Chapter 5 - Electrons In Atoms
5.1 Light and Quantized Energy
5.2 Quantum Theory and the Atom
5.3 Electron Configurations
Section 5.1 Light and Quantized Energy
Light, a form of electronic radiation, has
characteristics of both a wave and a
particle.
• Know the names and how to recognize the parameters that
describe an electromagnetic wave (period, wavelength,
amplitude, frequency, crest, etc.)
• Perform calculations involving c = 
• Describe the experiments and the interpretation of both
Planck’s experiment and the photoelectric effect.
• Compare the wave and particle natures of light.
Section 5.1 Light and Quantized Energy
• Define a quantum of energy and explain how it is related to
an energy change of matter.
• Contrast continuous electromagnetic spectra and atomic
emission spectra and provide examples of each.
Section 5.1 Light and Quantized Energy
Key Concepts
• All waves are defined by their wavelengths, frequencies
(periods), amplitudes, and speeds.
• In a vacuum, all electromagnetic waves travel at the speed
of light. The relationship between this speed and the
wavelength and frequency is c = λν
• All electromagnetic waves have both wave and particle
properties. This is known as wave-particle duality.
• Matter emits and absorbs energy in quanta. The amount of
energy can be calculated using Equantum = hν
• White light produces a continuous spectrum. An element’s
emission spectrum consists of a series of lines of
individual colors.
Nuclear Atom Unanswered Questions
Rutherford model said that:
• Positive charge and virtually all mass
concentrated in nucleus
• Most of volume occupied by electrons
But still unknown
• Spatial arrangement of electrons?
• Why doesn’t positive nucleus pull
negative electrons into itself?
How to Explain Chemical Behavior?
Elements with AN = 17, 18, 19
• Chlorine – highly reactive gas
• Argon – unreactive gas
• Potassium – very reactive metal
Wave Nature of Light
Electromagnetic “radiation”
• Exhibits wavelike behavior
• Consists of oscillating (periodically
varying) magnetic and electric fields
• Oscillations  to direction of travel
Moving Electromagnetic Wave
Magnetic Field
Electric
Field
Propagation
direction
Wavelength ()
Wave Characteristics
Wavelength
 (lambda)
m
Frequency
Period
 (nu)
T
Hz
s
Speed
c (for EM waves; m/s
speed of light)
Amplitude
[symbol not used for
this course]
Wavelength 
Shortest distance between equivalent
points on a continuous wave
View of wave frozen in space: x axis - distance
Wavelength 
Has unit of length
•m
• cm
• nm
• etc.
Frequency 
Number of waves that pass given point
per second
• SI unit: Hz


Waves per second
In calculations, use 1/s or s-1
Period (T)
• T = 1/ 
• Units of time (s)
• Time between crests for wave sampled
over time at one point in space
Wave period

Period 
View of wave sampled over time at 1 point
in space: x axis = time (oscilloscope trace)
EM Wave Speed
If electromagnetic (EM) wave in
vacuum, travels at speed of light, c
• 3.00  108 m/s
All EM waves in vacuum travel at c
• Doesn’t vary with wavelength,
frequency, or amplitude of wave
For all practical purposes, air is
equivalent to a vacuum for EM wave
Frequency / Wavelength
For EM waves in vacuum
c=
c constant , so as  increases,  must
decrease and vice versa
Frequency / Wavelength
c=
Longer

Red light versus blue light
Shorter Lower Frequency

Higher Frequency
Amplitude
Wave height measured from origin to
crest (or trough)
• Not concerned about units for
amplitude in this course
EM Wave Energy
Energy increases with increasing
frequency (or decreasing wavelength)
High energy
• Gamma rays
• X-rays
Low energy
• Radio waves
Electromagnetic Spectrum
Visible Light
104
Frequency (Hz)
Increasing Energy
1022
Calculating Wavelength
Practice problem 5.1, page 140
Wavelength () of microwave having
frequency () = 3.44109 Hz?
