Biology: Atoms, Electrons, and Bonding

Atoms, Isotopes, Electrons, and Bonding

  • Isotopes and nuclear composition

    • Isotopes differ in neutron number but have the same number of protons. The speaker notes: "This is an isotope" and discusses neutrons and protons in the nucleus.
    • An atom’s identity is defined by its protons (the atomic number). The nucleus contains positively charged protons and neutral neutrons.

  • Normal atoms and ions

    • A neutral atom has the same number of protons and electrons; the electrons balance the positive charge of the nucleus.
    • Ions are formed when the number of electrons changes (not just the nucleus). If electrons are lost, the atom becomes a positively charged cation; if electrons are gained, it becomes a negatively charged anion.

  • Example: Carbon (as a reference)

    • Carbon-12 (C-12): atomic mass ~ mass number A = 12, atomic number Z = 6 (protons), neutrons N = A − Z = 6, electrons = 6 when neutral.

  • Electrons and the nucleus: scale and uncertainty

    • Electrons are very small and carry negative charge; protons are positively charged.
    • We cannot know the exact position of an electron; we speak of probability spaces around the nucleus where electrons are likely to be found.
    • These probability spaces are called energy shells (orbitals) that surround the nucleus and are outside the nucleus.

  • Energy shells and shell capacities

    • First shell (K shell) is closest to the nucleus and holds 22 electrons.
    • Second shell (L shell) holds 88 electrons.
    • Third shell (M shell) also holds 88 electrons for the biologically relevant atoms discussed.
    • We fill shells from the inside out (innermost first).
    • The spoken shorthand uses K, L, M for the first, second, and third shells.

  • Energy concepts: kinetic vs potential energy

    • Kinetic energy: energy of motion.
    • Potential energy: stored energy that can be converted to kinetic energy.
    • Example parallels:
    • A pen held above the floor has potential energy; when dropped, gravity converts it to kinetic energy.
    • A ball rolling down steps stops when it hits a step due to discrete energy states (steps) that resemble discrete electron shells.
    • Water in rapids also shows energy transformation with velocity changes along the course.
    • In atoms, electrons can move outward to higher shells by gaining energy; they can also fall back to lower shells, releasing energy.

  • Fluorescence as an energy-shelled process

    • Fluorochrome fluorescein absorbs energy from a photon, promoting outer-shell electrons to a higher energy state.
    • When electrons collapse back toward the nucleus, they emit a photon with lower energy than the absorbed one.
    • In fluorophores, some energy is lost as heat to the surroundings during this process.
    • Example energy relation: a higher-energy photon of frequency (
      u{initial} ) is absorbed, and a lower-energy photon of frequency ( u{emitted} ) is emitted, with the difference going into heat: hν<em>initial=hν</em>emitted+Q  (Q≥0).h\nu<em>{initial} = h\nu</em>{emitted} + Q\; (Q\ge 0).
    • Fluorescence typically shifts from short wavelength (high energy) to longer wavelength (lower energy) emission.

  • Stability and filled electron shells

    • Atoms are generally more stable when their electron shells are filled.
    • Hydrogen tends to seek two electrons to fill its first shell (2 electrons total) and is therefore more reactive.
    • Helium has a filled first shell (2 electrons) and is very stable/rarely reactive.
    • Other elements shown: electrons per shell for several atoms illustrate how shell filling trends influence reactivity (e.g., Na with 2 in the first shell, 8 in the second, 1 in the third; Ne with 2 in the first shell and 8 in the second shell, a filled second shell).
    • In the context of biology, oxygen is highly reactive because its second shell is short by two electrons to full; sodium has one extra electron in the outer shell, making it reactive as well.

  • Covalent bonds: sharing electrons

    • Covalent bond: a pair of valence electrons shared between two atoms.
    • Examples of covalent bonds:
    • ( \mathrm{H_2} ): two hydrogens share one electron each to satisfy a shell.
    • ( \mathrm{O_2} ): two oxygens share electrons.
    • ( \mathrm{CH_4} ): carbon shares electrons with four hydrogens.
    • Water: ( \mathrm{H_2O} ) involves sharing between one oxygen and two hydrogens.
    • Electron sharing details:
    • Hydrogen: each H has 1 electron; to fill its first shell (2 electrons) it shares with another H so both see a full shell: a covalent bond forms when electrons are at an optimal distance (bond length).
    • Oxygen: outer shell needs 2 more electrons to reach a filled second shell (8 electrons total in second shell). In ( \mathrm{H_2O} ), O shares with two hydrogens so that O effectively has 8 electrons in its outer shell.
    • Polar vs nonpolar covalent bonds
    • Nonpolar covalent bonds: electrons shared equally (e.g., in ( \mathrm{H2} ), ( \mathrm{O2} ), and ( \mathrm{CH_4} )).
    • Polar covalent bonds: electrons shared unequally due to differing electronegativities; strongest example is water, where oxygen is more electronegative than hydrogen.

