Science PDF - Atomic Structure and Periodicity

1. Major Historical Figures in Atomic Theory Democritus: Early idea of indivisible particles ("atomos"). John Dalton: Proposed the atomic theory stating that atoms of each element are identical and atoms combine to form compounds. J.J. Thomson: Discovered the electron using cat...

1. Major Historical Figures in Atomic Theory Democritus: Early idea of indivisible particles ("atomos"). John Dalton: Proposed the atomic theory stating that atoms of each element are identical and atoms combine to form compounds. J.J. Thomson: Discovered the electron using cathode rays, leading to the "plum pudding" model. Ernest Rutherford: Discovered the nucleus through the gold foil experiment, proposed a dense, positively charged center. Niels Bohr: Developed the planetary model with electrons in fixed orbits around the nucleus. Erwin Schrödinger: Formulated the quantum mechanical model using wave functions. Werner Heisenberg: Established the uncertainty principle, indicating that one cannot know both the position and momentum of an electron simultaneously. 2. Mass Spectrometer and Mass Spectrum Mass Spectrometer: Device that separates ions based on mass-to-charge ratio. Main components include: ○ Ionization Chamber: Converts atoms/molecules into ions. ○ Acceleration: Ions are accelerated into a magnetic field. ○ Deflection: Ions are separated based on mass and charge. ○ Detection: Ions are detected, and their abundances are recorded. Mass Spectrum: A chart showing the relative abundances of ions based on their mass-to-charge ratio. Used to identify isotopes and calculate the relative atomic mass of an element. 3. Relative Atomic Mass, Isotopes, and Average Atomic Mass Isotope: Atoms of the same element with different numbers of neutrons. Relative Atomic Mass (Ar): Weighted average mass of all isotopes of an element. Average Atomic Mass Calculation: (mass of isotope 1×abundance)+(mass of isotope 2×abundance)(\text{mass of isotope 1} \times \text{abundance}) + (\text{mass of isotope 2} \times \text{abundance})(mass of isotope 1×abundance)+(mass of isotope 2×abundance), where abundance is a decimal. Example: For chlorine isotopes 35Cl^35Cl35Cl (75%) and 37Cl^37Cl37Cl (25%), average atomic mass = 35×0.75+37×0.25=35.535 \times 0.75 + 37 \times 0.25 = 35.535×0.75+37×0.25=35.5. 4. Radioisotopes and Types of Radioactive Decay Radioisotopes: Unstable isotopes that decay, emitting radiation. Types of Decay: ○ Alpha Decay: Emission of an alpha particle (24He^4_2He24​He). Reduces mass number by 4 and atomic number by 2. ○ Beta Decay: Emission of a beta particle (electron) as a neutron converts to a proton. Increases atomic number by 1. ○ Gamma Decay: Emission of gamma rays (high-energy photons), usually following alpha or beta decay, with no change in atomic or mass numbers. Balancing Decay Equations: Ensure mass and atomic numbers are conserved. 5. Quantum Mechanical (Probability) Model and Quantum Numbers Quantum Mechanical Model: Describes electrons as wave-like and particle-like, with probabilities for finding electrons in specific regions (orbitals). Quantum Numbers: ○ n (Principal Quantum Number): Indicates energy level (e.g., n = 1, 2, 3). ○ l (Azimuthal Quantum Number): Defines subshell (s, p, d, f) and shape of the orbital (e.g., l = 0 for s, l = 1 for p). ○ ml_ll​(Magnetic Quantum Number): Describes the orientation of the orbital. ○ ms_ss​(Spin Quantum Number): Electron spin, either +1/2 or -1/2. 6. Wave Mechanical Model, Heisenberg Uncertainty Principle, and Orbitals Heisenberg Uncertainty Principle: States that it is impossible to know both the exact position and momentum of an electron at the same time. Orbitals: ○ s-Orbitals: Spherical shape; each energy level has one s-orbital. ○ p-Orbitals: Dumbbell-shaped; starts at n=2 with three orientations (px_xx​, py_yy​, pz_zz​). ○ d-Orbitals: Complex shapes; five orientations starting at n=3. ○ f-Orbitals: Even more complex shapes; seven orientations starting at n=4. 7. Electron Configuration and Principles Electron Configuration: Shows the arrangement of electrons in orbitals (e.g., Carbon: 1s22s22p21s^2 2s^2 2p^21s22s22p2).  Abbreviated Configuration: Uses noble gases to shorten notation (e.g., Iron: [Ar] 3d64s23d^6 4s^23d64s2). Orbital Diagrams (Box Diagrams): Illustrate electron filling in orbitals with boxes and arrows for spins. Principles: ○ Aufbau Principle: Electrons fill the lowest energy orbitals first. ○ Hund’s Rule: Electrons fill orbitals singly before pairing within a subshell. ○ Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers (each orbital holds a max of 2 electrons with opposite spins). Octet Rule: Atoms tend to fill their outermost s and p orbitals with 8 electrons. 8. Periodic Table Structure Groups: Vertical columns. Elements in the same group have similar chemical properties. Periods: Horizontal rows. Properties vary across a period due to changing effective nuclear charge. Sections: ○ Metals: Found on the left side; typically shiny, conductive, and malleable. ○ Non-Metals: Found on the right side; usually poor conductors and brittle. ○ Metalloids: Along the "stair-step" line; have mixed properties. 9. Effective Nuclear Charge (Zeff_{eff}eff​) and Shielding Zeff_{eff}eff​: The effective positive charge felt by outer electrons, reduced by inner electron shielding. Shielding: Inner electrons reduce the attraction between the nucleus and outer electrons. 10. Periodic Trends Atomic Radius: Decreases across a period (due to increasing Zeff_{eff}eff​) and increases down a group (due to additional electron shells). Ionic Radius: Cations are smaller than their parent atoms; anions are larger. Ionization Energy: Energy required to remove an electron; increases across a period and decreases down a group. Electron Affinity: Tendency of an atom to gain an electron; generally increases across a period. Electronegativity: Measure of atom’s ability to attract electrons in a bond; increases across a period, decreases down a group. 11. Reactivity of Metals and Non-metals Metals: Reactivity increases down a group (e.g., alkali metals) due to lower ionization energy. Non-Metals: Reactivity generally increases up a group (e.g., halogens) as atoms more readily attract electrons. This guide summarizes key concepts. Focus on understanding trends and practicing electron configurations, orbital diagrams, and isotope calculations, as these are foundational for mastering atomic theory and periodic trends.