c=
=c/
= 3.00108 m/s
3.44108 s-1
= 8.7210-2 m
Practice
Electromagnetic Waves
Practice problems, page 140
Problems 1-4
Chapter assessment page 166
Problems 34 - 36 (concepts)
Problems 45 – 48
Supplemental Problems 1-2, page 978
History of Development of
Human Understanding of the Atom
- +
+ +
- + + - +
-
-
+++
++ -
Dalton Thomson Rutherford
-
+
+++++ -
-
Bohr
-
+
+++++ -
-
Quantum
Pre 1900 View of Matter & Energy
Matter and EM energy are distinct
• Matter consists of particles
Have mass
 Position in space can be specified

• Energy in form of EM radiation is wave
Massless
 Delocalized – position in space can’t be
specified

Pre 1900 View of Matter & Energy
View of energy also included idea that
energy can be gained or lost in
continuous manner
• Heating water up with Bunsen burner –
can change temperature by any arbitrary
amount simply by changing size of flame
and/or heating time
Pre 1900 View of Matter & Energy
Attempt to explain results of 2 key
experiments challenged neat
particle/wave division and idea of
continuous energy transfer
• Change in intensity and wavelength of
radiation emitted by heated objects as a
function of temperature
• Emission of electrons from metals when
certain frequencies of light shown on
metal
EM Radiation of Heated Objects
Heated objects emit differing
wavelengths of light depending
upon temperature
Max Planck measured (~1900)
energy vs wavelength profile at
different temperatures
EM Radiation of Heated Objects
Predictions of
classical theory
based on wave
model fails to
explain intensity
profile
Quantum Nature of Energy
Based on analysis of heated object
data, Planck concluded that matter can
gain or lose energy only in small
increments call quanta
Single quantum minimum amount of
energy atom can lose or gain
Quantum Nature of Energy
Energy of quantum related to frequency
of radiation
Equantum = h 
h = Planck’s constant = 6.626 x10-34 J s
If know energy, can get frequency
Consistent with idea that high
frequency EM radiation has high
energy
Quantum Nature of Energy
Equantum = h 
Above gives magnitude of quantum of
energy
Planck’s theory:
Energy transfer only happens in
increments = integer multiple of Equantum
Etransfer = n h 
n = 1, 2, …
Photoelectric Effect
Measure number and energy of
electrons ejected when light of certain
wavelength strikes metal surface
Electron
ejected
from
surface
Electrons
Beam of light
Metal surface
Nuclei
Photoelectric Effect
Photoelectric Effect
Classical wave model (continuous
energy) prediction:
• Given enough time, low frequency (=
low energy) light will eventually transfer
enough energy to metal to eject an
electron
• Analogous to heating water model
Actual (quantum) results
• No electrons unless  > threshold
Photoelectric Effect
Metal only ejects electrons when energy
(frequency) of light above minimum
threshold
Even dim light above threshold ejects
electrons
Increasing intensity of light of given
frequency above threshold ejects more
electrons but energy of electrons same
Increasing light frequency above threshold
ejects higher energy electrons but not more
electrons
Photoelectric Effect
Basis for Einstein proposal (1905) that
photons have both a wavelike and
particle nature
Einstein extended Planck’s quantum
idea to photons
Ephoton = h 
Photon = particle (packet) of EM
energy with no mass that carries a
quantum of energy
Wave/Particle Views
Light as a wave phenomenon
Light as a stream of photons
Calculating Energy of Photon
Practice problem 5.2, page 143
Energy of photon from violet part of
rainbow if  = 7.231014 s-1?
