30 terms

# Chapter 6 : Electronic Structure of Atoms

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Electromagnetic spectrum
gamma, X ray, IntraR, ROYGBV, UltraV, Mircowaves, Radio
Max Plank
Quantised energy (absorbed in discrete chunks)
Planks constant: h (E=hv)
Emission of light from hot objects
Photoelectric effect
Emission of electrons from metals on which light shines
Albert Einstein
Explained photoelectric effect
Named energy packets "photons"
Electrons are not capable of absorbing energy
Work function
Energy required to release an electron from a metal
Line Spectra
Energized gasses produce light of distinct wavelengths
Electrons falling from energized states
Niels Bohr
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.
Balmer
Wrote eqn for hydrogen emission spectra that evolved into Rydberg equation
Trial and error
3 things that describe light
wavelength (lamda)
frequency (v)
amplitude
Rutherford
Gold foil experiment
Dense nucleus, electron cloud
Louis de Broglie
Matter waves
Wavelength=h/momentum
Returns to Bohr postulate
Proof: Electrond diffract through crystal
Schrodinger
Equations to find probabilities of electron location for each orbital: Quantum/wave mechanics
Probability density: square of wave function
Heisenburg
Uncertainty principle: Limit to knowledge of electrons momentum and location
Things explained by Quantum model for electrons
Periodic Table
Colors of Transition metal aqeus ions
Bonding patterns of ions
Bond angles
Paramagnetic behavior of some atoms
Valence electrons
Outside Noble gas core
Ionization energy
Amount of energy to remove one electron (valence)
Degenerate orbitals
Orbitals that have the same energy
Aufbau process
Electrons always fill the lowest energy orbitals first (violated with some transition metals - Cr)
Hund's rule
For degenerate orbitals, one electron each, then double up
Pauli exclusion principle
No two electrons can have the same four quantum numbers
"n"
Positive integers, Defines energy of electron
Principal quantum number
"l"
0-(n-1), Defines shape of orbital
Angular momentum (azimuthal) quantum number
ml
-l...0...+l, defines orientation of orbital in space
distingishes between orbitals in the same subshell
Magnetic quantum number
ms
-1/2 or +1/2, spin up of spin down
electron shell
subshell
Orbitals with the same value of n
Orbitals with the same n and l values
node
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
Valence rules
For representative elements, d and f orbital electrons do not count as valence electrons.
For transition metals, f orbital electrons don't count.