  • Electronegativity and partial charges

    • In polar bonds, electrons spend more time closer to the more electronegative atom, creating partial charges rather than full charges.
    • Oxygen in ( \mathrm{H_2O} ) holds a partial negative charge: δ−\delta^-; hydrogens hold partial positive charges: δ+\delta^+.
    • These partial charges enable dipole-dipole interactions and hydrogen bonding between molecules.
    • Hydrogen bonding (weak, but biologically important) arises from attractions between partial charges in polar molecules (e.g., between water molecules or within biopolymers containing polar groups).

  • How polar and nonpolar bonds relate to biology

    • Polar covalent bonds abound in carbohydrates and other molecules rich in oxygen and nitrogen; these molecules are typically water-attracting (hydrophilic).
    • Nonpolar covalent bonds are common in lipids and cell membranes, which are largely hydrophobic (lipid-rich, water-insoluble environments).
    • The polarity of molecules influences interactions inside cells and in the extracellular environment.

  • Ionic bonds and electrolytes

    • Ionic bonds form when atoms exchange electrons to achieve full outer shells, producing ions with full charges.
    • Example: Sodium (Na) loses an electron to become Na⁺; Chlorine (Cl) gains an electron to become Cl⁻. They attract electrostatically to form NaCl (table salt).
    • Salt crystals in the absence of water are held together by electrostatic attractions between opposite charges.
    • In water, hydration shells around ions reduce the strength of the ionic attraction and help dissolve salts; salts can crystallize again if water is removed (evaporation).
    • Ions like Na⁺, K⁺, and others are biologically important; ions play crucial roles in signaling, membrane potential, and enzyme activity.

  • Bonds in context: strengths and roles

    • Covalent bonds are among the strongest bonds in biology, requiring substantial energy to break.
    • Weak interactions include hydrogen bonds and Van der Waals forces; they accumulate across many interactions to shape structures and dynamics (e.g., protein folding, DNA base pairing, lipid bilayer cohesion).
    • Ionic bonds are relatively weaker than covalent bonds but can be strong in low-dielectric environments (like dry crystals) and are modulated by solvent (e.g., water).

  • Quick recap of terms and concepts

    • Nucleus: protons (+) and neutrons (neutral).
    • Electrons: negatively charged; occupy probability spaces around the nucleus in discrete energy shells.
    • Shell capacities (for the shells discussed): 2,8,82, 8, 8 electrons respectively.
    • Covalent bonds: share electrons; can be nonpolar (equal sharing) or polar (unequal sharing).
    • Ionic bonds: electrons are exchanged; atoms become ions that attract electrostatically.
    • Hydrogen bonds: weak attractions between partial charges in polar molecules; essential for water structure and biomolecules.
    • Van der Waals forces: very weak interactions that contribute to molecular packing and interactions.

  • Foundational connections

    • Fillings of electron shells explain chemical reactivity patterns across elements (e.g., why Na and O are highly reactive, why noble gases like Ne are inert).
    • The behaviors of water (high polarity, hydrogen bonding, density anomaly upon freezing) underlie many biological processes and the structure of life.
    • The distinction between polar and nonpolar regions of molecules helps explain membrane structure, protein folding, carbohydrate interactions, and nucleic acid chemistry.

  • Practical implications highlighted in the talk

    • Water’s polarity and hydrogen bonding explain its solvent properties and the behavior of biological macromolecules in aqueous environments.
    • Understanding covalent vs ionic bonding helps explain the formation and stability of biomolecules and minerals in organisms and their environments.
    • The interplay of strong covalent bonds with weak hydrogen bonds and ionic interactions shapes molecular recognition, assembly of macromolecular complexes, and cellular structures.

  • Note on notation used in the lecture

    • Shells named as K (1st), L (2nd), M (3rd).
    • Atomic number Z (protons) and mass number A; neutron count N = A − Z.
    • Partial charges in polar covalent bonds denoted as δ−\delta^- (partial negative) and δ+\delta^+ (partial positive).
    • Chemical formulas and molecules expressed in standard form, e.g., ( \mathrm{H2} ), ( \mathrm{O2} ), ( \mathrm{H2O} ), and ( \mathrm{CH4} ).

  • Suggested connections for exam prep

    • Be able to explain the difference between isotopes and ions with simple examples.
    • Memorize the first three shell capacities: 2,8,82, 8, 8 and describe how shells fill from inside out.
    • Describe how fluorescence demonstrates energy absorption and emission, including energy conservation and heat dissipation.
    • Distinguish polar vs nonpolar covalent bonds and give molecular examples (e.g., water vs methane).
    • Explain how ionic bonds form and how hydration in water affects ionic interactions.
    • Understand why covalent bonds are generally stronger than hydrogen bonds or ionic attractions in dry environments, but how weak interactions collectively contribute to macromolecular structures.