Ephoton= h 
h = 6.626 x10-34 J s
Ephoton= (6.626 x10-34 J s)(7.231014 s-1)
= 4.7910-19 J
Practice
Photon Energy
Practice problems, page 143
Problems 5 - 7
Section assessment page 145
Problem 13
Chapter assessment page 166
Problems 37 - 40 (concepts)
Problems 49 – 53, 55 – 57
Supplemental, page 978 Problems 3 - 4
Atomic Emission Spectra
The fact that only certain colors are
seen in fireworks, neon signs, etc. is a
further indication that energy can’t
come out of an atom in arbitrary
amounts
Atomic Emission Spectra
Fact that only
certain colors
seen means only
certain distinct
frequencies are
possible
Energy = n h 
Strontium
Copper
Continuous Spectrum
Discrete Spectrum
Atomic Emission Spectra
Use optical spectroscopes and diffraction glasses
to view:
• Light from fluorescent tubes
• Incandescent light from overhead projector
• Emission from spectrum tubes (H, He)
Atomic Emission Spectra of
Hydrogen
Hydrogen gas discharge tube
Prism
Slit
Spectrum
Hg, Sr, H Spectra Comparison
Emission / Absorption Spectra - Na
Emission / Absorption Spectra – He
Fig 5.9
Chapter 5 - Electrons In Atoms
5.1 Light and Quantized Energy
5.2 Quantum Theory and the Atom
5.3 Electron Configurations
Section 5.2 Quantum Theory and the Atom
Wavelike properties of electrons help relate
atomic emission spectra, energy states of
atoms, and atomic orbitals.
• Compare the Bohr and quantum mechanical models of the
atom.
• Describe the process of atomic emission and calculate the
wavelength of an emitted photon given the energy levels
• Explain the impact of de Broglie's wave particle duality and
the Heisenberg uncertainty principle on the current view of
electrons in atoms.
• Identify the relationships among a hydrogen atom's energy
levels, sublevels, and atomic orbitals.
Section 5.2 Quantum Theory and the Atom
Key Concepts
• Bohr’s atomic model attributes hydrogen’s emission
spectrum to electrons dropping from higher-energy to
lower-energy orbits.
∆E = E higher-energy orbit - E lower-energy orbit = E photon = hν
• The de Broglie equation relates a particle’s wavelength to
its mass, its velocity, and Planck’s constant.
λ = h / mν
• The quantum mechanical model of the atom assumes that
electrons have wave properties.
• Electrons occupy three-dimensional regions of
space called atomic orbitals.
Rutherford Model of Atom - Limitations
Accelerated charged particles will radiate
EM waves
Electron in orbit; therefore accelerated
Should lose energy and spiral inwards to
nucleus
This doesn’t
happen!
EM waves
Bohr Model of Atom
1913, Niels Bohr proposed that H
atom has only certain allowable
(quantized) energy states
Lowest state = ground state
Energy gains promote electrons in
atom to excited state
Electrons confined to distinct
circular orbits
Smaller orbits  lower energy
Bohr’s Model
(nucleus way out of scale)
Nucleus
Electron
Orbit
Energy Levels  orbit radius
Quantum
Staircase
Bohr Atom Picture of Energy Transfer
Electron: Red Photon: Orange
Photon Absorption
Photon Emission
Emission Spectrum of Hydrogen
Bohr Model of Atom
3 Series of Atomic Emission Lines
Visible Series (Balmer) final state n=2
UV Series
(Lyman) final
state n=1
Infrared Series
(Paschen) final
state n=3
Bohr Energy Formula (not in book)
Energy change = D E
= E (final) – E(initial)
= E photon
=h
 1
E n   R H 
2
n



RH = Rydberg constant
= 2.18  10-18 J
E < 0 (Energy at infinite separation = 0)
DE  E f  E i  h
Bohr Freq. Formula (not in book)
Energy transitions for
hydrogen’s four visible
spectra lines - Balmer
series


D E  R H  1
1 
 




2
2
h
h

 n
n 
f 
 i
Bohr Model of Atom - Limitations
Explained emission spectra of
hydrogen very well
Failed to explain spectrum of other
elements
Did not fully account for chemical
behavior of atoms
Bohr Model of Atom
Has been shown that this model is
fundamentally incorrect – electrons not
particles in fixed orbits (classical model)
Schematic for X-ray Diffraction
X-ray beam with continuous range of
wavelengths incident on crystal
Diffracted radiation intense in certain directions
Intense spots
correspond to
constructive
interference from
waves reflected from
layers of crystal
Diffraction pattern
detected by
photographic film
Electrons as Waves - Evidence
Diffraction observed when electrons with
sufficient momentum strike an ordered crystal
lattice
Electrons
Nickel
Crystal
Detector
Screen
Diffraction
pattern on
detector screen
Comparing X-Ray and Electron
Diffraction Patterns in Al Foil
X-Ray
Electron
Electrons - particles with mass and charge create diffraction patterns in a manner similar to
electromagnetic waves!
