Comprehensive Introduction to General Chemistry and Atomic Structure for Biology

Supplemental Instruction and Course Logistics

• Supplemental Instructor (SI) Support:

  • An announcement regarding Supplemental Instruction is available on Canvas.

  • The SI holds $2$ study sessions per week:

    • Session 1: Held regularly on Sunday.

    • Session 2: Scheduled on either Monday or Tuesday, determined by student voting via a link provided in the Canvas announcement.

• Instructor Availability and Communication:

  • Questions regarding chemistry or course material are welcomed during lecture, office hours, or via email.

Scientific Questions and Falsifiability

• Falsifiability:

  • The core criterion for evaluating scientific questions, hypotheses, and theories is falsifiability rather than absolute correctness.

  • A robust hypothesis or theory must generate clear, testable, and falsifiable predictions that can be independently evaluated and verified or disproved by others.

  • Falsifiability forms the primary framework for conceptualizing scientific inquiries within the lecture and laboratory components of the course.

Fundamental Atomic Structure and Subatomic Particles

• Etymology and Early Atomic Theory:

  • The word atom originates from the Greek term meaning "uncuttable" or "indivisible."

  • Historically, early natural philosophers postulated that dividing matter continuously would eventually yield a fundamental particle that could not be reduced or divided further.

  • Early chemical investigators originally assumed that elements such as hydrogen ($\text{H}$), carbon ($\text{C}$), and phosphorus ($\text{P}$) represented these indivisible fundamental particles.

• Subatomic Structure:

  • Atoms are composed of smaller subatomic particles organized into a central core and surrounding regions.

  • The Nucleus:

    • Positions at the core of the atom, analogous to the position of the sun at the center of the solar system.

    • Contains two primary subatomic particles:

      • Protons: Subatomic particles possessing a positive electrical charge ($+1$).

      • Neutrons: Subatomic particles that are electrically neutral ($0$ charge).

  • Electrons:

    • Negatively charged subatomic particles ($-1$) surrounding the nucleus.

    • Organized and distributed within defined regions described as electron shells or electron orbitals.

    • Possess negligible, inherently almost non-existent mass relative to protons and neutrons.

• Atomic Mass and Atomic Number:

  • Protons and Neutrons: Massive particles ("chunky") that account for virtually all of an atom's mass.

  • Atomic Mass: Defined mathematically as the sum of protons and neutrons within the nucleus:         $\text{Atomic Mass} = \text{Protons} + \text{Neutrons}$

  • Atomic Number: Defined strictly as the total number of protons in the nucleus:         $\text{Atomic Number} = \text{Protons}$

  • An element's atomic number is immutable; changing the number of protons fundamentally changes the identity of the element.

  • Atoms can alter their number of electrons (forming ions) and neutrons (forming isotopes), but cannot alter proton count without altering elemental identity.

  • Hydrogen Example: Hydrogen has an atomic number of $1$ ($1$ proton). Its standard reported atomic mass of $1.0079$ represents a weighted average of all naturally occurring hydrogen isotopes.

The Periodic Table of Elements and Biological Composition

• Organization of the Periodic Table:

  • Categorizes and structures all known atoms in the universe based on atomic configuration and chemical behavior.

  • Groups (Columns):

    • Elements arranged within the same column exhibit identical or nearly identical chemical properties and reactivity.

    • Column position corresponds directly to the number of electrons residing in an atom's outermost electron orbital (valence shell).

    • Column $1$ elements possess $1$ electron in their outer orbital.

    • Column $2$ elements possess $2$ electrons in their outer orbital.

    • Columns $3$ through $8$ possess $3$, $4$, $5$, $6$, $7$, and $8$ outer electrons, respectively.

• Elemental Composition of Biological Organisms:

  • Major Biological Elements ($98\%$ of biological mass):

    • Hydrogen ($\text{H}$)

    • Carbon ($\text{C}$)

    • Nitrogen ($\text{N}$)

    • Oxygen ($\text{O}$)

  • Minor Biological Elements ($2\%$ of biological mass):

    • Comprises essential minerals, salts, and specific reactive atoms.

    • Sodium ($\text{Na}$)

    • Potassium ($\text{K}$)

    • Magnesium ($\text{Mg}$)

    • Sulfur ($\text{S}$)

    • Phosphorus ($\text{P}$)

Atomic Depictions and the Bohr Model

• The Bohr Model:

  • Developed historically by Niels Bohr, depicting atoms as a central nucleus surrounded by concentric circular electron shells.

  • While not a completely accurate physical representation of quantum electron behavior, it serves as an indispensable tool in biology for predicting chemical behavior and bonding capacity.

