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Chapter 7
Atomic Structure and
Periodicity
7.1 Electromagnetic Radiation

electromagnetic radiation:
 form

of energy that acts as a wave as it
travels
 includes: visible light, X rays, ultraviolet
and infrared light, microwaves, and radio
waves
 travel at a speed of 2.9979 x 108 m/s in a
vacuum
All forms are combined to form
electromagnetic spectrum
Electromagnetic Spectrum
Low
Energy
High
Energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Visible Light Wavelength
- Page 139
“R O Y
Frequency Increases
Wavelength Longer
G
B I
V”
Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough
Electromagnetic radiation propagates through
space as a wave moving at the speed of light.
Equation:
c =
c = speed of light, a constant (2.998 x 108 m/s)
 (lambda) = wavelength, in meters
 (nu) = frequency, in units of hertz (hz or sec-1)
Wavelength and Frequency
 Are
inversely related
• As one goes up the other goes
down.
 Different frequencies of light are
different colors of light.
 There is a wide variety of frequencies
 The whole range is called a spectrum
The Wave-like Electron
The electron propagates
through space as an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
Electromagnic
radiation
propagates
through space
as a wave
moving at the
speed of light.
Wave nature of electromagnetic Radiation


wavelength:
λ
= Greek letter lambda
 distance between points on adjacent waves
(consicutive peaks or troughs)
 in nm (109nm = 1m)

frequency:
  = Greek letter nu
 number of wave cycles that passes a point
in a second. 108 cycles/s= 108 s-1
 =108 Hertz = 108 Hz
 in 1/second (Hertz = Hz)
c  

c

C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Calculate the energy of red light vs. blue light.
red light: 700 nm
blue light: 400 nm
red:
hc
E

hc
E
 blue:
hc
E

34
8
(6.62x1034 J  s)( 3.00x108 m / s)
(
6
.
62
x
10
J

s
)(
3
.
00
x
10
m / s)
E
E
700x109 m
400x109 m
E = 2.85 x 10-19 J
E = 4.96 x 10-19 J
sunburn????? uv
7.2 Nature of Matter

Before 1900, scientists thought that matter
and energy were totally different
matter
particles
mass
position
energy
wave
massless
delocalized
In 1900
Matter and energy were seen as
different from each other in
fundamental ways.
 Matter was particles.
 Energy could come in waves, with any
frequency.
 Max Planck found that as the cooling
of hot objects couldn’t be explained by
viewing energy as a wave.

Explanation of atomic spectra
 When
we write electron
configurations, we are writing the
lowest energy.
 The energy level, and where the
electron starts from, is called it’s
ground state - the lowest energy
level.
Changing the energy
 Let’s
look at a hydrogen atom, with
only one electron, and in the first
energy level.
Changing the energy

Heat, electricity, or light can move the
electron up to different energy levels.
The electron is now said to be “excited”
Changing the energy
As the electron falls back to the ground 
state, it gives the energy back as light
Changing the energy
They may fall down in specific steps 
Each step has a different energy 
Ultraviolet
 The
Visible
Infrared
further they fall, more energy is
released and the higher the
frequency.
 This is a simplified explanation!
 The orbitals also have different
energies inside energy levels
 All the electrons can move around.
What is light?
Light is a particle - it comes in chunks.
 Light is a wave - we can measure its
wavelength and it behaves as a wave
2
 If we combine E=mc , c=, E = 1/2
mv2 and E = h, then we can get:

 = h/mv
(from Louis de Broglie)
 called de Broglie’s equation
 Calculates the wavelength of a particle.
Wave-Particle Duality
J.J. Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The
electron is
a particle!
The electron
is an energy
wave!
Confused? You’ve Got Company!
“No familiar conceptions can be
woven around the electron;
something unknown is doing we
don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
The physics of the very small
 Quantum
mechanics explains
how very small particles behave
• Quantum mechanics is an
explanation for subatomic
particles and atoms as waves
 Classical mechanics describes
the motions of bodies much
larger than atoms
Heisenberg Uncertainty
Principle
 It
is impossible to know exactly the
location and velocity of a particle.
 The better we know one, the less
we know the other.
 Measuring changes the properties.
 True in quantum mechanics, but
not classical mechanics
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an
electron.”
Werner Heisenberg
You can find out where the
electron is, but not where it is
going.
OR…
You can find out where the
electron is going, but not where
it is!
It is more obvious with the
very small objects
 To measure where a electron
is, we use light.
 But the light energy moves the
electron
 And hitting the electron
changes the frequency of the
light.
After
Before
Photon
Moving
Electron
Photon
wavelength
changes
Electron
velocity changes
Fig. 5.16, p. 145
Light is a Particle?
 Energy
is quantized.
 Light is a form of energy.
 Therefore, light must be quantized
 These smallest pieces of light are
called photons.
 Photoelectric effect? Albert Einstein
 Energy & frequency: directly related.
The energy (E ) of electromagnetic
radiation is directly proportional to the
frequency () of the radiation.
Equation:
E = h
E = Energy, in units of Joules (kg·m2/s2)
(Joule is the metric unit of energy)
h = Planck’s constant (6.626 x 10-34 J·s)
 = frequency, in units of hertz (hz, sec-1)
Nature of Matter

