Chapter 4 Arrangement of
Electrons in Atoms
4.1 The Development of a New
Atomic Model
Properties of Light
• Electromagnetic Radiation:
EM radiation are forms of energy which move
through space as waves
• There are many different types of EM waves
• visible light
• x-rays
• ultraviolet light
• infrared light
• radio waves
EM Waves
• Move at speed of light: 3.00 x 108 m/s
• Speed is equal to the frequency times the
wavelength c = v
• Frequency (v) is the number of waves passing
a given point in one second
• Wavelength () is the distance between peaks
of adjacent waves
• Speed of light is a constant, so v is also a
constant; v and  must be inversely proportional
Light and Energy:
The Photoelectric Effect
• Electrons are emitted from a
metal when light shines on the
metal
• Incoming EM radiation from
the left ejects electrons,
depicted as flying off to the
right, from a substance.
• Radiant energy is transferred
in units (or quanta) of energy
called photons
(Max Planck)
Photoelectric Effect
energy
p+
absorption
spectrum
no
e-
ground
state e-
When a specific or
quantized amount of
energy is exposed to
the atom, the electron
jumps from its “ground”
or original state to an
“excited” state
Photoelectric Effect
When the “excited”
electron returns to
lower energy levels,
it releases energy in
the form of light
energy
photon
p+
no
e-
emission
spectrum!
travels at the
speed of light
excited
state e-
(3.00 x 108 m/s)
• A photon is a particle of energy having a
rest mass of zero and carrying a quantum
of energy
• A quantum is the minimum amount of
energy that can be lost or gained by an
atom
• Energy of a photon is directly proportional
to the frequency of radiation
• E = hv (h is Planck’s constant,
6.62554 x 10 -24 J * sec)
Electromagnetic Spectrum
• Wavelength increases→
• Frequency decreases→
• Energy decreases→
Electromagnetic Spectrum
Wave-Particle Duality
• Energy travels through space as
waves, but can be thought of as a
stream of particles (Einstein)
• Each particle has 1 quantum of
energy.
Line Spectrums
• Ground State: The lowest energy state of an atom
• Excited State: A state in which an atom has a
higher potential energy than in its ground state
example: Neon lights
Emissions
Spectrum
• Bright line spectrum: Light is given off by excited
atoms as they return to lower energy states
• Light is given off in very definite wavelengths
• A spectroscope reveals lines of particular colorslight passed through a prism; specific frequencies
given off.
The Hydrogen Line Spectrum
• Definite frequency
• Definite wavelength
http://student.fizika.org/~nnctc/spectra.htm
Bohr Model
e-
p+
Energy
levels
Niels Bohr
no
Electrons circle
around the nucleus
on their energy level
The Bohr Model of the Atom
• Electron Orbits, or Energy Levels
• Electrons can circle the nucleus only in allowed
paths or orbits
• The energy of the electron is greater when it is in
orbits farther from the nucleus
• The atom achieves the ground state when atoms
occupy the closest possible positions around
the nucleus
• Electromagnetic radiation is emitted when
electrons move closer to the nucleus.
The Bohr Atomic Model
Energy transitions
• Energies of atoms are
fixed and definite
quantities
• Energy transitions
occur in jumps of
discrete amounts of
energy
• Electrons only lose
energy when they
move to a lower
energy state
Shortcomings of the Bohr Model
• Doesn't work for atoms larger than
hydrogen
(more than one electron)
• Doesn't explain chemical behavior
Chapter 4 Arrangement of
Electrons in Atoms
4.2 The Quantum Model of the
Atom
Electrons as Waves and Particles
• Louis deBroglie (1924)
• Electrons have wavelike properties
• Consider the electron as a wave confined to
a space that can have only certain
frequencies
The Heisenberg
Uncertainty Principle
"It is impossible to determine
simultaneously both the
position and velocity of an
electron or any other
particle.”
Werner Heisenberg- 1927
• Electrons are located by their interactions with photons
• Electrons and photons have similar energies
• Interaction between a photon and an electron knocks the
electron off of its course
The SchrØdinger
Wave Equation
• Proved quantization of electron
energies and is the basis for
Quantum Theory
• Quantum theory describes
mathematically the wave properties
of electrons and other very small
particles
• Electrons do not move around the
nucleus in "planetary orbits"
????????