• Rules for Drawing Bohr Models:

  1. Represent the central nucleus with a positive charge equal to the total proton count.

  2. Fill electron shells sequentially outward from the nucleus.

  3. Innermost Shell Capacity: Holds a maximum of $2$ electrons. If an atom has $3$ or more electrons, additional electrons must occupy outer shells.

  4. Subsequent Shell Capacities: Second and outer shells hold a maximum of $8$ electrons.

The Octet Rule, Chemical Reactivity, and Ion Formation

• The Octet Rule:

  • Atoms dynamically gain, lose, or share electrons to achieve a completely filled outer electron shell.

  • For most valence shells, stability is achieved when $8$ electrons reside in the outermost shell.

• Inert Elements (Noble Gases):

  • Elements born with naturally complete outer valence shells ($8$ valence electrons, or $2$ for Helium).

  • Do not react chemically with other atoms; existing in a state of maximum stability as inert monoatomic gases.

  • Examples include Helium ($\text{He}$), Neon ($\text{Ne}$), Argon ($\text{Ar}$), and Krypton ($\text{Kr}$).

• Ion Formation and Reactivity:

  • Atoms lacking a complete valence shell display chemical reactivity to resolve their electron excess or deficiency.

  • Case Study: Sodium ($\text{Na}$):

    • Sodium possesses a single electron in its outermost shell.

    • To reach a full octet, sodium can either gain $7$ electrons or lose $1$ electron.

    • Losing $1$ electron is energetic and physically favorable because the sodium nucleus lacks sufficient electrostatic pull to hold $7$ additional electrons.

    • Upon losing its outer electron, sodium achieves stability as a positively charged cation ($\text{Na}^+$) with a $+1$ net charge.

  • Ionic Attraction: Positively charged cations ($\text{Na}^+$) electrostatically interact with negatively charged anions, such as chloride ($\text{Cl}^-$), to form neutral ionic complexes.

Isotopes, Carbon-14, and Radiometric Dating

• Isotopes:

  • Variants of the same chemical element that contain an identical number of protons but differing numbers of neutrons.

  • Represent distinct structural "flavors" of a single element.

• Carbon Isotopes:

  • All carbon atoms possess exactly $6$ protons by definition.

  • Carbon-12 ($^{12}\text{C}$):

    • Composed of $6$ protons and $6$ neutrons ($6 + 6 = 12$).

    • Makes up greater than $90\%$ of all biological carbon.

  • Carbon-14 ($^{14}\text{C}$):

    • Composed of $6$ protons and $8$ neutrons ($6 + 8 = 14$).

    • Exists in extremely small trace quantities in nature.

• Biospheric Incorporation and Decay of Carbon-14:

  • Cosmic radiation striking the upper atmosphere converts atmospheric nitrogen ($\text{N}$) into unstable $^{14}\text{C}$.

  • Atmospheric $^{14}\text{C}$ is fixed into plant tissues via photosynthesizing organisms.

  • Herbivores (such as deer) ingest plants, incorporating $^{14}\text{C}$ into their tissues; carnivores and humans subsequently ingest these organisms.

  • Living organisms maintain a steady-state equilibrium ratio of $^{14}\text{C}$ through continuous dietary intake.

  • Upon death, uptake ceases, and $^{14}\text{C}$ undergoes predictable radioactive decay over time.

  • Measuring the remaining proportion of $^{14}\text{C}$ in organic matter allows scientists to determine the age of fossils.

Overview of Chemical Bonds and Molecules

• Definitions:

  • Molecule: Any structure composed of two or more atoms held together by chemical bonds.

  • Compound: A molecule containing two or more different chemical elements bonded together (e.g., carbon bonded to hydrogen, or oxygen bonded to nitrogen).

• Four Primary Biological Chemical Bonds:

  1. Covalent Bonds

  2. Hydrogen Bonds

  3. Ionic Bonds

  4. Van der Waals Interactions

Covalent Bonds: Nonpolar vs. Polar

• Mechanisms of Covalent Bonding:

  • Formed when two atoms share valence electrons.

  • Electrons continuously move or "ping-pong" back and forth between the nuclei of both participating atoms.

  • The total number of covalent bonds an atom forms equals the number of additional electrons required to fulfill the octet rule:

    • Carbon ($\text{C}$): Possesses $4$ valence electrons; requires $4$ additional electrons; consistently forms $4$ covalent bonds.