Max Planck: a German physicist
suggested that an object emits energy in the
form of small packets of energy called
quanta

Quantum- the minimum amount of energy
that can be gained or lost by an atom
E  h
Planck’s constant (h): 6.626 x 10-34 J*s
Nature of Matter

Einstein proposed that radiation itself is
really a stream of particles called photons

Energy of each photon is :
E photon  hv 

also showed that energy
has mass
E  mc
2
hc

Nature of Matter
c
v
c  v

E  hv
h

mc
hc

 mc
E
hc

2
E  mc
shows that anything with both mass and velocity
has a corresponding wavelength
2
Nature of Matter
In 1924, Louis de
Broglie
(French scientist)
 suggested that matter
has both particle-like
and wave-like
characteristics

h

mv
Main Ideas:




matter and energy are not distinct
energy is a form of matter
larger objects are mostly particle-like
smaller objects are mostly wave-like
7.3 The Atomic Spectrum of Hydrogen
Spectroscopic analysis of the visible
spectrum…
White light
…produces all of the colors in a continuous spectrum
Atomic Spectra
 White
light is
made up of all
the colors of the
visible
spectrum.
 Passing it
through a prism
separates it.
• These are called
the atomic
emission
spectrum
• Unique to each
element, like
fingerprints!
• Very useful for
identifying
elements
Continuous Spectra
White light
passed
through a
prism
produces a
spectrum –
colors in
continuous
form.
The Continuous Spectrum
 ~ 650 nm
 ~ 575 nm
 ~ 500 nm
 ~ 480 nm
 ~ 450 nm
The different
colors of light
correspond to
different
wavelengths
and
frequencies
Continuous Emission Spectrum



line-emission spectrum- series of
wavelengths of light created when visible
portion of light from excited atoms is
shined through a prism
scientists using classical theory expected
atoms to be excited by whatever energy they
absorbed
continuous spectrum emission of continuous range of
frequencies of EM radiation
 contains all wavelengths of visible light
Spectroscopic analysis of the
hydrogen spectrum…
H receives a high energy
spark
H-H bonds
Are broken and H atoms
are excited
…produces a “bright line” spectrum
Line Spectra
Light passed
through a
prism from
an element
produces a
discontinuous
spectrum of
specific
colors
Hydrogen
only four lines are observed
Line Spectra
The pattern of lines emitted by excited atoms of
an element is unique
= atomic emission spectrum
• These are called
the atomic
emission
spectrum
• Unique to each
element, like
fingerprints!
• Very useful for
identifying
elements
H Line-Emission Spectrum
 light
is emitted by excited H atoms
when bond is broken in the
diatomic molecule
 ground state- lowest energy state
of an atom
 excited state- when an atom has
higher potential energy than it has
at ground state
H Line-Emission Spectrum
 when
an excited electron falls back
to ground state, it emits photon of
radiation
 the photon is equal to the
difference in energy of the original
and final states of electron
 since only certain frequencies are
emitted, only certain energies are
allowed for electrons in H atom
Electron
transitions
involve jumps
of
definite
amounts
of bands
This produces
of light with definite
energy.
wavelengths.
7.4 The Bohr Model



Niels Bohr (Danish physicist) in 1913
Developed a quantum model
for H atom that explained the emission
line spectrum
Electron moves around the nucleus
only in certain allowed circular orbits,
in which it has a certain amount of
energy
The Bohr model
Energy level of an electron analogous
to the steps of a ladder
 The electron cannot exist between
energy levels, just like you can’t stand
between steps on a ladder
 A quantum of energy is the amount of
energy required to move an electron
from one energy level to another

Niels Bohr




Developed the quantum model of the
hydrogen atom.
He said the atom was like a solar
system.
The electrons were attracted to the
nucleus because of opposite charges.
Didn’t fall in to the nucleus because it
was moving around.
The Bohr Atom





He didn’t know why but only certain
energies were allowed.
He called these allowed energies
energy levels.
Putting Energy into the atom moved the
electron away from the nucleus.
From ground state to excited state.
When it returns to ground state it gives
off light of a certain energy.
The Model: Summary



