quantum chem

spectral lines

light emitted from an excited element

no range of frequencies

no 2 heated elements glow the same color, explained by specific electron orbits (electrons jump energy levels)

energy level

low energy- close to nucleus (ground state)

high energy- further to nucleus (excited state)

farther fall, greater energy released

nanometers

meters x 10^-9

E = h v

in joules/mol

heisenberg uncertainty principle

the exact location and momentum of an electron cannot be determined

quantum numbers

1. principal quantum (energy levels, n = 1, 2, … 8)

2. orbital number (shape of cloud/subshell, s p d f)

3. magnetic (3d orientation of orbital, s = 1 orbital, p = 3, etc)

4. spin (counter/clockwise, up down arrows)

*energy level corresponds to number of orbital shapes

degenerate orbitals

no difference in energy

pauli exclusion principle

there cannot be more than 2 electrons in one orbital

max 2 electrons in same orbital if they have opposite spin

4 unique quantum numbers

aufbau principle

electrons occupy lowest energy orbitals available, predictable pattern

hund’s rule

each orbital must have one electron before filling second spot for electron

all single electrons must have same spin

ions

symmetrical distribution favored compared to electron repulsion, allowing for a lower energy configuration

lower energy configuration favored, assymetrical is higher energy configuration

remove electrons from highest principal number energy level, even if its not the highest energy sublevel

shielding effect

core electrons repel outermost valence electrons, reducing attraction to nucleus

removing an electron increases attraction

#energy levels = same as period