Electrons in Atoms
Ch. 13
Models of the Atom
13-1
The evolution of Atomic Models
•  Dalton (1766-1844): atom indivisible
•  J.J. Thomson (1856-1940):
–  “Plum-pudding” model – negative electrons
stuck in positively charged material
•  Rutherford (1871-1937):
–  Electrons surround dense nucleus, rest of atom
is empty space
•  Bohr (1885-1962):
–  “Planetary model,” electrons fixed in energy
levels around nucleus
The Quantum Mechanical
Model
•  Quantum Mechanical Model:
– Estimates the probability of finding an
electron in a certain place using the
Schrodinger equation.
– “Fuzzy cloud” model; where the
cloud is more dense the probability of
finding the electron is high, where the
cloud is less dense the probability is
low.
Energy Levels
•  Electrons move around the
nucleus in energy levels.
•  Quantum of energy = amount of E
required to move to a higher level.
•  When they move away from the
nucleus (up a level) they require
energy.
•  When they move towards the
nucleus (down a level) they
release energy
•  The farther from the nucleus the
energy level is, the more energy is
required to move up a level (away
from nucleus).
Energy Levels
•  Quantum number (n) refers to an energy level
–  n = 1, 2, 3, 4, …7, values increase going away
from nucleus.
–  Each energy level fits a certain amount of
electrons:
•  Level 1 = 2 electrons
•  Level 2 = 8 electrons
•  Level 3 = 18 electrons
•  Level 4 = 32 electrons
Sublevels + Orbitals
•  Within each energy level there are sublevels; the
number of sublevels is equal to the quantum
number.
–  Ex: Energy level 4 has 4 sublevels within it.
•  A sublevel is made up of atomic orbitals: s, p, d, f
–  Orbital s fits 2 electrons total
–  Orbital p fits 6 electrons total
–  Orbital d fits 10 electrons total
–  Orbital f fits 14 electrons total
–  s fills up first, then p, then d, then f
Energy
Level (n)
Sublevel/
Orbital
1
1s
2
2s
2p
3s
3p
3d
4s
4p
4d
4f
3
4
Electrons in
Total # of
each Sublevel/ electrons in
Orbital
Level
2
2
2
6
2
6
10
2
6
10
14
8
18
32
Atomic Orbitals
Orbital
Shape
# of Electrons
s
Spherical
2
p
Dumbbell
6
d
Clover-leaf
10
f
Complex
14
f - orbital
s - orbital
d - orbital
DRAW!
p - orbital
Electron Arrangement in
Atoms
13-2
d-1
f-2
On blank periodic table: write the configuration of
each element, color the 4 sublevels + make a key
Electron
Arrangement in Atoms
• 
• 
• 
• 
• 
• 
• 
• 
Period 1 - 1s2
Period 2 - 2s22p6
Period 3 - 3s23p6
Period 4 - 4s23d104p6
Period 5 - 5s24d105p6
Period 6 - 6s24f145d106p6
Period 7 - 7s25f146d107p6
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p67s25f146d107p6
Electron Configuration Notation
•  Notation used to represent electron
configurations:
– H: 1s1 # of electrons in sublevel/orbital
Energy level
Sub level/orbital
– He: 1s2
– Li: 1s2 2s1
– Be: 1s2 2s2
– B: 1s2 2s22p1
d-1
f-2
You Try!
• 
• 
• 
• 
• 
Write the electron configurations for the following:
•  Ca (20):
C (6):
22s22p63s23p64s2
1s
2
2
2
1s 2s 2p
•  Ir (77):
F (9):
22s22p63s23p64s2
1s
2
2
5
1s 2s 2p
104p65s24d105p66s2
3d
Ne (10):
145d7
4f
2
2
6
1s 2s 2p
•  Cm (96):
Na (11):
22s22p63s23p64s2
1s
2
2
6
1
1s 2s 2p 3s
104p65s24d105p66s2
3d
P (15):
145d106p67s25f8
4f
2
2
6
2
3
1s 2s 2p 3s 3p
d-1
f-2
•  Abbreviated form: shows preceding noble gas and
the configuration of only the last energy level!
–  Mg: 1s2 2s22p63s2
•  or [Ne] 3s2
–  B: 1s2 2s2 2p1
•  or [He] 2s2 2p1
–  Si: 1s2 2s2 2p6 3s2 3p2
•  or [Ne] 3s2 3p2
–  Al: 1s2 2s2 2p6 3s2 3p1
•  [Ne] 3s2 3p1
–  Xe: 1s2 2s2 2p6 3s23p64s23d104p65s24d105p6
•  [Kr] 5s24d105p6
d-1
f-2
•  What happens in the fourth period?
– After 4s2, comes 3d10, then 4p6
– Scandium (#21): 1s2 2s2 2p6 3s2 3p6 4s2
3d1
• or [Ar] 4s2 3d1
– Copper is [Ar] 4s2 3d9
– Bromine is [Ar] 4s2 3d10 4p5
•  What happens in the sixth period?
– After 6s2, comes 4f14, then 5d10, then 6p6
– Tungsten (W) is [Xe] 6s24f145d4
Orbital Notation Rules
1)  Aufbau principle: electrons enter orbitals
of lowest energy first.
2)  Pauli exclusion principle: an atomic orbital
may describe at most 2 electrons.
3)  Hund’s rule: one electron enters each
orbital until ALL orbitals contain 1 electron
with parallel spins.
•  Why is 3d on the
4th row after 4s?
•  3d has ______ energy than 4s
and ______energy than 5p.4p
•  Why is 4f on the 6th row after 6s?
•  4f has ______ energy than 6s
and ______energy than 5d
•  Aufbau principle: lowest energy orbitals are
filled first!
•  Sublevel order =
•  1s,2s,2p,3s,3p,4s,4p,5s,4d,5p,6s,4f,5d,6p
Orbital Notation
O
8
1s22s22p4
Light and Atomic Spectra
13-3
(only pg. 372-375)
Electromagnetic Spectrum
•  Energy in the form of electromagnetic
radiation (radiant energy) travels in waves
•  Waves transfer the energy from one place
to another
•  Ex: radio waves, TV, microwave, visible
light, x-rays, gamma rays, infrared, UV
•  All forms of radiant energy are part of the
electromagnetic spectrum
Low energy
High Energy
Wavelength + Frequency
•  Two main properties of electromagnetic
waves:
1)  Frequency
2)  Wavelength
•  Wavelength is the distance between two
corresponding peaks or troughs.
•  Frequency is the number of wave cycles per
second.
•  Wavelength is inversely proportional to
frequency
•  wavelength
frequency
Wave length
Frequency
Wave length
Frequency
•  Higher frequency waves (short wavelength)
have high energy
–  Ex: gamma rays, x-rays, ultraviolet rays
–  Ex: Violet light in visible spectrum
•  Low frequency waves (long wavelength)
have low energy
–  Ex: radio waves, microwaves, infrared (heat)
waves
–  Ex: Red light in visible spectrum
Light and Atomic Spectra
•  Electrons absorb energy and move to
higher energy states/levels
•  Electrons give off that energy in the form
of light when they fall back down to lower
energy states, or ground state.
•  ALL electromagnetic waves travel at the
speed of light in a vacuum – 300 million
meters per second or (3.0 x 108 m/s)
• When atoms are
energized by an electric
current they emit light.
• When this light is
passed through a prism
they produce an
emission spectrum.
• Each element has its
own unique atomic
emission spectrum
fingerprint
http://phys.educ.ksu.edu/vqm/html/emission.html