Electrons as Waves
Louis De Broglie (1924)
Wavelength  associated with particle
of mass m moving at velocity v
 = h/ mv
de Broglie equation
 All moving particles have wave
characteristics
Idea for Particle as Wave
Vibrating guitar string
Only multiples of /2
allowed
Orbiting electron
Only multiples of
 allowed
Heisenberg Uncertainty Principle
1927
Fundamentally impossible to know
precisely both velocity and position of
particle at same time
Not just a technical limitation
Finding Out Where an Electron Is
Act of measuring changes properties
To determine electron location, use
light as probe
But light moves electron
And hitting the electron changes the
frequency of the light
Heisenberg Uncertainty Principle
Impact of photon on knowledge of
location and velocity of electron in an
atom
Before
Before collision
collision
After collision
Quantum Mechanical Model of Atom
1926 – Schrödinger wave equation
• Treated hydrogen atom’s electron as a
wave
• Unlike Bohr model, worked for other
atoms besides hydrogen
• Also limited electron’s energy to
quantized values
• Makes no attempt to describe electron’s
path around the nucleus
Bohr Model
According to Bohr’s
atomic model, electrons
move in definite orbits
around nucleus, much
like planets circle sun
These orbits, or energy
levels, are located at
certain fixed distances
from nucleus
Wave Model (Electron Cloud)
Quantum
mechanical (wave /
electron cloud)
model of atom
specifies only
probability of finding
electron in certain
regions of space
Marble
Model
Plum Pudding
Model
Nuclear Model
Planetary
(Bohr) Model
Quantum Mechanical
(wave / electron
cloud) Model
Classical to Quantum Theory
Indivisible Electron Nucleus Orbit
Greek
X
Dalton
X
Thomson
X
Rutherford
X
X
Bohr
X
X
Wave
X
X
Electron
Cloud
X
X
Summary
Major Observations & Theories Leading
from Classical to Quantum Theory
Observation
Scien
tist
Planck
1900
Theory
Spectral
Energy quantized
shape of
blackbody
radiation
Photoelectric Einstein Light has particle
effect
1905 behavior (photons)
Summary
Major Observations & Theories Leading
from Classical to Quantum Theory
Observation
Atomic line
spectra
Scien
tist
Bohr
1913
Bohr model de Broglie
works for H
1924
Theory
Energy of atoms
quantized; photon
emitted in orbit energy
transition
All moving particles
have wave-like nature
Summary
Major Observations & Theories Leading
from Classical to Quantum Theory
Observ
ation
?
?
