Lecture Notes: Model Systems, Atoms, Bonds, and Thermodynamics
Model systems and the reductionist approach in molecular biology
Molecular biologists are by nature reductionists: they break complex processes into smallest components to understand how parts work and then assemble them to picture the whole
Emergent properties mean the whole system can behave in ways not predictable from parts alone
Still, starting with parts and players is essential to understand the process
Detailed understanding matters in this course: this class goes deeper than a lightweight intro; expect lots of detail about each process
Post-course, if you only care about large-scale organisms (e.g., giraffes), some details may be less relevant
When building models, start with the simplest model and assume the most efficient behavior unless data say otherwise
In multicellular organisms, what’s most efficient for the organism as a whole may differ from what’s efficient for individual cells
The instructor’s focus: molecular biology in the context of very early animal embryogenesis; sometimes examples differ from bacteria (e.g., E. coli) where the organism is the cell
Living systems obey the physical laws of the universe; cannot dematerialize atoms, create energy, or violate conservation of energy
Every equation must be balanced; no creating/destroying matter or energy inappropriately
Use of model systems is central in this course
A model system is an organism or in vitro system chosen to study a process with as little other complexity as possible
Begin with a simple organism to isolate the process of interest away from confounding factors
Example: studying gene expression changes in response to environmental cues is easier in unicellular organisms (bacteria or yeast) than in a mouse with endocrine system, stress, and diet complications
Criteria when choosing a model system:
Simplicity of the organism and ease of isolating the process
Ease of laboratory handling and generation time
Ability to obtain genetically identical individuals to reduce background genetic variation
Generation times (illustrative):
Escherichia coli: generation time ≈
Mouse: generation time ≈
Mustard relative (Arabidopsis): generation time ≈
Commonly used unicellular organisms and systems:
Prokaryotes: bacteria (E. coli) and bacteriophages (viruses that infect bacteria)
Unicellular eukaryotes: yeast (Saccharomyces cerevisiae), Tetrahymena (unicellular eukaryote), Hydra (model for multicellular tissues in the instructor’s work)
Common multicellular model systems:
Mouse (Mus musculus)
Drosophila melanogaster (fruit fly)
Arabidopsis thaliana (mustard relative)
Why model systems? Nature is conservative: once an approach to performing a process evolves, it tends to be conserved across later generations
The same molecular machinery for DNA replication, etc., tends to be used in descendants
This conservation allows conclusions drawn in simple systems to be extrapolated to more complex organisms, with appropriate caveats
How model systems progress from simple to complex understanding:
Start in bacteria/phage (simplest system), then test whether the same mechanism holds in eukaryotes, then in multicellular eukaryotes
If differences arise, identify what factors differ (e.g., additional cellular compartments, tissue context)
Final takeaway: model systems enable experiments and data-driven conclusions about how processes work; human testing follows after animal and cellular work
Atoms, isotopes, and radioactive tracers
Fundamental unit of matter (in this course): atom with three main parts: protons, neutrons (in the nucleus), and electrons (orbiting the nucleus)
Key quantities:
Atomic number = number of protons; also equals the number of electrons in a neutral atom
Mass number where is the number of neutrons; determines atomic mass
Electron arrangement in orbitals and energy levels: electrons occupy shells with capacity limits
Shell capacity (simplified for course):
Innermost shell: up to electrons
Next shell: up to electrons
Third shell and beyond: many more; in practice, the course deals with atoms with relatively low (often for examples)
Isotopes: same element, same (and same chemical behavior) but different neutron number; some isotopes are radioactive
Radioactive tracers and isotopes commonly discussed:
Isotopes used as tracers: ${}^3H$, ${}^{14}C$, ${}^{32}P$, ${}^{125}I$, among others
Common light elements used: hydrogen, carbon, oxygen, sulfur, phosphorus; iodine-125 used for gamma emission tagging of proteins; phosphorus-32 and carbon-14 used for beta emitters in many experiments
Radioactive decay types (three main types discussed):
Alpha decay: emission of a helium nucleus (); not used in class experiments here
Gamma decay: emission of high-energy electromagnetic radiation (gamma rays); ${}^{125}I$ is a gamma emitter used in some contexts
Beta decay: emission of a beta particle (electron $e^-$) or positron ($e^+$); common in many biologically relevant tracers
In biology, tracers often rely on beta decay (electrons or positrons) or gamma emitters depending on detection method
Three detection methods for beta-decay tracers:
Geiger counter: quick, qualitative/semi-quantitative; good for contamination checks and rough ranking; crude localization; analog dial; not precise
Scintillation counting: highly quantitative; sample is placed in a scintillation floor that emits photons when decays occur; a photodetector counts photons to yield a precise activity value; best for exact quantitation but poor for precise localization
Autoradiography: localization technique; samples (e.g., cells or colonies) are placed on a film or emulsion; decays expose the film, creating autoradiograms; good for spatial localization but not highly quantitative
Autoradiograms and examples:
Bacterial colonies with radioactive tracer on filter paper pressed onto X-ray film show which colonies took up the tracer
DNA gels can be exposed to X-ray film to reveal which fragments carried the radioactive label
Practical note on safety and application:
The instructor used to be a radiation safety officer; modern practice in the campus reduces or eliminates the need for radioactive tracers in many experiments
Important takeaway on tracers:
Different detection methods serve different purposes: quick checks, precise quantitation, or precise localization
Atoms to molecules: bonds, polarity, and interactions in water
Ionic vs covalent vs hydrogen bonds
Ionic bonds: form when atoms transfer electrons, producing ions with opposite charges; electrostatic attraction between ions leads to bond formation
