Studied by 4 peopleRed light
Lowest energy (700-650 nm)
Orange light
Low energy (600 nm)
Yellow light
Medium energy (550 nm)
Green light
Medium energy (500 nm)
Blue light
High energy (450 nm)
Violet light
Highest energy (400 nm)
Subshells
Have the same principal and angular momentum quantum numbers (n & l)
Shells
Have the same principal quantum number (n)
S orbital
Spherical
Px orbital
Two ovals on the x-axis
Py orbital
Two ovals on the y-axis
Pz orbital
Two ovals on the z-axis
One orbital
The s subshell has
3 orbitals
The p subshell has
5 orbitals
The d subshell has
7 orbitals
The f subshell has
Aufbau principle
Electrons fill lowest energy subshells first
Pauli exclusion principle
No two electrons have the exact same quantum numbers
Hund’s rule
Electrons fill subshells (with the same spin) before pairing
Valence electrons
The electrons in the last filled subshell are…
Diamagnetic
All electrons are paired
Paramagnetic
One or more unpaired electrons
Cr & Mo
Take one electron from the s subshell and put it into the d subshell, making d half-full
Cu, Ag, & Au
Take one electron from the s subshell and put it into the d subshell, making d full
La & Yb
Take one electron from 4f to create 5d
Ac & No
Take one electron from 5f to create 6d
Ground state
Lowest energy
Excited state
Higher energy than the ground state
Dyz
Dxz
Dz^2
Dxy
Dx^2-y^2
Balmer series
When visible light is produced (falls to n=2)
Lyman series
When energy levels fall to n=1
Paschen series
When energy levels fall to n=3
Brackett series
When energy levels fall to n=4
Pfund
When energy levels fall to n=5
Note
Flashcard