Space around nucleus is divided into
spherical (circualr) paths (orbits) each has a
number called “Principal Quantum
number”
The electron can exist only in one of these
orbitals but not in between
Orbits possess fixed size and energy,
therefore electron has a definite energy
characteristic of its orbit
R Z
E E 
 Energy of electron
n
J
R  Rydberg constant  2.180x10
particle
2
H
n
orbit
2
-18
H
En  2.180 X 10
23
18
J
6.02 X 10 particles
1kJ
X
X
 1312kJ / mol
particle
1mol
1000J
 1312kJ
En 
n mole
2


Orbits allowed for electron are those in
which electron has an angular momentum= nh
1


An electron can pass only from one bit to
another. Absorption or emission will occur
Energy of the outermost orbit is zero
The Bohr Atom
n=4
n=3
n=2
n=1
Bohr Model


To create an accurate model, he had to use
quantum theory instead of classical
created an equation used to calculate the energy
levels available to electrons in a certain atom:
E  2.178 10
18
2
Z
J( 2 )
n
where n= integer and Z=atomic number
negative sign makes the energy more negative the
closer it is to the nucleus
Bohr Model


Can gain energy by
moving to a higher
energy level
Can lose energy by
moving to lower
energy level
Bohr Model
a photon is
released that has
an energy equal to
the difference
between the initial
and final energy
orbits
Bohr Model

equation can be used twice to find the ∆E
when an electron moves energy levels
E  2.178 10
E  [2.178 10
18
J(
Zf
2
nf
2
E  2.178 10
18
2
Z
J( 2 )
n
)]  [2.178 10
18
1
1
J( 2  2 )
nf
ni
18
2
Zi
J ( 2 )]
ni
Bohr Model

can wavelength of photon released by using
hc

E

E=0 is set at an distance of ∞ away from the
nucleus and becomes more negative as the
electron comes closer to the nucleus
E  2.178 10
18
1
J( )  0

Example 1
1
1
Calculate the
18
E  2.178 10 J ( 2  2 )
energy required to
nf
ni
move the hydrogen
electron from n=1
1 1
18
E  2.178 10 J ( 2  2 )  1.633 10 18 J
to n=2. Find the
2 1
wavelength of
radiation that had to
m
34
(6.626 10 J  s)( 2.9979 )
hc
s
be absorbed by the


E
1.633 10 18 J
electron.
9
10
nm
  1.216 107 m 
 121.6nm
1m
Calculate the
energy required
to remove the
electron from the
hydrogen atom
in its ground
state.
Example 2
E  2.178 10
18
E  2.178 10
E  2.178 10
18
1
1
J( 2  2 )
nf
ni
18
1 1
J(  2 )
 1
J (0  1)  2.178 10
18
Energy was absorbed by the electron so the
value of ∆E value is positive.
J
The Bohr Model




Doesn’t work.
Only works for hydrogen atoms.
Electrons don’t move in circles.
The quantization of energy is right, but
not because they are circling like
planets.
Bohr Model

problems:
 did not work for other atoms
 did not explain chemical
behavior of atoms
Heisenberg’s Uncertainty Principle





According to de Broglie: Electron behaves like a
wave
It is possible to specify the position of a wave at a
particular instant?
Energy, wavelength and amplitude can be
determined
But exact position is impossible to be determined
The electron cannot be imagined as :




moving particle
In a path of the same radius (well defined orbits)
Thus, location, direction and speed of motion of a
particle cannot be determined
Then Bohr Model had to be “Abandoned
Heissenberg Uncertainty Principle
“It is impossible to determine both the
position and momentum of a subatomic
particle (such as the electron) with
arbitrarily high accuracy”
 The
effect of this principle is to convert the
laws of physics into statements about relative,
instead of absolute, certainties.
Heisenberg Uncertainty Principle



we cannot know the
exact position and
momentum (motion)
of the electron
as more is known
about position, less is
known about
momentum
uncertainties are
inversely proportional
h
 x   ( m ) 
4
where
∆x: uncertainty in
position
∆m : uncertainty in
mometum
minimum uncertainty
is h/4
7.5 The Quantum Mechanical Model

Exactt position of electron can not be
defined



Exact bath of electron about nucleus can not be
defined
Werner Heisenberg, Louis de Broglie and
Erwin Schrodinger made the approach
called “Quantum Mechanics”
They assumed that the electron is a standing
wave
The Quantum Mechanical Model
Waves are associated with electrons
 Information about energies of
electrons and their positions are
obtained from studying the associated
waves
 Description of electron is based upon
“ Probability of finding a particle
within a given region of space” “ but
not on the exact position”

