Biology Week 1 part 3
Ant Experiment and Scientific Practice
Experimental design described: ants were divided into three groups to test how leg length affects their movement and distance to a feeder.
Group 1 (green): legs left untouched.
Group 2 (yellow): legs cut short.
Group 3 (right): legs lengthened by adding material.
A total of 75 ants were used.
They observed the ants leaving the nest and traveling to the feeder, then returning to the nest.
Prediction: leg length would influence travel distance/time; shorter legs would lead to different travel outcomes compared to longer legs.
Result discussed: the distance traveled from the nest to the feeding mass was measured, and the distance reported remained at a fixed value of 10 meters for the feeding mass distance.
They repeated measurements to reduce distortions due to small sample size and to check significance.
Summary of the scientific process highlighted:
Biologists practice evidence-based decision making: ask questions about how organisms work, pose hypotheses, and use experimental or observational evidence to decide whether hypotheses are correct.
There are two broad kinds of science:
Applied science
Basic (pure) science: aims to expand knowledge regardless of short-term applications (e.g., studying dolphin physiology to understand how they hold their breath underwater).
Dissemination of results is essential: scientists publish findings in scientific journals to let the community learn, build upon work, and continue progress. A scientific paper usually answers: What question was asked? How was it tested (experiments)? What were the results? What conclusions were drawn?
The publication process includes describing the methods and results in detail so others can replicate and evaluate.
Basic science vs. Applied science contrasted:
Basic science seeks to expand fundamental understanding (e.g., physiology, energy use in organisms).
Applied science uses knowledge to address practical problems or applications.
Scientific communication and journals:
Findings are disseminated to the scientific community to advance knowledge and enable subsequent research.
Quick hormonal primer for building blocks of life (transition to chemistry):
There are ions and molecules that serve as building blocks in chemical evolution.
About 96% of the matter in living organisms is composed of four elements: ext{H}, ext{C}, ext{N}, ext{O}.
Matter, Atoms, and Elements: Quick Primer
Matter is anything that takes up space and has mass.
Mass vs. weight:
Mass is the quantity of matter in an object and is constant regardless of location.
Weight depends on gravity: W = m g where W is weight, m is mass, and g is the acceleration due to gravity.
Example: an object has the same mass on Earth and the Moon, but its weight differs because gravity differs.
Element defined:
An element is a substance that cannot be broken down into simpler substances by chemical means.
Atoms of an element are characterized by their atomic number (number of protons).
Atomic structure basics (as introduced):
Nucleus contains protons (positive charge) and neutrons (neutral).
Electrons are negatively charged and reside in electron shells around the nucleus.
The total mass of an atom is largely due to protons and neutrons (collectively, nucleons).
Hydrogen is noted as an exception in some simplifications because its most common isotope has 1 proton and 1 electron and no neutrons.
Notation concepts:
Elements are represented by chemical symbols (e.g., H, C, N, O).
Atomic number Z equals the number of protons.
Mass number A equals the total number of protons and neutrons: A = Z + N where N is the number of neutrons.
In isotopes, A can vary while Z stays the same.
The mass number is often written as a superscript (e.g., ^{A}_{Z} ext{X} for an element symbol X).
Electron configuration (illustrative for key elements):
In carbon (atomic number 6): the distribution is 2 electrons in the first shell, and 4 electrons in the second shell (2, 4).
Oxygen (atomic number 8) has electrons arranged to fulfill the octet principle through sharing and bonding.
Energy basics (context for chemical bonding):
Energy is the capacity to do work or to transfer heat.
Two main types of energy:
Kinetic energy: E_k = frac{1}{2} m v^2
Potential energy: stored energy due to position or configuration.
Chemical energy is a form of potential energy stored in chemical bonds.
Conceptual note: the transcript emphasizes that the more stable a position in space or the more complex a structure, the more potential energy is present. (This reflects the context of chemical energy in bonds—though in standard chemistry, bond formation typically lowers potential energy as systems become more stable; the notes reflect the transcript's wording.)
Bonds, Electronegativity, and Bond Types
Covalent bonds: formed when atoms share electrons to fill outer shells and achieve stability.
Example molecules formed by covalent bonds include:
Water: ext{H}_2 ext{O}
Ammonia: ext{NH}_3
Methane: ext{CH}_4
Oxygen gas: ext{O}_2
In these cases, the outer electron shells become more stable as electrons are shared.