In solutions like water, ionic bonds tend to dissociate because ions form hydration shells with water molecules; not reliable for stable biomolecules in aqueous environments
Covalent bonds: form when atomic nuclei share electrons; the shared electron pair spends time around both nuclei
Covalent bonds are strong and require significant energy to form; breaking a covalent bond releases a lot of energy
Sharing is not equal in all cases; electronegativity differences matter (e.g., oxygen tends to hog electrons)
Hydrogen bonds: weaker than covalent and ionic bonds; arise from electrostatic attraction between partial charges (e.g., partial positive on H and partial negative on O in water)
These are dynamic and continually form/break with molecular motion
Polar vs nonpolar molecules and solubility in water
Polar (hydrophilic) molecules have poles or partial charges that interact favorably with water
Water is a polar solvent: the oxygen bears a partial negative charge; hydrogens carry partial positive charges; forms hydrogen bonded networks
Nonpolar (hydrophobic) molecules lack regional partial charges and tend to cluster in water, driven by the disruption of water’s hydrogen-bond network
Hydrophobic clustering example: nonpolar molecules coalesce in water not due to direct attraction between them, but due to water’s preference to maximize hydrogen bonding with itself
Water as solvent and its implications for biology
Almost all biological processes occur in aqueous environments, so polarity and hydrogen bonding govern molecular interactions
Key qualitative comparisons (without numerical details):
Covalent bonds: strong, high energy to form, large energy release upon breaking
Ionic bonds: medium strength in general, but easily broken in water due to hydration
Hydrogen bonds: weak individually, but numerous bonds yield substantial collective effects
Hydrophobic interactions: not a bond per se; aggregation reduces disruption of the water network
Quick memorable phrases:
“Oxygen hogs the electrons” describes why many bonds involving oxygen are polar
Polar = hydrophilic; Nonpolar = hydrophobic (water-averse)
Takeaway about polarity and structure:
Polar molecules (like water, many biomolecules) interact through hydrogen bonding and dipole interactions
Nonpolar molecules tend to aggregate in water, forming hydrophobic clusters
Chemistry 105/106: reaction kinetics and thermodynamics (summary of key concepts)
Core idea: cannot create or destroy energy; energy can be moved or transformed; track energy changes along a reaction path
Reaction progress and energy diagrams (conceptual):
Reactants start with a certain energy level; products finish at another energy level
Some reactions absorb energy to reach products (endergonic/endergonic) and have products at higher energy than reactants
Endergonic (endothermic): ext{Δ}G > 0; the products are at higher energy than reactants; non-spontaneous under the given conditions
Some reactions release energy (exergonic/exothermic) with products at lower energy than reactants
Exergonic: ext{Δ}G < 0; spontaneous under the given conditions
The net energy change is the Gibbs free energy change:
If ΔG < 0, the reaction tends to proceed; if ΔG > 0, the reaction tends to not proceed without energy input; ΔG = 0 corresponds to equilibrium (the transcript notes spontaneous behavior at ΔG = 0, which is scientifically inaccurate; at ΔG = 0 the system is at equilibrium and has no net tendency to proceed in either direction)
Path dependence and rate vs spontaneity
Knowing only start and end states does not reveal the path taken from reactants to products
Activation energy (Ea): energy barrier to reach the transition state; the difference between the energy of the transition state and the reactants
A reaction can be spontaneous (ΔG < 0) but slow if Ea is large; conversely, a non-spontaneous reaction can be accelerated by pushing conditions or catalysis
Catalysts and their role
Catalyst lowers the activation energy by providing an alternative pathway to the transition state
Catalyst is not consumed; it does not change the overall ΔG or the product composition
By lowering Ea, a catalyst increases the rate of the reaction under specific conditions; the final equilibrium composition remains the same
Classic example: sodium (in water) reaction
An exergonic reaction that can be very fast due to a very low Ea in that environment (highly favorable to proceed quickly)
Important nuance about ΔG and spontaneity (clarifications):
Correct thermodynamic statement: if ext{Δ}G < 0, the reaction tends to be spontaneous under the given conditions; if ext{Δ}G > 0, the reaction is non-spontaneous under those conditions; if , the system is at equilibrium
The lecture notes mention a common point of confusion: some claims say “ΔG = 0 means spontaneous,” which is not correct in standard thermodynamics; notes here reflect the instructor’s wording but include the correction for accuracy
Practical note for future exam questions
Expect questions about how catalysts affect reaction rates vs ΔG, how to interpret ΔG diagrams, and the conceptual difference between rate vs spontaneity
Brief segue to acid-base topics
The plan for the next class: cover acid-base kinetics and rederive the Henderson–Hasselbalch equation; note: class schedule includes a Wednesday session due to a Monday cancellation
Connections, themes, and practical implications
Modeling and experiments rely on simplifying assumptions (reductionism) but must acknowledge emergent properties and limits of simple models
Model systems enable controlled experiments to reveal mechanisms that often transfer to more complex organisms, with careful validation
Understanding atomic-level interactions (bonds, polarity, solvation) is essential to predict macroscopic properties like solubility, reaction rates, and protein folding
Radiotracers and detection methods illustrate different levels of information gathering: qualitative presence/absence, quantitative activity, and spatial localization
Thermodynamics and kinetics together determine whether a reaction occurs and how fast it proceeds under physiological conditions; catalysts are central to metabolism because they tune reaction rates without altering the final outcomes
Reminders for the course roadmap
Expect future discussions on acid-base kinetics and Henderson–Hasselbalch equation derivation
The material covered here ties into concepts from chem 105 and 106 and will underpin upcoming laboratory and problem-solving contexts