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Chapter 9
Electronic Structure
and Periodic Trends
Homework
 Assigned Problems (odd numbers only)
 “Questions and Problems” 9.1 to
9.71 (begins on page 258)
 “Additional Questions and Problems”
9.81 to 9.115 (page 284-286)
 “Challenge Questions” 9.119, 9.121,
9.123 (page 286)
Electromagnetic Radiation
 Matter is anything that has mass and occupies
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space.
Nearly all changes that matter undergoes
requires the absorption or release of energy
Energy is the capacity to do work
 The process of moving matter against an
opposing force.
Forms of energy include heat, electrical, and light
One way energy is transmitted through space is
by Electromagnetic Radiation
 Transmits from one place to another in the form
of a wave
 Given off by atoms when they have been
excited by any form of energy
Electromagnetic Radiation
 Light (radiant) energy, which is visible and
invisible
 Classified into types according to the
frequency of the wave
 Sunlight, visible light, radio waves,
microwaves (ovens), X-rays, and heat from
a fire (infrared), are all forms of this radiant
energy
 These forms of radiant energy exhibit the
same wavelike characteristics
Wavelength and Frequency
 Electromagnetic radiation is radiant (light) energy that travels in
waves at the speed of light
 The waves have three basic characteristics: wavelength,
frequency, and speed
 The highest point on the wave is a peak
 Wavelength (l = distance between neighboring peaks)
 generally measured in nanometers (1 nm = 10-9 m)
 Velocity (v = how fast the wave is moving)
 c = speed of light
 3.00 x 108 m/s
 Amplitude (how tall the waves are)
 Frequency (u = the number of waves that pass a point in a
given time)
 generally measured in Hertz (Hz),
 1 Hz = 1 wave/sec = 1 sec-1
c = u x l
Waves
frequency
wavelength
frequency
wavelength
C = speed of light
Electromagnetic Spectrum
 Classified by wavelength:
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 Lower energy (longer wavelength, lower frequency)
 Higher energy (shorter wavelength, higher frequency)
Radiowaves: AM/FM/TV signals, cell phones, low frequency
and energy
Microwaves: Microwave ovens and radar
Infrared (IR): Heat from sunlight, infrared lamps for heating
Visible: The only EM radiation detected by the human eye
 ROYGBIV
Ultraviolet: Shorter in wavelength than visible violet light,
sunlight
X-rays: Higher in energy than UV
Gamma rays: Highest in energy, harmful to cells
Wavelengths of EM Radiation
Atomic Spectra and Energy Levels
 When white light passes through a prism
it produces a continuous rainbow of
colors from (red to violet)
 From red to violet the wavelength
becomes shorter
Atomic Spectra and Energy Levels
 When an element is
heated (strontium
and barium) light is
produced
 If this light is passed
through a prism, it
does not produce a
continuous rainbow,
only certain colors
Atomic Spectra and Energy Levels
 Only specific colors are produced in the
visible region. This is called a “bright-line
spectrum”
 Each line produced is a specific color,
and thus has a specific energy
 Each element produces a unique set of
lines (colors) which represents energy
associated with a specific process in the
atom
Light Energy and Photons
 Scientists associated the lines of an atomic
spectrum with changes in an electrons energy
(“Bohr Model”)
 An electron in a higher energy state will return
to a lower energy state
 The energy that is given off (emitted)
corresponds to the energy difference between
the higher and lower energy states
 The light emitted behaves like a stream of
small particles called “photons”
Electron Energy Levels
 Electrons possess energy; they are in constant motion in
the large empty space of the atom
 The arrangement of electrons in an atom corresponds to
an electron’s energy
 The electron resides outside the nucleus in one of seven
fixed energy levels
 Energy levels are quantized: Only certain energy values
are allowed
Light Energy and Photons
 The energy of a photon is related by
the equation
E = hν
 “The energy of a photon is directly
proportional to its frequency”
 “The energy of a photon is inversely
proportional to its wavelength”
Electron Energy Levels
