Name_____________________________ Electron configuration – NOTES
Period_________
A. Waves



Wavelength () - length of one complete wave
Frequency () - # of waves that pass a point during a certain time period (measured in hertz
Amplitude (A) - distance from the origin to the trough or crest
L
O
W
H
I
G
H
E
N
E
R
G
Y
E
N
E
R
G
Y

Frequency & wavelength are inversely proportional



c:
:
:
speed of light (3.00  108 m/s)
wavelength (m, nm, etc.)
frequency (Hz)
c = 
B. Quantum Theory

Planck (1900)




Observed - emission of light from hot objects
Concluded - energy is emitted in small, specific amounts (quanta)
Quantum - min imum amount of energy change
Einstein (1905)
 Observed - photoelectric effect
 Concluded - light has properties of both waves and particles
“wave-particle duality”
 Photon - particle of light that carries a quantum of energy
 The energy of a photon is proportional to its frequency.
 E:
energy (J, joules)
 h:
Planck’s constant (6.6262  10-34 J·s)
 :
frequency (Hz)
C. Bohr’s Model
 e- exist only in orbits with specific amounts of energy
called energy levels
 Therefore…
 e- can only gain or lose certain amounts of energy
 only certain photons are produced
E = h
(Hz) = 1/s)
D. Line-Emission Spectrum
Quantum Model of the Atom
A. Electrons as waves


Louis de Broglie (1924)
 Applied wave-particle theory to e e- exhibit wave properties
Heisenberg Uncertainty Principle
 Impossible to know both the velocity and position
of an electron at the same time
B. Quantum mechanics



Heisenberg Uncertainty Principle
 Impossible to know both the velocity and position of an electron at the same time
Schrödinger Wave Equation (1926)
 finite # of solutions  quantized energy levels
 defines probability of finding an eOrbital (“electron cloud”)
 Region in space where there is 90% probability of finding an e-
C. Quantum numbers
Four Quantum Numbers: Specify the “address” of each electron in an atom
1. Principal Quantum Number ( n )
 Energy level
 Size of the orbital
 n2 = # of orbitals in / 2n2 = # e- in level
 the energy level
2. Angular Momentum Quantum # ( l )
 Energy sublevel
 Shape of the orbital

Orbitals combine to form a spherical shape.
4. Spin Quantum Number ( ms )
 Electron spin  +½ or -½
 An orbital can hold 2 electrons that spin in opposite directions.

Pauli Exclusion Principle
 No two electrons in an atom can have the same 4 quantum numbers.
 Each e- has a unique “address”:
Electron Configurations
A. General Rules

Pauli Exclusion Principle
 Each orbital can hold
TWO electrons with opposite spins.

Hund’s Rule
 Within a sublevel, place one e- per
orbital before pairing them.
 “Empty Bus Seat Rule”

Electron Filling order
Aufbau Principle
 Electrons fill the lowest energy orbitals first.
B. Notation – EXAMPLE (Oxygen)

Orbital diagram

Electron Configuration
Each s sublevel has _1_ orbital and can hold __2___ e-.
Each p sublevel has _3_ orbitals & can hold __6___ e-.
Each d sublevel has _5_ orbitals & can hold __10__ e-.
Each f sublevel has _7_ orbitals & can hold __14__e-.
EXAMPLE (SULFUR)

Longhand Configuration

Shorthand Configuration
C. Period Patterns




Period #
 energy level (subtract for d & f)
A/B Group #
 total # of valence eColumn within sublevel block
 # of e- in sublevel
Shorthand Configuration
 Core e-: Go up one row and over to the Noble Gas.
 Valence e-: On the next row, fill in the # of e- in each sublevel.
 Example: Germanium
D. Electron Configuration exceptions:
Example:
Write the electron configuration and orbital notation of a neutral atom of Arsenic (As).