Scien
Theory
tist
Schrodinger Quantum mechanical
description of atom using
1926
wave function
Heisenberg Fundamentally
impossible to know
1927
precisely both velocity &
position at same time
Wave Function (From Schrödinger
Wave Equation)
Y is wave function
(solution to wave
equation)
Square of Y gives
probability of finding
electron within particular
volume of space around
nucleus
Wave Function and Orbital
Y(wave function) defines atomic orbital –
3D description of electron’s probable
location
Entire family of wave functions exists, each
having particular set of quantum numbers
• “n” example of a quantum number
Quantum numbers determine electron
energy and shape/size of probability
distribution
Quantum Mechanical Model
Nucleus found inside
blurry “electron
cloud”
Orbital describes
chance of finding
electron in a region
Draw line/surface at
90% probability
• Shape may be
complex
Electron Density – Hydrogen Atom
A – likelihood of finding electron at
particular point  dot density
B – Orbital boundary: volume encloses 90%
probability of finding electron inside
A
B
Radial Probability Distribution
Most
probable
distance
Distance from nucleus (r)
Hydrogen Atom
Schrödinger wave equation can be
analytically solved for the H atom
Energy levels same as Bohr model –
also labeled by “n”
Position of electron no longer described
by circular orbit
Position specified by probability only –
details described by orbital
Electron Density – Hydrogen Atom
Some (small) probability electron will be
found a large distance from nucleus or
very close to nucleus
Hydrogen’s Atomic Orbitals
Quantum mechanical model assigns
principal quantum number (n)
• Indicates relative size and energy of
orbitals
• As n increases, electron spends more
time farther from nucleus
• n=1 is lowest = ground state
Hydrogen’s Atomic Orbitals
Principal energy levels contain energy
sublevels
Labeled as s, p, d, or f
• also g, etc but not concerned with in
this course
Sublevel value determines orbital
shape
Hydrogen’s Atomic Orbitals
n=1
Only one s sublevel with one s orbital
n=2
s and p sublevels; three p orbitals in p sublevel
n=3
s, p, and d sublevels; 5 d orbitals in d sublevel
n=4
s, p, d and f sublevels; 7 f orbitals in f sublevel
Hydrogen’s Atomic Orbitals
For hydrogen only (special case), all
sublevels of a given principal quantum
number n have the same energy
• For n=2, the 2s and 2p sublevels have
the same energy
• For n=3, the 3s, 3p, and 3d sublevels
have the same energy
Special case means energy levels have
same pattern as Bohr atom model
Hydrogen’s Atomic Orbitals
For a given sublevel, energy increases
with increasing n
Energy: 5p > 4p > 3p >2p (no 1p !)
Orbital Energies for Hydrogen Atom
(Aufbau diagram)
Orbital Notation
n sublevel direction (as subscript)
n = principal quantum number
sublevel = s, p, d, or f
direction = not applicable for s
x, y, z for p
xy, xz, yz, x2-y2, z2 for d
forget it for f
Orbital Notation
n sublevel direction (as subscript)
1s
3s
2px
3dxy
4pz
4dx2-y2
s and p Orbitals
(p shape exaggerated)
Node
Hydrogen
1s, 2s, 3s
Orbitals
1s
2s
3s
1s
2s
3s
1s
2s
3s
2pz
d Orbitals (also exaggerated)
The odd
one
(different
shape)
3dxz
3dz2
Hydrogen’s Atomic Orbitals
A given orbital 2s, 3px, 4dyx, 5s, etc
can be occupied by at most two
electrons
However, hydrogen has only one
electron to worry about
Summary of Sublevels
# of
orbitals
Max #
electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Summary by Energy Level
1st Energy Level
n=1
Only s orbital
Holds 2 electrons
1s2
2 total electrons =2n2
2d Energy Level
n=2
s and p orbitals are
available
2 in s, 6 in p
2s22p6
8 total electrons =2n2
By Energy Level
3d energy level
n=3
s, p, and d orbitals
2 in s, 6 in p, and 10
in d
3s23p63d10
18 total electrons
=2n2
4th energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d,
and 14 in f
4s24p64d104f14
32 total
electrons=2n2
Orbital Summary for Hydrogen
Chapter 5 - Electrons In Atoms
5.1 Light and Quantized Energy
5.2 Quantum Theory and the Atom
5.3 Electron Configurations
Section 5.3 Electron Configuration
A set of 3 rules determines the
arrangement of electrons in an atom. This
arrangement is called the electron
configuration.
• Apply the Pauli exclusion principle, the aufbau principle,
and Hund's rule to write electron configurations using orbital
diagrams and electron configuration notation (including
noble gas notation).
• Define valence electrons, and draw electron-dot structures
representing an atom's valence electrons.
Section 5.3 Electron Configuration
Key Concepts
• The arrangement of electrons in an atom is called the atom’s
electron configuration.
• Electron configurations are defined by the aufbau principle,
the Pauli exclusion principle, and Hund’s rule.