The SchrØdinger
Wave Equation
• Electrons exist in regions called orbitals
• An orbital is a three-dimensional region
around the nucleus that indicates the
probable location of an electron
• SchrØdinger equation for probability of a
single electron being found along a single
axis (x-axis)
Atomic Orbitals & Quantum Numbers
• Quantum Numbers specify the properties
of atomic orbitals and the properties of the
electrons in orbitals:
Principal Quantum Number (n)
Angular Momentum Quantum Number (l)
Magnetic Quantum Number (m)
Spin Quantum Number
Principal Quantum Number (n)
• Indicates the main energy levels occupied
by the electron
• Values of n are positive integers
• n = 1 is closest to the nucleus, and lowest
in energy
• The number of orbitals possible per
energy level (or "shell") is equal to n2
Quantum Theory
Energy
levels, n
en=1
p+
no
n=2
n=3
n=4
Angular Momentum Quantum Number (l)
•
•
•
•
Indicates the shape of the orbital
Number of orbital shapes = n
Possible values are l = 0, 1, 2, or 3
Shapes are designated s, p, d, f
Click Here!
• S shape is spherical
• P shape is a dumbbell, or figure 8
Magnetic Quantum Number (m)
• The orientation of the orbital around the nucleus
• s orbitals have only one possible orientation,
m=0
• p orbitals have three possible,
m = +1, 0 or -1
• d orbitals have five possible,
m = -2, -1, 0, +1, or +2
• f orbitals have 7 possible orientations
Spin Quantum Number
• Indicates the fundamental spin states of
an electron in an orbital
• Two possible values for spin, +1/2, -1/2
• A single orbital can contain only two
electrons, which must have opposite spins
Diagrams of the Orbitals
Summary of the Quantum Numbers
Chapter 4 Arrangement of
Electrons in Atoms
4.3 Electron Configurations
Energy Levels
Energy Levels
(7)
Sublevels
(4)
s
1 orbital
p
3 orbitals
2 e-
d
5 orbitals
6 e-
f
7 orbitals
10 e-
14 e-
Writing Electron Configurations
Pauli
Exclusion
Principle
No two
electrons in
the same
atom can
have the
same set of
four quantum
numbers
Rules:
Hund's Rule
Orbitals of equal
energy are each
occupied by one
electron before any
orbital is occupied
by a second
electron, and all
electrons in singly
occupied orbitals
must have the same
spin
Aufbau
Principle
An electron
occupies the
lowest-energy
orbital that
can receive it
Orbital Notation (1 of 3)
• Unoccupied orbitals are represented by a
line, _____
• Lines are labeled with the principal
quantum number and the sublevel letter
• Arrows are used to represent electrons
• Arrows pointing up and down indicate
opposite spins: Pauli Exclusion Principle
No two electrons in the same
atom can have the same set
of four quantum numbers
(occupy the same space @
the same time)
Writing Electron
Configurations
1st e-
2nd e-
Hund's Rule
Orbitals of equal energy are each
occupied by one electron before
any orbital is occupied by a
second electron, and all electrons
in singly occupied orbitals must
have the same spin
3rd e-
4th e-
Configuration Notation (2 of 3)
• The number of electrons in a sublevel is
indicated by adding a superscript to the
sublevel designation
• Hydrogen = 1s1
Aufbau Principle
• Helium =
1s2
An electron occupies the
• Lithium =
1s2 2s1
lowest-energy orbital
that can receive it
(Always start with n = 1
and work your way up)
1 2
3 4 5
6 7
8 9 10 11 12 13 14 15 16 17 18
Sublevel Blocks on
the Periodic Table
Fill-In the Periodic Table using Energy Levels & Sublevels
Orbital Filling Order
START
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
FINISH
Noble Gas Notation (3 of 3)
• The configuration begins with the
preceding noble gas’s symbol in brackets
and is followed by the rest of the
configuration for the particular element.
• [Ne] 3s23p5
Terms
• Highest occupied energy level:
The electron containing energy level with
the highest principal quantum number
• Inner shell electrons:
Electrons that are not in the highest
energy level
• Octet Rule:
Highest energy level s and p electrons are
filled (8 electrons)
Octet Rule
• Characteristic of noble gases, Group 18
• Exceptions: Hydrogen & Helium
• Noble gas configuration:
Outer main energy level fully occupied,
usually (except for He) by eight electrons
This configuration has extra stability
Survey of the Periodic Table
Elements of the Fourth Period
• Irregularity of Chromium
• Expected:
• Actual:
1s22s22p63s23p64s23d4
1s22s22p63s23p64s13d5
Rule:
Sublevels are most stable when
they are either half or
completely filled. Electrons will
shift to different energy levels to
accommodate this stability
whenever possible.
Survey of the Periodic Table
• Several transition and rare-earth elements
borrow from smaller sublevels in order to
half fill larger sublevels
i.e. d borrows 1 e- from s
• This accounts for some of the unexpected
electron configurations found within the
transition elements.
Magnetic Fields & Electrons
• Paramagnetic:
When an atom has unpaired electrons, it
will be attracted into a magnetic field
Ex: 1s22s22p2
• Dimagnetic:
When an atom has only paired
electrons, it will be slightly repelled by a
magnetic field
Ex: 1s22s22p63s2