 The different lines in an atomic spectrum
are associated with changes in an
electrons energy
 Each electron resides in a specific E
level called it’s principal quantum number
(n, where n=1, n=2…)
 Electrons closer to nucleus have lower energy
(lower n values)
 Electrons farther from the nucleus have higher
energy (higher n values)
Electron Energy Levels
 Electrons can be “excited”
to a higher E level with the
absorption of E
 The energy absorbed is
equal to the difference
between the two E states
 When an electron loses E
and falls to a lower E level,
it emits EM radiation
(photon)
Electron Energy Levels
 If the EM radiation wavelength is in the
visible spectrum a color is seen
Energy Levels of Hydrogen:
The Bohr Model
 In 1913 Bohr
developed a
quantum model
based on the
emission spectrum
for hydrogen
 The proposal was
based on the
electron in hydrogen
moving around the
nucleus in a circular
orbit
Energy Levels of Hydrogen/
The Bohr Model
 The Bohr atom
nucleus
Energy Levels of Hydrogen/
The Bohr Model
 The Bohr atom has several
orbits with a specific radius
and specific energy
 Each orbit or energy level is
identified by “n” the principal
quantum number
 Electrons can be “excited”
to a higher energy level with
nucleus
absorption of energy
 The energy absorbed and
released is equal to the
energy difference between
the two states
Energy Levels of Hydrogen/
The Bohr Model
 The energy levels calculated by the Bohr model closely
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agreed with the values obtained from the hydrogen
emission spectrum
The Bohr model did not work for other atoms
Energy levels were OK but another model was needed to
describe the location of the electron about the nucleus
Shrodinger in 1926 (DeBroglie, Heisenberg) developed
the more precise quantum mechanical model
The quantum (wave) mechanical model is the current
theory of atomic structure
Quantum Mechanical Model
 The electron is treated not as a particle but as a
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wave bound to the nucleus
The electron does not move around the nucleus in a
circular path (orbit)
Instead, the electron is found in orbitals. It is not an
circular path for the electron
An orbital indicates the probability of finding an
electron near a particular point in space
An orbital is a map of electron density in 3-D space
Each orbital is characterized by a series of numbers
called quantum numbers
Electron Energy Levels
 The energy of an electron and its distances
from the nucleus can be grouped into levels
 Principal quantum number “n” is the major
energy level in the atom
 It has values of n =1, 2, 3, etc.
 As “n” increases the size of the principal
energy level (shell) increases
Principal E level electron capacity = 2n2
Electron Sublevels
 All electrons in a principal E level (shell)
do not have the same energy
 Each principal level is divided into 1, 2, 3,
or 4 sublevels (subshells)
 An E level contains the same number of
sublevels (s, p, d, and f) as its own pr.
energy level number
# of sublevels in a principal E level = n
Electron Sublevels
 The order of the increasing energy for sublevels
(within an E level)
 The sublevels with the lowest to highest energy:
 s sublevel (holds up to 2 electrons)
 p sublevel (holds up to 6 electrons)
 d sublevel (holds up to 10 electrons)
 f sublevel (holds up to 14 electrons)
Lowest
energy
s<p<d<f
Highest
energy
Orbitals
 The third term used to describe electron
arrangement about the atomic nucleus (shells,
subshells) is the orbital
 Since the electron location cannot be known
exactly, the location of the electron is
described in term of probability, not exact
paths
 Region in space around the nucleus where
there is a high (90%) probability of finding an
electron of a specific energy
Orbitals
 Orbital shapes are 3-D regions where the
highest probability exists
 Each orbital is represented by four quantum
numbers
 Orbitals within the same sublevel differ mainly
in orientation
 Orbitals of the same type, but in different E
levels (i.e. 1s, 2s, 3s) have the same general
shape, but differ in size
s-Orbitals
 Only one type of orbital
 Spherical in shape
The larger the energy level,
the larger the sphere
 Holds two electrons
s-Orbitals
1s
Fig10_23
2s
3s
p-Orbitals
