1. Electronic Structure of Atoms
Chapter 6

2. 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

3. 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

4. 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

6. Bohr Model

7. 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

8. 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

9. Schrödinger’s Plot of Hydrogen

10. 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

11. Sublevel shapes

12. Sublevel shapes

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

15. 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

16. 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

17. Types of Electron Configurations
Quantum number
Orbital diagram
Shorthand
Noble Gas