.
  • Reactivity is largely governed by electrons, especially those in the outermost shell (the valence shell).
• The octet rule is introduced as a driving idea for chemical stability (atoms gain, lose, or share electrons to achieve a full valence shell).
• The valence shell is the outermost electron shell; its occupancy largely determines how the atom behaves chemically.
• A common analogy used: the nucleus is the center of a football field; electrons revolve around the nucleus and spectators sit in the surrounding seats (electrons in orbit around the nucleus).
• Electron shell capacities mentioned:
  • First shell: 2 electrons.
  • Second shell and higher: up to 8 electrons per shell—this arrangement tends toward stability.
• Orbitals in space concept:
  • s orbitals are spherical.
  • p orbitals are oriented along x, y, and z axes (often described as two-lobed shapes along three perpendicular axes).
  • The statement “two x, two y, and two z” reflects the arrangement of the three p orbitals (each can hold 2 electrons).
• The lecture notes that the exact numerical values for some of these orbital capacities aren’t required to memorize here, as they will be shown on slides.

Electrons, energy, and electronegativity (qualitative overview)

• Electrons contribute to an atom’s reactivity; the energy needed to hold electrons or the energy of the electrons influences how strongly an atom attracts electrons in bonds.
• The instructor lists a rough ordering of electronegativity (note: the order presented may differ from standard conventions):
  • Chlorine appears as highly electronegative in the discussion, greater than oxygen, greater than nitrogen, greater than carbon, greater than hydrogen.
  • The idea conveyed: some elements are strongly electronegative (tend to attract electrons in bonds), while others are weakly electronegative.
• This leads to differences in bond formation and polarity between atoms.

Bonds and interactions: covalent, ionic, and van der Waals forces

• Covalent bonds:
  • Polar covalent bonds: electrons are shared unequally between atoms with differing electronegativities, leading to partial charges (one atom becomes more negative).
  • Nonpolar covalent bonds: electrons are shared more equally, resulting in no significant partial charges.
  • The degree of polarity depends on the electronegativity difference between the bonded atoms.
• Ionic bonds:
  • One atom (often highly electronegative) gains an electron from another atom, creating charged species (ions).
  • Example conceptually discussed: formation of anions via electron transfer; ionic interactions contribute to bonding in many contexts.
• Relation to water and polarity:
  • Polar molecules and ions interact strongly with water; hydrophilic substances tend to be polar or ionic and dissolve in water.
  • Hydrophobic substances tend to be nonpolar and do not interact favorably with water.
  • DNA and hydrogen bonds interact within a polar environment (water) yet also rely on hydrogen bonds for base pairing (A-T and G-C) as described in the next point.
• Hydrogen bonds in biology:
  • DNA double helix features hydrogen bonds between nitrogenous bases (A with T and G with C).
  • These bonds contribute to the stability of the DNA structure in conjunction with the hydrophobic/internal environment of the double helix.
  • The bases themselves are discussed as part of the molecular context where multiple forces (hydrogen bonding, hydrophobic interactions) stabilize the overall structure.

Van der Waals and dispersion forces

• Van der Waals forces and London dispersion forces are weak attractions that arise from proximity and temporary dipoles.
• These forces are generally weaker than covalent and ionic bonds but can be important for molecular packing and interactions in close contact.

Water, hydrophilicity, and hydrophobicity in context

• Water is a polar solvent; it strongly interacts with polar and ionic species.
• Hydrophilic vs. hydrophobic:
  • Hydrophilic substances tend to be polar and interact with water;
  • Hydrophobic substances tend to be nonpolar and avoid water.
• The interplay of polar interactions, hydrogen bonding (e.g., in DNA base pairing), and hydrophobic effects contributes to the stability and behavior of biological molecules in aqueous environments.

Practical notes on the course format and study approach

• Slides are posted on Canvas; students are encouraged to download and study them, and take notes.
• The instructor plans about 10 questions per session, mixing question types (recall, conceptual, application).
• Students are encouraged to discuss with neighbors and raise questions; the instructor will circulate to help.
• The plan is to work through the questions and then review the answers together.
• The instructor emphasizes focusing on big-picture understanding rather than memorizing isolated details.

Key takeaways and connections to broader biology

• Chemistry is the foundation of biology: atoms form molecules, which form macromolecules, which assemble into cells and organisms.
• Understanding atomic structure (protons, neutrons, electrons) explains why atoms bond, how they bond (covalent, ionic, van der Waals), and how polarity and charge influence molecular interactions in water and biological contexts.
• The valence shell and octet rule provide a conceptual framework for predicting bonding patterns and molecule formation.
• Hydrogen bonding, hydrophobic/hydrophilic balance, and van der Waals forces collectively shape the structure and behavior of DNA, proteins, membranes, and other biomolecules in aqueous environments.

Quick reference to concepts mentioned (for study)

• Atomic mass unit (AMU) and Dalton (Da) are used to express atomic/molecular masses.
• Protons define the element; neutrons define isotopes; electrons determine ions and reactivity.
• Valence shell: outermost electron shell; governs bonding behavior.
• First electron shell capacity: 2 electrons; second and higher shells typically hold up to 8 electrons per shell.
• Octet rule: atoms tend to gain, lose, or share electrons to achieve a full valence shell.
• Covalent bonds: polar vs nonpolar; polarity depends on electronegativity differences.
• Ionic bonds: electron transfer creates ions; strong interactions in polar contexts.
• Hydrogen bonds: specific case of polar interactions relevant in DNA base pairing.
• Hydrophilic vs hydrophobic: interactions with water influence solubility and molecular assembly.
• Van der Waals and London dispersion forces: weak, proximity-driven attractions.
• DNA base pairing: A–T and G–C hydrogen bonding stabilizes the double helix.

End of notes