 Can only occur in n=2 or
higher levels
 Are dumb-bell in shape
 Three sub-orbitals (px, py and
pz) each holding 2 electrons for
a total of 6 electrons in a porbital
pOrbitals
z
y
x
(a)
z
z
y
y
x
x
(b)
Fig10_21
(c)
d-Orbitals
 Five possible d-orbitals
 Odd shapes
 Only possible in n=3 and
larger energy levels
 Holds a total of 10 electrons
z
z
y
z
y
y
dOrbitals
x
x
x
dyz
dxz
dxy
z
z
y
y
x
x
dx2 - y2
Fig10_24
dz2
f-Orbitals
 Seven possible types of forbital
 Shapes very difficult, so don’t
have to know
 Can hold a total of 14 electrons
 Only possible for energy levels
n=4 and higher
Writing Orbital Diagrams and
Electron Configurations
 To show how the electrons are
distributed in the E levels within an
atom
 Orbital diagrams
 Electron configurations
 The most stable arrangement of
electrons is one where the electrons
are in the lowest energy sublevels
possible
Writing Orbital Diagrams and
Electron Configurations
 The most stable arrangement of
electrons is called “ground-state
electronic configuration”
 The most stable, lowest E
arrangement of the electrons
 The GS configuration for an element
with many electrons is determined by
a building-up process
Writing Orbital Diagrams and
Electron Configurations
 For the building-up process, begin by
adding electrons to specific E levels
beginning with the 1s sublevel
 Continue in the order of increasing
sublevel energies:
1s→2s →2p →3s →3p →4s →3d →4p →5s →4d →etc.
Orbital Diagram
 The notation illustrating the electron
arrangement in terms of which energy
levels and sublevels are occupied
 Uses the building-up principal
 Hund’s Rule: When electrons are placed
in a set of orbitals of equal energy, the
orbitals will be occupied by one electron
each before pairing together
Notation
 Draw a box for each orbital
 Use an arrow up or down to
represent an electron
 Only one up and one down arrow is
allowed in a box
1s
2s
2p
Filling of Orbitals
 In General:
 Begin filling from the lowest to the
highest energy level
 If there are more than one sub-orbital
possible, electrons will spread out
first instead of doubling up
 Once each sub-orbital is filled with one
electron, they will double up, but MUST
have opposite spins (Hund’s Rule)
Orbitals Review
 s-orbitals
 Only one per n
 Can hold two electrons for a total of 2
electrons in an s-orbital
 p-orbitals
 Three per n
 Can each hold two electrons for a total
of 6 electrons in a p-orbital
Orbitals Review
 d-orbitals
 Five per n
 Can each hold two electrons for a total
of 10 electrons in a d-orbital
 f-orbitals
 Seven per n
 Can each hold two electrons for a total
of 14 electrons in an f-orbital
Orbital Diagram
 hydrogen
 Only one electron
 Occupies the 1s orbital
 helium
 Two electrons
 Both occupy the 1s orbital
 lithium
 Three electrons
1s
1s
1s
2s
 Two occupy the 1s orbital, one occupies the 2s
orbital
Electron Configurations and the
Periodic Table
 No need to memorize the filling order of the
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electron, use the periodic table
The atomic numbers are in order of
increasing sublevel
Can “build-up” atoms by reading across the
periods from left to right
By following a path of increasing atomic
number and note the various subshells as
they are encountered
Each box in the table (across a period) is an
increase in one electron
Electron Configurations and the
Periodic Table
 The elements are arranged by increasing
atomic number
 The periodic table is divided into sections
based on the type of subshell (s, p, d, or f)
which receives the last electron in the build
up process
 Different blocks on the periodic table
correspond to the s, p, d, or f sublevels
Electron Configurations and the
Periodic Table
 s-block elements (Groups 1A and 2A) gain
their last electron in an s-sublevel
 p-block elements (Groups 3A to 8A) gain
their last electron in a p-sublevel
 d-block elements (transition metals) gain
their last electron in a d-sublevel. First
appear after calcium (element 20)
 d-sublevel is (n-1) less than the period
number
 f-block elements are in the two bottom rows
of the periodic table
 f-sublevel is (n-2) less than the period
number
Subshell Filling Order
1
2
3
4
5
6
7
(n-1)d
np
ns
(n-2) f
Writing Electronic Configurations
Using Sublevel Blocks