In covalent bonds, electrons are shared between atoms. If the sharing is equal, the bond is nonpolar covalent; if sharing is unequal due to differing electronegativities, the bond is polar covalent.
Electronegativity:
Electronegativity is the tendency of a nucleus to attract shared electrons toward itself.
It varies by element and is related to the number of protons and the distance of the outer shell from the nucleus.
In water (H–O), oxygen is more electronegative than hydrogen, pulling shared electrons closer to oxygen and creating a polar covalent bond with partial charges.
Polar vs nonpolar covalent bonds:
Polar covalent bonds: unequal sharing of electrons, creating partial charges on atoms (e.g., in H2O, NH3, H–O bonds).
Nonpolar covalent bonds: equal sharing of electrons, little to no partial charge separation (e.g., many C–H bonds in hydrocarbons).
The degree of polarity is tied to electronegativity differences; larger differences yield more polar bonds.
Ionic bonds (contrast with covalent):
Occur when electrons are transferred from one atom to another, creating ions (cations and anions).
Example: sodium chloride (table salt) is formed by ionic bonds between Na^+ and Cl^-.
Ionic bonds are generally weaker in aqueous solutions (in water, ions dissociate and interact with solvent).
The formation process can be summarized: a sodium atom loses an electron to form a cation (Na^+), while chlorine gains an electron to form an anion (Cl^-).
When dissolved in water, NaCl dissociates into Na^+ and Cl^- ions.
Additional terminology and examples:
Molecules with covalent bonds (e.g., ext{CO}_2) can be compounds if they consist of more than one element; CO2 is a molecule with covalent bonds, not an ionic compound.
A molecule like ext{O}_2 is a diatomic molecule held together by covalent bonds.
In the example of methane ext{CH}_4, carbon forms covalent bonds with four hydrogens; these are predominantly nonpolar, contributing to the molecule's stability and significant chemical energy in nonpolar bonds.
“Building blocks” and relevance to biology:
Covalent and ionic bonds determine how atoms connect to form macromolecules and metabolic pathways.
The polarity of bonds influences solubility, reactivity, and interactions in biological systems (e.g., water's polarity drives hydrogen bonding and solvent properties).
Molecules, Compounds, and Their Distinctions
Molecules: two or more atoms held together by covalent bonds. Examples: ext{CO}2, ext{H}2 ext{O}, ext{NH}3, ext{CH}4, ext{O}_2.
Compounds: substances composed of two or more elements chemically bonded together; all salts like ext{NaCl} are compounds (involving ionic bonds between different elements).
Key contrasts:
Covalent bonds involve sharing electrons; molecules can be covalently bonded entities (e.g., CO2, H2O).
Ionic bonds involve transfer of electrons, creating charged ions that attract each other (e.g., NaCl).
Why Bonds Matter: Energy, Stability, and Biological Relevance
Bonding affects energy storage and release: chemical reactions involve making/breaking bonds, changing potential energy.
The role of polarity and bond type influences:
Molecular shape and dynamics
Solubility in water and interactions with other molecules
Reactions and energy changes in metabolism
Quick References and Formulas
Mass–weight relationship: W = m g
Kinetic energy: E_k = frac{1}{2} m v^2
Potential (chemical) energy: considered as energy stored in bonds; specific expressions depend on the system.
Atomic notation: mass number and atomic number relationships: A = Z + N
Common building-block elements of life: ext{H}, ext{C}, ext{N}, ext{O}
Notation for a diatomic molecule example: ext{O}_2
Examples of molecules and ions mentioned: ext{H}2 ext{O}, ext{NH}3, ext{CH}4, ext{CO}2, ext{NaCl}
Practical Takeaways for Exam Preparation
Understand the experimental design and what constitutes a controlled comparison (three groups, manipulated variable, constant distance, replication for significance).
Differentiate between covalent and ionic bonding and recognize examples of each.
Recognize the concepts of electronegativity, polar vs nonpolar bonds, and how these influence molecular properties.
Distinguish between molecules and compounds and know examples of each from biology (e.g., ext{H}2 ext{O}, ext{CO}2, ext{NaCl}).
Recall the basic definitions of matter, mass, weight, elements, atoms, isotopes, and the nucleus composition.
Be able to connect chemical bonding concepts to biological contexts such as energy storage in chemical bonds and protein/nucleic acid interactions.
Practice applying the notation and simple formulas to describe energy, mass, and bonding: for example, using W = m g and E_k = frac{1}{2} m v^2 when appropriate.