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Chapter 9
Electrons in Atoms
and the
Periodic Table
Homework
 Assigned Problems (odd numbers only)
 “Questions” (page 310-11)
 “Problems” 31 to 89 (page 311-14)
 “Cumulative Problems” 91-113
(page 315-17)
 Highlight Problems 115 (optional)
Light: Electromagnetic Radiation
 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
 A form of energy that travels through space at the
speed of light
 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
Light: Electromagnetic Radiation
 Electromagnetic radiation is radiant (light) energy
that travels in waves at the speed of light
 Waves are periodic: The pattern of peaks and
troughs repeats itself at regular intervals
 The waves have three basic characteristics:
wavelength, frequency, and speed
 Wavelength (l) is the distance (in nm) between
neighboring peaks in a wave
 The highest point on the wave is a peak
 The shorter the wavelength, the higher the energy of
the light
 The longer the wavelength, the lower the energy of the
light
Light: Electromagnetic Radiation
 Frequency (u) is the number of waves that
pass a point in a given time
 measured in Hertz (Hz),
 1 Hz = 1 wave/sec = 1 sec-1
 Velocity (v = how fast the wave is moving)
 c = speed of light
 3.00 x 108 m/s
 Amplitude the height of the wave. It is the
distance from crest to crest or from rest
position to trough position
Wavelength and Frequency
 Because all EM radiation travels at the speed of
light (c), a relationship exists between
wavelength and frequency
C = λѵ
 This is an inverse relationship so that if the
wavelength doubles, the frequency is halved. If
the wavelength is halved, the frequency doubles
C = λѵ
Waves
frequency
wavelength
frequency
wavelength
C = speed of light
The Electromagnetic Spectrum
 Light (radiant) energy is the energy of
electromagnetic waves
 It is visible and invisible and 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
The Electromagnetic Spectrum
The electromagnetic spectrum range:
 from high-energy gamma and X-rays
 short wavelength, λ
 high frequency, ѵ
 to very low-energy radio and TV waves
 long wavelength, λ
 low frequency, ѵ
 The visible region of light is a narrow range
of wavelengths between these two extremes
The Electromagnetic Spectrum
 EM radiation is classified by wavelength:







 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
Light Emission by Different Elements
 When white light passes through a prism it
separates and produces a continuous rainbow
of colors from (red, orange, yellow, green,
blue, indigo, and, violet)
 From red light to violet light the wavelength
becomes shorter (700 nm to 400 nm)
Light Emission by Different Elements
 When an element is
heated its atoms
absorb energy and
re-emits that energy
 Light is produced
 If this light is passed
through a prism, it
does not produce a
continuous rainbow,
only certain colors
Emission Spectra
 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
 Lines are also produced in the infrared and
ultraviolet regions
Emission Spectra
 Scientists first detected the line
spectrum of hydrogen (mid-1800’s)
which produced only four lines
Emission Spectra
 Scientist could not explain why atoms excited
with energy produced discontinuous spectra
 After the discovery of the nuclear structure of
the atom (Rutherford, 1911), scientist thought
of the atom as a microscopic solar system
with electrons orbiting the nucleus
 To explain the line spectrum of hydrogen,
Bohr’s theory of the hydrogen atom began
with this idea and assumed the electrons
move in circular orbits around the nucleus
Emission Spectra for 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
The Bohr Model: Atoms with Orbits
 The Bohr atom has several orbits with

nucleus




a specific radius and specific energy
Each orbit or energy level is identified
by “n” the principal quantum number
The values of n are positive, whole
numbers 1, 2, 3, etc.
The principal energy level (n =1) has
the lowest energy and the smallest
radius
Electrons can be “excited” to a higher
energy level with 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 Bohr atom
nucleus
The Bohr Model: Atoms with Orbits
 The different lines in an emission spectrum are associated with
changes in an electron’s 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)
nucleus
The Bohr Model:
Excitation and Emission
 Scientists associated the lines of an atomic
spectrum with changes in an electrons energy
(“Bohr Model”)
 An electron excited to a higher energy state
will return to a lower energy state
 The energy that is given off (emitted) is a
photon of light that corresponds to the energy
difference between the higher and lower
energy states
 This precise amount of energy called a
quantum
The Bohr Model: Excitation and
Emission
 The energy of a photon is related by the
equation:
E = hѵ
 “The energy of a photon is directly proportional
to its frequency”
 “The energy of a photon is inversely
proportional to its wavelength”
 Energy transitions between orbits closer
together produce photons of light with longer
wavelengths (lower energy)
The Bohr Model: 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
The Bohr Model: 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)
The Bohr Model: Electron Energy Levels
 If the EM radiation wavelength is in the
visible spectrum a color is seen
The Bohr Model: Electron Energy Levels
 The energy levels calculated by the Bohr model