Schrödinger Equation





Wave equation describing electron as being
a wave
The amplitudes (height), , of electron
wave at various points of space are
calculated
 commonly called “wave function”
 provides information about the allowable
energies for an electron in H atom.
 corresponds to a certain energy and
describes a region around nucleus “Orbital”
where the electron having that energy may
be found
Orbital: Region around the nucleus
where the electron can be expected to
be found
 The Function 2
2 describes the probability of the
position of the electron at a particular
point
 2  Probablity of finding a particle in
a given region of space
 2  Electric charge density at a given
region of space

Orbital
 A 3 dimensional space around a nucleus in
which electrons are most likely to be found
 Shape represents electron density (not a
path the electron follows)
 Each orbital can hold up to 2 electrons.
Thus,



The charge can be assumed to be spread out
as a charge cloud by rapid motion of
electron
The cloud is denser in some regions than
others
The probability of finding electron in a
given region in space is proportional to the
density of the cloud
Meaning of Wave Function



the wave function itself does not have
concrete meaning
the square of the wave function represents
the probability of finding an electron at a
certain point
easily represented as probability distribution
where the deepness of color indicates the
probability
Meaning of Wave Function



(a) electron density map
probability of finding an
electron is highest at short
distances from nucleus
(b) calculated probability of
finding an electron at
certain distances from
nucleus in the 1s orbital
7.6 Quantum Numbers




There are many solutions to Schroedinger’s
equation for H atom
Each solution is a wave function called
Orbital.
Each solution can be described with
quantum numbers that describe some
aspect of the solution.
Schrödinger’s equation requires 3
quantum numbers
7.6 Quantum Numbers


Quantum numbers specify the properties of
atomic orbitals and of electrons in orbitals
the first three numbers come from the
Schrödinger equation and describe:
main energy level
 shape
 orientation


4th describes state of electron
1st Quantum Number
Principal Quantum Number: n
 Main energy level (or shell) occupied by
electron. They are called atomic orbitals



regions where there is a high probability of
finding an electron.
values are all positive integers >0 (1,2,3,…)
As n increases
 size of orbital is larger
 electron has higher energy
 the electron’s average distance from
the nucleus increases
Principal Quantum Number
Maximum number
of electrons that
can fit in an
energy level:
2n2
st
1
Energy
Quantum Number
nd
2
Quantum Number
Angular Momentum Quantum Number: l
 indicates the shape of the orbital
(sublevel or subshell)
 the number of possible shapes (or l
values) for an energy level is equal to n
 the possible values of l are 0 and all
positive integers less than or equal to n
-1
 l has integer values from 0 to n-1
l
= 0 is called s
 l = 1 is called p
 l =2 is called d
 l =3 is called f
 l =4 is called g
2nd Quantum Number
s orbitals: 1: s
 spherical
 l value of 0
 1st occur at n=1
2nd Quantum
Number
p orbitals: 3
2px, 2py, 2pz
 dumbbellshaped
 l value of 1
 1st occur at
n=2
 for n>2,
shape is
same but
size
increases
nd
2
Quantum Number
d orbitals: 5: 3dxz, 3dyz, 3dxy, 3dx2-y2, dz2
 mostly cloverleaf
 l value of 2
 1st occur at n=3
 for n>3, same shape but larger size
nd
2
Quantum Number
f orbitals: 7 types
various shapes 
l value of 3 
begin in n=4 
nd
2




Quantum Number
Other shapes can exist in energy levels as
long as they follow the rules
g (l=4) starts in 5 with 9 orbitals
h (l=5) starts in 6 with 11 orbitals, etc
but no known elements have electrons in
them at ground state
nd
2
Level
Quantum Number
Sublevels
Sublevels
0
1
2
0
1
2
0
1
0
3
3rd Quantum Number
Magnetic Quantum Number: ml
 indicates the orientation of an orbital around the
nucleus