 Locate the element, the number of
electrons is equal to the atomic number
 Lowest energy sublevel fills first, then the
next lowest following a path across each
period
 The configuration of each element builds
on the previous element
 The p, d, or f sublevels must completely
fill with electrons before moving to the
next higher sublevel
Electron Configuration Example #1
 Write the complete electron
configuration for chlorine
 Chlorine is atomic number 17 (on the
periodic table) so the neutral atom
has 17 electrons
 Writing sublevel blocks in order up to
chlorine gives:
1s22s22p63s23px
Electron Configuration Example #1
1
2
3
4
5
6
7
(n-1) d
np
ns
(n-2) f
Electron Configuration Example #1
2
2
6
2
5
Cl : 1s 2s 2p 3s 3p
2
5
or [Ne] 3s 3p
1s
2s
2p
3s
3p
Electron Configuration Example #2
 Write the complete electron
configuration for calcium
 Calcium is atomic number 20 (on the
periodic table) so the neutral atom
has 20 electrons
 Writing sublevel blocks in order up to
calcium gives:
1s22s22p63s23p64sx
Electron Configuration Example #2
1
2
3
4
5
6
7
(n-1) d
np
ns
(n-2) f
Electron Configuration Example #2
2
2
6
2
6
Ca : 1s 2s 2p 3s 3p 4s
or [Ar] 4s
1s
2s
2p
2
2
3s
3p
4s
Electron Configs
Examples
2
2
6
2
6
2
Ca : 1s 2s 2p 3s 3p 4s
or [Ar] 4s
2
Periodic Trends of the Elements
 Per. Table: Graphically represents the
behavior of the elements
 Elements are arranged by increasing
atomic number
 In the periodic table, elements with
similar properties occur at regular
intervals
 The arrangement of electrons and not
the mass that determines chemical
properties of the elements
Periodic Trends of the
Elements/Valence Electrons
 Valence electrons: The electrons in the
outermost energy level “n” (where n = 1,
2, 3 …)
 The most important (chemically)
 Always found in the outermost s or p
sublevels
 Group number equals the valence
electrons for each element in that group
 Applies only to the groups 1A-8A
Periodic Trends of the
Elements/Valence Electrons
 Group IA elements have one valence
electron: ns1
 Group IIA elements have two
valence electron: ns2
 Group VIIA elements have seven
valence electron: ns2np6
Periodic Trends of the
Elements/Valence Electrons
 Write the electron configuration for lithium
Li: 1s22s1
 Write the electron configuration for sodium
Na: 1s22s22p63s1
 Each group 1A element has a single
electron in an s-sublevel. This is the (one)
valence electron
Atomic Size
 For representative (main
group) elements only
 Describes the volume of the
electron cloud in the atoms
 Dependent upon the electron
configuration of the atoms
Atomic Size
 Within groups: The atomic
radius increases from top to
bottom
 Increase in the period number
 Principal E level (n) increases
 Valence electron is further
from the nucleus
Atomic Size
 Across periods: The atomic
radius decreases from L to R
with increasing atomic number
 Each element increases in
proton and electron number
 Increase in + nuclear charge
 Valence electrons pulled closer
to the nucleus
Size of Atoms and
Their Ions
 The formation of a positive ion
requires the loss of one or more
valence electrons
 Loss of the outermost (valence)
causes a reduction in atomic size
 Positive ions are always smaller
than their parent ions
Size of Atoms and
Their Ions
 The formation of a - ion requires
the addition of one or more
electrons to the valence shell of
an atom
 There is no increase in + nuclear
charge to offset the added
electron’s - charge
 Increase in size due to repulsion
between electrons
Ionization Energy
 The minimum energy required to
remove one electron from an atom of
an element
 The more tightly an electron is held,
the higher the ionization energy
Ionization Energy
 In the same group (top to bottom)
Ionization Energy decreases
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Energy required to remove an electron decreases
Due to larger principal energy level (larger n value)
This puts outer electron farther from nucleus
As n increases, ionization energy decreases
 Across same period (left to right)
Ionization Energy increases
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Metals (left end) have lower ionization E
Tend to lose electrons to form + ions
Nonmetals (right end) have higher ionization E
Tend to gain electrons in chemical reactions
 End