Electronic Structure of Atoms Chapter 6

Wave Nature of Light Visible light is part of the electromagnetic spectrum Speed of ALL light = 3.00 x 108 m/s Speed of light (c) is a constant that connects wavelength (λ) and frequency (v) c = λ x v

Quantized Energy and Photons Relationship between temperature, intensity, and frequency of light emitted Planck assumed that energy could be released in small pieces or chunks (quantum) E = h x v High frequency wavelengths have high energy and can cause damage to tissue

Bohr’s Model of the Hydrogen Atom Continuous spectra- created by most sources of radiation Discontinuous spectra – created by sources that only emit a few wavelengths of light Bohr’s Model was used to explain discontinuous spectra Orbits correspond to specific energies When electrons move from higher to lower orbits, a specific amount of energy (light) is emitted which corresponds to specific wavelengths of light

Bohr Model

Wave Behavior of Matter de Broglie proposed that an electron in its movement about the nucleus has a wavelength that is dependent upon its mass and velocity Heisenberg Uncertainty Principle Since electrons can behave as both wave and particle- it is not possible to know the position and momentum at the same time

Quantum Mechanics and Atomic Orbitals Schrodinger developed mathematical equations to incorporate both particle and wave nature of electron Solution leads to wave function psi (ψ) ψ2 gives information about the location of an electron (probability density) Orbitals Each orbital has characteristic energy and shape Quantum Numbers Principle quantum number (n) – level – 1, 2, 3… Second quantum number (l) – shape - 0 to n-1 Third quantum number (ml) – orientation – l to -l

Schrödinger’s Plot of Hydrogen

Representation of Orbitals S orbitals – spherical P orbitals – “dumb-bell” with 2 lobes 3 p orbitals – px, py, pz D orbitals – “cloverleaf” 5 d orbitals – dyz, dxz, dxy, d x2-y2, d z2 F orbitals 7 f orbitals

Sublevel shapes

Sublevel shapes

Sublevels combined… Energy Level 1 Energy level 2 Energy level 3

Orbitals in Many-Electron Atoms Effective nuclear charge Each electron is attracted by the nucleus and repelled by other electrons Effective nuclear charge is the net positive charge attracting an electron which equals the number of protons in the nucleus Electrons farther from the nucleus are less affected by the nuclear charge (screening or shielding effect) Energy of orbitals increases with distance from nucleus Within a given level s < p < d < f

Electron Configurations Description of the arrangement of electrons among the orbitals Stable electron configurations have the lowest possible energy states Rules Aufbau Hund Pauli exclusion principle Use the periodic table to assist with determining the electron configuration

Types of Electron Configurations Quantum number Orbital diagram Shorthand Noble Gas