has values from +l
 -l

specifies the exact orbital that the electron is
contained in



each orbital holds maximum of 2 electrons
total number of orbitals is equal to n2 for an energy
level
number of possible ml values for a certain
subshell is equal to 2l + 1
rd
3
Quantum Number
Energy
Level
(n)
1
2
3
4
Sublevels in
Level
# Orbitals
in Sublevel
Total # of
Orbitals in
Level
s
s
1
1
1
4
p
3
s
1
p
3
d
5
s
1
p
3
d
f
5
7
9
16
th
4
Quantum Number
Spin Quantum Number: ms
 indicates the spin state of the electron
 only 2 possible directions
 only 2 possible values: +½ and -½
 paired electrons must
have opposite spins
 maximum number of
electrons in an energy
level is 2n2
Quantum Numbers:
type
n = principle quantum no.
values
1,2,3,etc
meaning
shell (period)
l = angular momentum quant. #
or azimuthal q#
0,1,2,3...
s,p,d,f,g,h..
subshell
ml = magnetic q. #
0,1,2,3,..
orbital
ms = spin q. #
½
spin
These 4 Quantum numbers give the general
location of electrons within an atom and the
general shape of the orbital in which they
reside.
Electrons Allowed
 All electrons in the same sublevel have the same
energy.
 All 2s electrons have the same energy. All 2p
electrons have the same energy which is slightly
higher than the energy of the 2s electrons
s sublevel
2 electrons
p sublevel
6 electrons
d sublevel
10 electrons
f sublevel
14 electrons
7.9 Polyelectronic Atoms



Kinetic energy - as the electrons move
around the nucleus
Potential energy - from their
attraction to nucleus
Potential energy - from their repulsion
to each other
Electron Correlation Problem



can’t find the exact location of electrons
can’t find the specific repulsions between
electrons
so we must treat each electron as if it has an
average amount of attraction to nucleus and
repulsion to other electrons
Electron Shielding


occurs when an electron is not attracted to the
nucleus
because of electrons in lower energy levels
repelling it.
Penetration Effect

all orbitals in the same energy level do NOT
have the same amount of energy ( are not
degenerate)
Es < E p < Ed < E f

the amount of energy in each sublevel is
determined by its average distance from the
nucleus
n=4
4 subshells
f
f
f
f
f
f
f
f (7 orbitals)
d (5 orbitals)
d
p (3 orbitals)
d
px
d
d
py
s (1 orbital)
d
pz
n = 4, l = 0, ml = 0
s
n=3
3 subshells
d (5 orbitals)
d
p (3orbitals)
d
px
d
py
n = 4, l = 2, ml = +2
d
d
pz
+
s (1 orbital)
n=2
2 subshells
p (3 orbitals)
s (1 orbital)
n=1 (1st shell)
1 subshell
s (1 orbital)
s
px
py
pz
n = 2, l = 1, ml = -1
s subshell
s subshell
only holds 2 electrons
Gives: "electron Address" = (n)energy , (l)shape of house (orbital),
(ml) which , spin
n=4
4 subshells
f
f
f
f
f
f
f
f (7 orbitals)
d (5 orbitals)
d
p (3 orbitals)
d
px
d
d
py
s (1 orbital)
d
pz
s
n=3
3 subshells
d (5 orbitals)
d
p (3orbitals)
d
px
d
py
d
d
pz
+
s (1 orbital)
n=2
2 subshells
p (3 orbitals)
s (1 orbital)
n=1 (1st shell)
1 subshell
s (1 orbital)
l = 3 ml =-3,-2,-1,0,1,2,3
n = 4 l = 2 ml =-2,-1,0,1,2
l = 1 ml =-1,0,1
l=0
ml =0
l=2
n=3 l=1
l=0
s
px
py
pz
n=2
s subshell
s subshell
only holds 2 electrons
Gives: "electron Address" = (n)energy , (l)shape of house (orbital),
(ml) which , spin
n =1
n=4
4 subshells
f
f
f
f
f
f
f
f (7 orbitals)
d (5 orbitals)
d
p (3 orbitals)
d
px
d
d
py
s (1 orbital)
d
Pauli exclusion principle:
no 2 e-’s in the same atom
can have the same 4 quantum
numbers.
pz
s
n=3
3 subshells
d (5 orbitals)
d
p (3orbitals)
d
px
d
py
d
d
pz
+
s (1 orbital)
n=2
2 subshells
p (3 orbitals)
s (1 orbital)
n=1 (1st shell)
1 subshell
s (1 orbital)
s
px
py
Each compartment (orbital)
can only hold two electrons.
pz
s subshell
s subshell
only holds 2 electrons
Gives: "electron Address" = (n)energy , (l)shape of house (orbital),
(ml) which , spin
Hund’s rule: each orbital
of a subshell will get 1 e- of
parallel spin before they
are paired.
n = 2, l =1, ml =+1, ms = -1/2
7.11 Electron Arrangement in Atoms
The way electrons are arranged in
various orbitals around the nuclei of
atoms. Three rules tell us how:
1) Aufbau principle - electrons enter the
lowest energy first.