• An element’s valence electrons determine the chemical
properties of the element.
• Electron configurations can be represented using orbital
diagrams, electron configuration notation, and electron-dot
structures.
Rules for Filling Orbitals (1)
Aufbau principle – one by one build up
• Each electron occupies lowest energy
orbital available
• Energy level order determined from
diagram
Atomic Orbital Energies
For a given principal quantum number
n, all sublevels of a given type have the
same energy (said to be degenerate)
• All three p orbitals for n=2 are
degenerate
• All five d orbitals for n=3 are
degenerate
For hydrogen, all sublevels are
degenerate
• Levels are same as Bohr atom levels
Aufbau Diagram
for Hydrogen
Atom
Orbitals in Many-Electron Atoms
For n  2, the s- and p-orbitals are no
longer degenerate because the
electrons interact with each other
Unlike case for hydrogen, 3d > 3p >3s
Aufbau diagram looks different for
many-electron systems
• No longer follows simple Bohr model
Increasing energy
Aufbau Diagram – Multi-Electron Atoms
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Aufbau Diagram – Multi-Electron Atoms
Orbital
Energies:
Multi-Electron
Atoms
Aufbau Diagram Features – Table 5.3
Orbital Filling Order (see page 160)
Spin of the Electron
Associated with
electrons is a
property called spin
Spin generates a
magnetic field which
can be oriented up or
down
Use  or  to indicate
Rules for Filling Orbitals (2)
Pauli Exclusion Principle
• A maximum of two electrons may
occupy a single atomic orbital provided
the electrons have opposite spins
Rules for Filling Orbitals (3)
Hund’s Rule – how to handle
degenerate orbitals
Single electrons with same spin must
occupy each degenerate orbital before
additional electrons with opposite spins
share an orbital
• Rule arises because increased
electron-electron repulsion occurs when
2 electrons occupy the same orbital
Hund’s Rule for p Orbitals
Second electron
goes into empty
degenerate orbital
with spin in same
All
degenerate
direction as first
orbitals filled –
can start pairing
now
1
2
3
4
5
6
Electron Configurations
Periods 1, 2, and 3 (only)
Three rules:
• Electrons fill orbitals starting with lowest
n and moving upwards (Aufbau)
• No two electrons can fill one orbital with
the same spin (Pauli Exclusion
Principle)
• For degenerate orbitals, electrons fill
each orbital singly before any orbital
gets a second electron (Hund’s rule)
Constructing Orbital Diagrams
Use a box for each orbital
For 1 electron, box with single arrow
If 2 electrons in 1 orbital, use opposite arrows
Label each box with n and sublevel
C
1s
2s
2p
Electron Configuration Notation
Indicate n and sublevel for each orbital and
the total electron occupancy with a
superscript
C 1s22s22p2
• Distribution of electrons in the three p orbitals
(px, py, pz) not explicit in this notation
We write in n order, not energy order
• Ti 1s22s22p63s23p63d24s2
(note energy of 3d > 4s)
(Textbook uses energy order for configs.)
