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Chapter 4
Arrangement of Electrons in
Atoms
I. The Development of a New
Atomic Model
 Electromagnetic Radiation:
 Electromagnetic Spectrum:
 Wavelength (): corresponding points
on adjacent waves---Ex:
Frequency (): # of waves that pass a
point in a specific time
 c = () () ------inversely proportional
 c = () () ------inversely proportional
c : m/s
 : m, cm, nm
 : waves/second--Hertz (Hz)
 Photoelectric Effect: emission of e- by
certain metals when light shines on
them
 Quantum: min quantity of nrg that can
be lost or gained by an atom
 E = (h) ()
o J = (Js) (Hz)
o Planck’s constant: 6.626 X 10-34 Js
•
Einstein
dual wave-particle to describe light
 Photon: radiation with zero mass
carrying a quantum of nrg
o
o
packet of nrg emitted when an e- drops nrg
levels
 Ground state: lowest nrg state
 Excited state: higher potential nrg
--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
Line-Emission Spectrum
excited state
ENERGY IN
PHOTON OUT
ground state
Bohr Model
65
4
3
2
1
-Energy of photon
depends on the
difference in energy
levels
-Bohr’s calculated
energies matched
the IR, visible, and
UV lines for the H
atom
Other Elements
 Each element has a unique bright-line
emission spectrum.
 “Atomic Fingerprint”
Helium
Bohr’s calculations only worked for
hydrogen! ----pg 97
II. The Quantum Model of the
Atom
 A. Electrons as Waves
o Diffraction: bending of a wave as it passes
by the edge of an object
o Interference: results when waves overlap
EVIDENCE: DIFFRACTION PATTERNS
VISIBLE LIGHT
ELECTRONS
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 e-
Ψ 1s 

1 Z 3/2 σ
π a0
e
A. Atomic Orbitals and
Quantum Numbers
 Orbital: probable location of an e Quantum #: properties of atomic
orbitals and properties of e-’s in
orbitals
 Principal quantum #: (n), indicates
main nrg level occupied by the eo n = 1 -----occupies 1st nrg level
 Angular momentum quantum #: (l),
indicates shape of orbital
 Magnetic quantum #: (m), orientation
of an orbital
 Spin quantum #: which spin state (+)(-)
 ***See table 4-2 pg 104
Orbital (“electron cloud”)
Region in space where there is 90% probab
Orbital
Radial Distribution Curve
Four Quantum Numbers:
Specify the “address” of each electron in an
UPPER LEVEL
1. Principal Quantum Number ( n )
Energy level
Size of the orbital
n2 = # of orbitals in
the energy level
2. Angular Momentum Quantum # ( l )
Energy sublevel
Shape of the orbital
s
p
d
f
n =
# of sublevels per level
n2 =
# of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbital
within each sublevel
4. Spin Quantum Number ( ms )
Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in
opposite directions.
III. Electron Configuration
 Aufbau principle: lowest nrg orbits fill
first
 Pauli exclusion principle: no 2 e-’s can
have the same 4 quantum #’s. This is
where spin allows 2 e-’s to be in the
same orbit
o Ex: 
 Hund’s rule: orbital of equal nrg are
occupied by 1 e-, before any is
occupied by 2 e-’s
o Ex:   
 Orbital Notation:
ex: pg 107
 Electron Config Notation: pg 107
 Electron Dot diagram: ex
 Noble gases:
 are inert
 complete octet
 --show ex----
 Table 4-3 pg 110
1. Principal #  energy level
2. Ang. Mom. # sublevel (s,p,d,f)
3. Magnetic #  orbital
4. Spin #
 electron
Feeling overwhelmed?