•
This causes difficulties because of
the overlap of orbitals of different
energies – follow the diagram!
2) Pauli Exclusion Principle - at most 2
electrons per orbital - different spins
Electron Configurations
Electron configuration describes the distribution of
electrons among the various orbitals in the atom
The spdf notation uses
numbers to designate a
principal shell and the
letters to identify a subshell;
a superscript number
indicates the number of
electrons in a designated
subshell
Orbital Diagram
An orbital diagram uses boxes to represent
orbitals within subshells and arrows to represent
electrons:
Each box has arrows representing electron spins;
opposing spins are paired together
EOS
Rules for Electron Configurations
Electrons occupy the lowest available energy
orbitals
Pauli exclusion principle – no two electrons in the
same atom may have the same four quantum
numbers
•Orbitals hold a maximum of two electrons
•spins must be opposed
Rules for Electron Configurations
For orbitals of identical energy, electrons enter
empty orbitals whenever possible – Hund’s rule
Electrons in half-filled orbitals have parallel spins
EOS
Rules for Electron Configurations
Capacities of shells (n) and subshells (l)
EOS
Rules for Electron Configurations
Subshell filling order ...
Each subshell must
be filled before
moving to the next
level
Rules for Electron Configurations
Subshell filling order ...
Each subshell must be
filled before moving
to the next level
Aufbau Principle
Aufbau is German for building up.
 As the protons are added one by one,
the electrons fill up hydrogen-like
orbitals.
 Fill up in order of energy levels.

The Aufbau Principle
A hypothetical building up of an atom from the one that
precedes it in atomic number
(Z = 1) H 1s1
(Z = 3) Li 1s22s1  [He]2s1
(Z = 2) He 1s2
Abbreviated electron configuration
(Z = 3) Li 1s22s1
EOS
The Aufbau Principle ...
[He]2p2
[He]2p3
[He]2p4
[He]2p5
[He]2p6
EOS
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
3d
3p
3s
2p
2s
1s
aufbau diagram
5f
4f
Pauli Exclusion Principle
No two electrons in an
atom can have the same
four quantum numbers.
Wolfgang Pauli
To show the different
direction of spin, a pair
in the same orbital is
written as:
Quantum Numbers
Each electron in an atom has a
unique set of 4 quantum numbers
which describe it.
1)
2)
3)
4)
Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number
Electron Configurations
3) Hund’s Rule- When electrons

occupy orbitals of equal energy,
they don’t pair up until they have
to.
Let’s write the electron
configuration for Phosphorus
 We need to account for all 15
electrons in phosphorus
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2s
1s
4f
3d
4s
3p
5f
The first two electrons
go into the 1s orbital
2p
Notice the opposite
direction of the spins
 only 13 more to go...

Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
5f
4f
3d
4s
3p
3s
2p
2s
1s
The next electrons
go into the 2s orbital
 only 11 more...

Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next electrons
go into the 2p orbital
• only 5 more...
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next electrons
go into the 3s orbital
• only 3 more...
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The last three electrons
go into the 3p orbitals.
2p
They each go into
separate shapes (Hund’s)
• 3 unpaired electrons
• = 1s22s22p63s23p3
Orbital Diagram for A Nitrogen Atom
7N

1s

2s


2p

3s
Orbital Diagram for A Fluorine Atom
9F

1s

2s

2p


3s
Orbital Diagram for A Magnesium Atom
12Mg

1s

2s

2p



3s
8O


1s
2s


2p

3s
Write the orbital diagram for the electrons
in an iron atom 26Fe
 
1s
  
2s
2p
   
3d

3s

  
3p
Orbitals fill in an order
 Lowest
energy to higher energy.
 Adding electrons can change the
energy of the orbital. Full orbitals
are the absolute best situation.
 However, half filled orbitals have a
lower energy, and are next best
• Makes them more stable.
• Changes the filling order
Write the electron configurations
for these elements:
Titanium - 22 electrons
 1s22s22p63s23p64s23d2
 Vanadium - 23 electrons
 1s22s22p63s23p64s23d3
 Chromium - 24 electrons
 1s22s22p63s23p64s23d4 (expected)
But this is not what happens!!

