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)

Blackbody radiation

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.

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