unit 3 review research chem

Overview

• Unit 3 covers Periodic Table organization, electron configurations, orbital diagrams, PES, and periodic trends (Coulombic attraction, atomic radius, ionization energy, electronegativity, shielding, ionic radius).

• Also includes light (wave–particle duality, EMR, Planck’s constant, Bohr model, emission spectra).

Organization of the Periodic Table

• Mendeleev (1869): Organized by atomic mass, left gaps.

• Moseley (1914): Reorganized by atomic number (protons) - modern PT.

• Current Goals: Elements with similar properties in same columns (groups); rows (periods) correspond to principal energy levels.

Terminology: Vertical Columns (Groups) and Horizontal Rows (Periods)

• Groups (vertical): Same number of valence electrons, similar properties.

  • Group 1: Alkali Metals

  • Group 2: Alkaline Earth Metals

  • Groups 3–12: Transition Metals

  • Group 17: Halogens

  • Group 18: Noble Gases

• Periods (horizontal): Each corresponds to a new principal energy level being filled (n = 1, 2, 3, …).

Valence Electrons

• Rule (s- and p-block): Last digit of group number = number of valence electrons.

  • Example: Group 1 has 1 val e-, Group 14 has 4 val e-.

• Exception: Helium (He) has 2 valence electrons.

Electron Configurations

• Basic Idea: Electrons fill energy levels/sublevels in specific order for stability.

  • Sublevel capacities: s: 2 e-, p: 6 e-, d: 10 e-, f: 14 e-.

  • Each orbital holds 2 electrons (opposite spins).

• Filling Rules:

  • Aufbau Principle: Fill lowest energy levels/subshells first (e.g., 1s before 2s).

  • Hund’s Rule: Occupy equal-energy orbitals singly before pairing, with parallel spins.

  • Pauli Exclusion Principle: No two electrons in same orbital have same quantum numbers (max one spin-up, one spin-down).

• Order of Filling (typical): $1s \to 2s \to 2p \to 3s \to 3p \to 4s \to 3d \to 4p \dots$

• Noble Gas Shorthand: Use closest prior noble gas in brackets, then continue config.

  • Example: Cl ($[\text{Ne}] 3s^2 3p^5$).

• Isoelectronic: Ions/atoms with the same number of electrons (e.g., F$^-$, Na$^+$, Ne).

• Ions: Configuration reflects new electron count; often matches a noble gas.

Orbital Diagrams

• Visualize all four quantum numbers: energy level (n), sublevel (l), orbital ($m<em>l$), spin ($m</em>s$).

• Follow Aufbau, Hund’s, Pauli rules for filling.

Photoelectron Spectroscopy (PES)

• What it is: Experimental technique measuring electron binding energies.

• Binding Energy: How tightly an electron is held.

• Key Concept: Electrons closer to nucleus (lower energy) have higher binding energies.

• PES diagrams show peaks corresponding to core/valence electrons; height indicates number of electrons.

Periodic Trends

• Core Idea: Coulombic Attraction: Attraction between protons (+) and electrons (-).

  • Coulomb’s Law (conceptual): $F = k \frac{q<em>1 q</em>2}{r^2}$

  • $q1, q2$: charges; $r$: distance between nucleus and electrons.

Trends Summary

1. Atomic Radius (size of atom)

   • Group: Increases down (more energy levels, e- farther).

   • Period: Decreases across (increasing nuclear charge pulls e- closer).

2. Ionization Energy (IE) (energy to remove an electron)

   • Group: Decreases down (valence e- farther, more shielded).

   • Period: Increases across (nuclear charge increases, pulling e- closer, harder to remove).

3. Electronegativity (EN) (ability to attract e- in a bond)

   • Group: Decreases down (increased shielding/distance reduces attraction).

   • Period: Increases across (higher effective nuclear charge, stronger attraction for bonding e-).

4. Shielding Effect (inner e- shield valence e- from nuclear charge)

   • Group: Increases down (more inner shells).

   • Period: Roughly constant (same inner shells, but increasing $Z_{\text{eff}}$).

5. Ionic Radius (size of ion)

   • Cations ($G^+$): Smaller than neutral atom (loss of e- reduces repulsion).

   • Anions ($G^-$): Larger than neutral atom (gain of e- increases repulsion).

Light: Wave–Particle Duality and Electromagnetic Radiation

• Light acts as both a wave (wavelength, frequency) and particle (photons).

• Wave Description: Electromagnetic spectrum (radio to gamma rays).

• Particle Description: Photon energy related to frequency and wavelength.

  • Planck’s constant (h): $E = h \nu$

  • Speed of light (c): $c = \lambda \nu$

  • Combined: $E = \frac{hc}{\lambda}$

  • $c = 3.00 \times 10^8 \text{ m s}^{-1}$; $h = 6.626 \times 10^{-34} \text{ J s}$.

• Bohr Model & Emission Spectra: Electrons absorb energy, jump to excited states; emit photons (light) when relaxing back down. Each element has a unique emission spectrum.

Key Formulas and Concepts

• Coulombic attraction: $F = k \frac{q<em>1 q</em>2}{r^2}$

• Speed of light: $c = \lambda \nu$; $c = 3.00 \times 10^8 \text{ m s}^{-1}$

• Photon energy: $E = h \nu$; $E = \frac{hc}{\lambda}$; $h = 6.626 \times 10^{-34} \text{ J s}$

• Sublevel capacities: $s:2, p:6, d:10, f:14$

How to Use These Notes Quickly

• Focus on the main ideas for each section.

• Understand the direction and reason for each periodic trend.

• Memorize the electron configuration filling order and rules.

• Connect light concepts (E, $\nu$, $\lambda$) to the Bohr model and spectra.

• Practice core examples for configurations and ions.