    • Oxygen ($\text{O}$): Possesses $6$ valence electrons; requires $2$ additional electrons; consistently forms $2$ covalent bonds (e.g., two single bonds in $\text{H}_2\text{O}$, or one double bond in $\text{O}_2$).

  • Covalent bonds can manifest as single, double, or triple shared pair interactions and form the essential molecular backbone of living matter.

• Nonpolar Covalent Bonds:

  • Occur when electron density is shared equally between atoms (e.g., diatomic gas molecules like $\text{Cl}_2$ or $\text{O}_2$).

  • Results in a uniform charge distribution across the molecule.

• Polar Covalent Bonds:

  • Occur when one atom exhibits high electronegativity, pulling the shared electron density heavily toward its nucleus.

  • Generates asymmetric charge distribution with partial negative charges ($\delta^-$) near the electronegative atom and partial positive charges ($\delta^+$) near the electron-starved atom.

  • Water ($\text{H}_2\text{O}$) as a Polar Molecule:

    • Oxygen strongly pulls electron density away from both hydrogen atoms.

    • Oxygen carries a permanent partial negative charge ($\delta^-$), while both hydrogens carry partial positive charges ($\delta^+$).

• Incompatibility of Oil and Water:

  • Water molecules are polar and behave like microscopic magnets, adhering strongly to one another.

  • Oils are completely nonpolar molecules lacking charge asymmetry.

  • Polar water molecules associate so tightly with one another that nonpolar oil molecules are physically excluded, preventing mixing.

Hydrogen Bonding

• Mechanism:

  • Requires a molecule containing a polar covalent bond between a hydrogen atom and a strongly electronegative atom.

  • Electronegative atoms—specifically Oxygen ($\text{O}$), Nitrogen ($\text{N}$), and Fluorine ($\text{F}$)—draw electron density away, leaving the bonded hydrogen with a strong partial positive charge ($\delta^+$).

  • The electron-starved hydrogen atom interacts electrostatically with the partial negative charge ($\delta^-$) on an electronegative atom ($\text{O}$, $\text{N}$, or $\text{F}$) of an adjacent molecule.

  • Common molecular motifs exhibiting hydrogen bonding include $\text{O}-\text{H}$, $\text{N}-\text{H}$, and $\text{F}-\text{H}$.

• Biological Significance:

  • Individual hydrogen bonds are weak and transient, but collectively exert powerful structural effects.

  • Gives liquid water its high surface tension.

  • Stabilizes the structural integrity of the DNA double helix; individual base pairs are held together by hydrogen bonds that are weak enough to be unzipped during replication/transcription, yet stable en masse.

Ionic Bonds and Water Dissolution

• Mechanism of Ionic Bonding:

  • Occurs via complete electron transfer from a donor atom to a recipient atom, rather than electron sharing.

  • The donor becomes a positively charged cation; the recipient becomes a negatively charged anion.

  • Oppositely charged ions adhere tightly via strong electrostatic attraction, functioning like opposing magnetic poles (e.g., $\text{Na}^+$ and $\text{Cl}^-$ forming table salt, $\text{NaCl}$).

  • In a dry state, ionic bonds are exceptionally strong, surpassing standard covalent bonds.

• Mechanism of Dissolution in Water:

  • Water rapidly disrupts ionic crystalline lattices.

  • The partial negative oxygen regions of water molecules surround positively charged cations ($\text{Na}^+$).

  • The partial positive hydrogen regions of water surround negatively charged anions ($\text{Cl}^-$).

  • Water isolates individual ions, encapsulating them in hydration shells and carrying them away, thereby dissolving the compound.

Van der Waals Interactions and Biological Applications

• Mechanism:

  • The weakest, most transient, and abstract type of chemical interaction.

  • Driven by short-lived, transient fluctuations in electron density within nonpolar or neutral atoms.

  • At any given instant, random electron movement creates a temporary, localized concentration of negative charge on one side of an atom, establishing a momentary induced dipole.

  • This momentary dipole induces a complementary instantaneous dipole in adjacent surrounding atoms, leading to a brief, weak magnetic attraction.

  • These interactions persist for only a tiny fraction of a second before breaking and dynamically reforming.

• Biological Case Study: Gecko Locomotion:

  • Geckos adhere to vertical walls and ceilings without employing suction cups, chemical adhesives, or sticky secretions.

  • Microscopic structures on gecko feet come into close contact with surface molecules of a wall.

  • Physical proximity induces instantaneous Van der Waals dipoles between the electrons in the gecko's foot structures and the electrons in the wall surface.

  • Every step generates millions of transient, additive magnetic interactions, sticking the foot to the wall dynamically before releasing seamlessly for the next step.