Chromium is actually:
 1s22s22p63s23p64s13d5
 Why?
 This
gives us two half filled
orbitals (the others are all still full)
 Half full is slightly lower in energy.
 The same principal applies to
copper.
Copper’s electron configuration
 Copper
has 29 electrons so we
expect: 1s22s22p63s23p64s23d9
 But the actual configuration is:
 1s22s22p63s23p64s13d10
 This change gives one more filled
orbital and one that is half filled.
 Remember
9
d
these exceptions:
4
d,
Electron Configurations of Ions
Anions: gain e– to complete the valence shell
Example:
-
EOS
Electron Configurations of Ions
Cations: lose e– to attain a complete valence
shell
Example:
(Z = 11) Na
(Z = 11) Na+
EOS
Electron Configurations of Ions
Cations formed from transition metals lose e–
from the highest principal energy level (n)
EOS
7.10 The history of the Periodic Table




Developed independently by German
Julius Lothar Meyer and Russian Dmitri
Mendeleev (1870”s).
Didn’t know much about atom.
Put in columns by similar properties.
Predicted properties of missing
elements.
The aufbau principle and
Periodic Table
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
• 38
20electrons
4212
Details of the Periodic Table
Elements in the same column have
the same electron configuration.
 Elements in columns have similar
properties.
 Similar properties because of
electron configuration.
 Noble gases have filled energy
levels.
 Transition metals are filling the d

Valence electrons- the electrons in
the outermost energy levels (not
d).
 Core electrons- the inner
electrons.
 Hund’s Rule- The lowest energy
configuration for an atom is the
one have the maximum number of
of unpaired electrons in the orbital.


C 1s2 2s2 2p2
Valence Electrons and Core Electrons
Valence electrons are those with the highest
principal quantum number
Sulfur has six valence electrons
EOS
Valence Electrons and Core Electrons
Electrons in inner shells are called core electrons
Sulfur has 10 core electrons
EOS
Periodic Relationships
We can deduce the general form of electron
configurations directly from the periodic table
The Periodic Table
Exceptions
2
2
 Ti = [Ar] 4s 3d
2
3
 V = [Ar] 4s 3d
1
5
 Cr = [Ar] 4s 3d
2
5
 Mn = [Ar] 4s 3d
Half filled orbitals.
 Scientists aren’t sure of why it
happens
1
10
 same for Cu [Ar] 4s 3d

Irregular configurations of Cr and Cu
Chromium steals a 4s electron to
make its 3d sublevel HALF FULL
Copper steals a 4s electron
to FILL its 3d sublevel
More exceptions
2
1
 Lanthanum La: [Xe] 6s 5d
2 1
1
 Cerium Ce: [Xe] 6s 4f 5d
Promethium Pr: [Xe] 6s2 4f3 5d0
2 7
1
 Gadolinium Gd: [Xe] 6s 4f 5d
2
14
1
 Lutetium Pr: [Xe] 6s 4f
5d
 We’ll just pretend that all except Cu
and Cr follow the rules.

Main Group (Representative) and
Transition Elements
Elements in which the orbitals being filled in the
aufbau process are either s or p orbitals of the
outermost shell are called main group
(Representative) elements
“A” group designation on the periodic
table
The first 20 elements are all main group
elements
Transition Elements
In transition elements, the sublevel (shell) being
filled in the aufbau process is in an inner
principal shell (d or f)
d-Block transition elements: Electrons enter the
d-sublevels.
f-Block transition elements: d sublevels are
completely filled. Electrons enter f-sublevels
Lanthanides: electrons fill 4f subleve
Actinides: electrons fill 5f sublevels
Periodic Relationships
General form of electron configurations can be deduced
directly from the periodic table
The Periodic Table
Periodic Relationships
Transition Elements
Completely
filled and halffilled sublevels
are more
energetically
favorable
configurations
7.12 Periodic Trends in
Atomic Properties
Ionization Energy

Ionization energy the energy required to remove
an electron form a ground state atom in the
gaseous phase
A (g) + energy  A+ (g) + e Highest energy electron removed first.
 First ionization energy (I1) is that required
to remove the first electron.
 Second ionization energy (I2) - the second
electron, etc.
Successive ionization Energies
Continual removal of electrons increases ionization
energy greatly
B  B+ + e–
I = 801 kJ mol–1
B+  B+2 + e–
I = 2427 kJ mol–1
B+2  B+3 + e–
I = 3660 kJ mol–1
B+3  B+4 + e–
I = 25,025 kJ mol–1
B+4  B+5 + e–
I = 32,822 kJ mol–1
Ionization Energy:
The energy required to completely remove an
e- from an atom in its gaseous state.
Mg(g)  Mg1+ + eMg1+(g)  Mg2+ + e-
1st ionization energy
2nd ionization energy
Question: Which of the above ionizations would have the
highest ionization energy and why?
Trends in ionization energy

for Mg
•
•
•


I1 = 735 kJ/mole
I2 = 1445 kJ/mole
I3 = 7730 kJ/mole
The effective nuclear charge increases as
electrons are removed
It takes much more energy to remove a core
electron than a valence electron because
there is less shielding.
Ionization Energy


if one electron is removed, the positive
charge binds the electrons more tightly so
2nd ionization energy must be higher
the largest jump in energy is when you
remove a core electron instead of valence
Ionization Energy
First ionization energies across Periods and Groups
Across Periods and Groups



