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Chapter 6 : Electronic Structure of Atoms
gamma, X ray, IntraR, ROYGBV, UltraV, Mircowaves, Radio
Quantised energy (absorbed in discrete chunks)
Planks constant: h (E=hv)
Emission of light from hot objects
Emission of electrons from metals on which light shines
Explained photoelectric effect
Named energy packets "photons"
Electrons are not capable of absorbing energy
Energy required to release an electron from a metal
Energized gasses produce light of distinct wavelengths
Electrons falling from energized states
Explained emission spectra with energy states of hydrogen
Assumed that electrons did not spiral into nucleus w/o explanation
No element except hydrogen explained, flaws in planetary model.
Wrote eqn for hydrogen emission spectra that evolved into Rydberg equation
Trial and error
3 things that describe light
Gold foil experiment
Dense nucleus, electron cloud
Louis de Broglie
Returns to Bohr postulate
Proof: Electrond diffract through crystal
Equations to find probabilities of electron location for each orbital: Quantum/wave mechanics
Probability density: square of wave function
Uncertainty principle: Limit to knowledge of electrons momentum and location
Things explained by Quantum model for electrons
Colors of Transition metal aqeus ions
Bonding patterns of ions
Paramagnetic behavior of some atoms
Outside Noble gas core
Amount of energy to remove one electron (valence)
Orbitals that have the same energy
Electrons always fill the lowest energy orbitals first (violated with some transition metals - Cr)
For degenerate orbitals, one electron each, then double up
Pauli exclusion principle
No two electrons can have the same four quantum numbers
Positive integers, Defines energy of electron
Principal quantum number
0-(n-1), Defines shape of orbital
Angular momentum (azimuthal) quantum number
-l...0...+l, defines orientation of orbital in space
distingishes between orbitals in the same subshell
Magnetic quantum number
-1/2 or +1/2, spin up of spin down
Orbitals with the same value of n
Orbitals with the same n and l values
probability function goes to zero for an orbital at this location (no nodes for s orbitals)
electron pairs explain what?
lines in emission spectra actually closely spaced pairs, split by opposite magnetic fields.
Weird orbital filling trend with certain transition metals
Cr and Mo, Cu and Au and Ag, others
Will fill both s and d orbitals singly rather than completely filling s orbital. ex: Cr= [Ar] 4s^1 3s^5
For representative elements, d and f orbital electrons do not count as valence electrons.
For transition metals, f orbital electrons don't count.