Orbital Diagram/ Electron Configuration
H
He 1s2
1s1
1s
1s
Li 1s22s1
Be 1s22s2
1s
1s
2s
222s
22s
222p
22p
513
24 6
B 1s
C
N
O
F
Ne
1s
1s
2s
2p
1s
2s
2p
2s
Electron Configuration
Determine the electron configuration for
Phosphorus (AN=15)
Need to account for 15 electrons
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2s
1s
First two electrons go
2p into the 1s orbital
Notice the opposite
spins
13 more to go
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2s
1s
Next two electrons go
2p
into the 2s orbital
11 more to go
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2s
1s
Next six electrons go
2p
into the 2p orbitals
5 more to go
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2s
1s
Next two electrons go
2p
into the 3s orbital
3 more to go
Increasing energy
7s
6s
5s
4s
7p
6p
5p
4p
2s
1s
5d
4d
5f
4f
3d
Last three electrons
go into the 3p orbitals
2p
Each go into separate
orbitals
3 unpaired electrons
1s22s22p63s23p3
3p
3s
6d
Using the Aufbau Diagram
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2
• 2 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2
• 4 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2
• 12 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2 3p6 4s2
• 20 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2 3p6 4s2
3d10 4p6 5s2
• 38 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2 3p6 4s2
3d10 4p6 5s2
4d10 5p6 6s2
• 56 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2 3p6 4s2
3d10 4p6 5s2
4d10 5p6 6s2
4f14 5d10 6p6 7s2
• 88 electrons
Fill from bottom up following
arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
1s2 2s2 2p6 3s2 3p6 4s2
3d10 4p6 5s2
4d10 5p6 6s2
4f14 5d10 6p6 7s2
5f14 6d10 7p6
• 108 electrons
• Maxed out
Noble Gas Notation
Na
1s22s22p63s1 (standard notation)
1s
Ne
2p
3s
1s22s22p6
1s
Na
2s
2s
2p
[Ne]3s1 (noble gas notation)
Noble Gas Configurations
Noble gases always have s and p
orbitals completely filled (He has no p)
2
6
ns np
Principal quantum number (n) of these
orbitals is same as the period in which
the gas is found - p orbitals are the last
filled (s for He)
He: 1s2
Ne: 2s22p6
Ar: 3s23p6
Kr: 4s24p6
Table 5.5 - Noble Gas Configurations
How to express noble gas using noble gas configuration
Determining Electron Configuration
Germanium (Ge), a semiconducting
element, is commonly used in the
manufacture of computer chips. What
is the ground state configuration for an
atom of germanium using noble gas
notation?
Atomic number of Ge = 32
Orbital Filling
Order (page 160)
[Ar]4s23d104p2
Above configuration
obtained from order
of filling. For this
class, use principal
energy level (n)
order
[Ar] 3d104s24p2
32
electrons
Practice
Electron Configurations
Practice problems, page 160
Probs 21(a-f), 22-25
Chapter assessment page 167
Probs 85(a-d), 86(a-d), 87(a-f), 88,
89
Supplemental Problems, page 978
Probs 5(a-d), 6
Orbital Filling Order
Lowest energy to higher energy.
Adding electrons can change energy of
orbital
Half filled sublevels have lower energy
Makes them more stable
Causes exceptions in the filling order
shown on the diagram
Exceptions to Filling Order
Aufbau diagram works to vanadium, AN 23
Half-filled and fully-filled set of d orbitals
have extra energy stability, so chromium is
Cr [Ar]3d54s1 (1/2 filled d)
Not [Ar]3d44s2
Next exception is copper:
Cu [Ar]3d104s1 (filled d)
Not [Ar]3d94s2
Valence Electrons
Defined as those electrons in the
atom’s outermost orbitals
• Orbitals with highest n
• If have a (n-1)d10 component in the
configuration, then ignore these electrons
for counting valence electrons, even if
higher energy than the ns2 electrons


Zn [Ar]3d104s2
2 valence electrons, not 12
Br [Ar] 3d104s24p5 7 valence electrons,
not 15 or 17
Electron-Dot Structures
Shorthand notation for indicating
valence electrons
• Write the element’s chemical symbol
• Add a dot for each valence electron


One at a time on all four sides of symbol
Then pair them up until all are used
Mg [Ne]3s2
S
[Ne]3s23p4
Mg
S
Valence Electrons
Determine the chemical (bonding)
properties of the element
S [Ne]3s23p4 6 valence electrons
Cs [Xe]6s1
1 valence electron
Dot Structures for Elements in 2d Period
Example Problem 5.3 - page 162
What is electron dot structure for tin?
Sn: [Kr]4d105s25p2 4 valence electrons
Sn
Practice
Electron-Dot Structures
Practice problems, page 162
Probs 26 (a-c), 27, 28
Section 5.3 Assessment, page 162
Prob 33
Chapter assessment pages 167-8
Probs 81(a-d), 90 (a-e), 91-93
Supplemental Problems, page 978
Probs 7, 8, 9 (a-d)
END