Generally from left to right, I1 increases
because
there is a greater nuclear charge with the
same shielding.
As you go down a group I1 decreases
because electrons are farther away from
nucleus
Ionization Energy

Across Period:
 requires more
energy to remove an
electron so increases
 because electrons
added in the same
energy level do not
shield electrons
from nuclear charge

Down Group:
 requires less energy
to remove electron
so decreases
 because the valence
electrons are farther
away from protons
attracting them
It is not that simple



Zeff changes as you go across a period, so
will I1
Half filled and filled orbitals are harder to
remove electrons from.
here’s what it looks like.
Atomic number
First Ionization energy
Atomic number
First Ionization energy
Atomic number
First Ionization energy
Ionization Energy
Ionization Energy
Ionization Energy
Electron Affinity
Electron Affinity – the energy change when an
electron is added to a gaseous neutral
atom
A + e-  A- + energy
Cl(g) + e-  Cl-(g) E = -349 kJ/mol
• Electron affinities are expressed as negative
because the process is exothermic
Electron affinity values
Electron Affinity
Electron Affinity

Across Period:
releases more
energy so number
increases (gets more
negative)
 because electrons
added in the same
energy level do not
shield electrons
from nuclear charge


Down Group:
 releases less energy
so number
decreases (gets less
negative)
 because the
electrons being
added are farther
away from the
attracting protons
Electron Affinity
Atomic Radii


Size of orbitals can not be
specified exactly, neither
can the size of atom
Atomic Radii – half the
distance between the
nuclei of identical atoms
that are bonded together
Atomic Radii

Across Period:
 atoms get smaller
 because of the
increased number of
protons attracting the
electrons
 the electrons added in
the same energy level
do not shield electrons
from nuclear charge

Down Group:
 atoms
get larger
 increases
 because the energy
levels being added
to the atom
Atomic Radii Properties
Ionic Radii
The ionic radius of each ion is
the portion of the distance
between the nuclei occupied
by that ion
Ionic Radii
Cations are smaller than the
atoms from which they are
formed
– the nucleus attracts the
remaining electrons more
strongly
Anions are larger than the atoms
from which they are formed
– the greater number of
electrons repel more strongly
EOS
Isoelectronic Configurations
• Elements that all have the same number of electrons
For isoelectronic species, the greater the
nuclear charge, the smaller the species
Effective nuclear charge
Atomic and Ionic Radii
Atomic/Ionic Radii
Summary of Periodic Trends
7.13 Properties of a group: Alkali metals
Information contained in the Periodic Table







I
Groups
n of representative elements exhibit similar
chemical
f properties that change in a regular way
Each o
group member has same valence electron
configuration
r
It is the number and type of valence electrons that
m an atoms chemistry.
determine
a get the electron configuration from the
You can
Periodic
t Table
Metals
i lose electrons have the lowest IE
Non metalsgain electrons most negative electron
o
affinities.
n
Various groups have special names
Special Names for Groups
in the Periodic Table
Metals, Nonmetals, and Metalloids
Metals have a small number of electrons in their
valence shells and tend to form positive ions
Except for hydrogen
and helium, all s-block
elements are metals
All d- and f-block
elements are metals
EOS
Metals, Nonmetals, and Metalloids
Atoms of a nonmetal generally have larger numbers
of electrons in their valence shell than do metals,
and many tend to form negative ions
Nonmetals are all pblock elements and
include hydrogen and
helium
Metalloids have
properties of both
metals and nonmetals
EOS
The Alkali Metals



Do not include hydrogen- it behaves as a nonmetal
On going down the group there is :
 decrease in IE
 increase in radius
 Increase in density
 decrease in melting point
They loose their valence electron readily. Thus
they behave as reducing agents
 They react readily with nonmetals
Reducing ability






Lower IE< better reducing agents
Cs>Rb>K>Na>Li
works for solids, but not in aqueous
solutions.
In solution Li>K>Na
Why?
It’s the water -there is an energy change
associated with dissolving
Potassium
Reacts
Violently
with Water
Hydration Energy






It is exothermic
for Li+ -510 kJ/mol
for Na+ -402 kJ/mol
for K+ -314 kJ/mol
Li is so big because of it has a high charge
density, a lot of charge on a small atom.
Li loses its electron more easily because of
this in aqueous solutions
The reaction with water





Na and K react explosively with water
Li doesn’t !!!
Even though the reaction of Li has a more
negative H than that of Na and K
Na and K melt easily compared to Li. Thus
they will have greater contact with water
than Li
H does not tell you speed of reaction