closely agreed with the values obtained from the
hydrogen emission spectrum
The Bohr model did not work for other atoms
Energy levels were OK but model could not predict
emission spectra for an element with more than one
electron
Shrodinger in 1926 (DeBroglie, Heisenberg)
developed the more precise quantum-mechanical
model
The quantum (wave) mechanical model is the
current theory of atomic structure
The Quantum-Mechanical Model:
From Orbits to Orbitals
 The quantum-mechanical model gives a new way to





view electronic structure
This model incorperates the wavelike and particle-like
behavior of the electron
For the hydrogen atom, the allowed energy states are
the same as that predicted by the Bohr model
The Bohr model assumes the electron is in a circular
orbit of some distance from the nucleus
In the quantum-mechanical model, the electron’s
location cannot be described exactly
The electron’s location is described as region of space
(probability) where the electron will be at any given
instant
The Quantum-Mechanical Model:
From Orbits to Orbitals
 The electron is treated not as a particle but as a
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
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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
The Quantum-Mechanical Model:
Electron Energy Levels
 Electrons with higher E will tend to be farther
from the nucleus than those of lower E
 The energy of an electron and its distances from
the nucleus can be grouped into levels or shells
 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
The Quantum-Mechanical Model:
Electron Sublevels
 All electrons in a principal E level (shell)
do not have the same energy
 The energy of electrons in the same shell
have energies close in magnitude, but
not identical
 The range of energies for electrons in a
shell is due to the existence of electron
subshells or energy sublevels
 An electron subshell is an energy level
within an electron shell in which electrons
all have the same energy
The Quantum-Mechanical Model:
Electron Sublevels
 The number of sublevels (subshells) within a pr.
energy level (shell), n, varies
 Each principal E level (shell) is divided into 1, 2,
3, or 4 sublevels (subshells)
 Sublevels (subshells) are identified by a number
and a letter: s, p, d, and f
 Each principal E level (shell) contains the same
number of sublevels (subshells) as its own
principal energy level number (shell):
# of sublevels in a principal E level = n
The Quantum-Mechanical Model:
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
Quantum-Mechanical 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
 The orbital is a region of space where an electron
assigned to that orbital is likely to be found
 Region in space around the nucleus where there
is a high (90%) probability of finding an electron
of a specific energy
Quantum-Mechanical Orbitals
 Each orbital can hold up to 2 electrons
 Each sublevel is composed of one or more orbitals
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
One orbital in an s-sublevel
Three orbitals in a p-sublevel
Five orbitals in a d-sublevel
Seven orbitals in an f-sublevel
 The orbitals in each of the four sublevels (subshells)
have characteristic shapes
 Orbitals within the same sublevel differ mainly in
orientation
Quantum-Mechanical Orbitals
 Orbitals of the same type, but in different E
levels (e.g. 1s, 2s, 3s) have the same general
shape, but differ in size
 Orbitals in an s-subshell do not have the same
shape as orbitals in a p-subshell, etc.
 The nucleus is located at the center of each
orbital
Quantum-Mechanical Orbitals:
s-Orbitals
 There is one s-orbital in each s-sublevel
 Every principal energy level contains
only one s-orbital in an s-sublevel
 S-orbitals are spherical in shape
 The larger the energy level, the larger
the sphere
 An s-sublevel can hold a total of two
electrons within the s-orbital
Quantum Mechanical Orbitals:
s-Orbitals
 The spherical s-orbital gets larger as n
increases
nucleus
1s
Fig10_23
2s
3s
Quantum Mechanical Orbitals:
p-Orbitals
 The p-orbitals come in sets of three within each
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
p-sublevel
All of equal energy
The three p-orbitals first occur in the n=2 (or
higher) levels
P-orbitals are dumb-bell in shape
The three orbitals within a p-sublevel are
oriented at right angles to one another and
labeled as (px, py and pz)
p-sublevel can hold a total of six electrons, two
electrons in each of the p-orbitals (px, py and pz)
Quantum Mechanical Orbitals:
p-Orbitals
p-orbitals have a two-lobe, dumbbell
shape. The nucleus is at the point where the two
lobes meet
nucleus
z
y
x
(a)
z
z
y
y
x
x
(b)
Fig10_21
(c)
Quantum Mechanical Orbitals:
d-Orbitals
 d-orbitals come in sets of five within
each d-sublevel
 All of equal energy
 The five orbitals first occur in the n=3
shell
 Odd shapes
 d-sublevel can hold a total of 10
electrons, 2 electrons in each of five
d-orbitals
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
 f-orbitals come in sets of seven within each




f-sublevel
All of equal energy
The seven orbitals first occur in the n=4
shell
Shapes very difficult, so don’t have to
know
f-sublevel can hold a total of 14 electrons,
2 electrons in each of seven f-orbitals
Electron Configurations:
How Electrons Occupy Orbitals
 Two ways too show how the
electrons are distributed in the
principal 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
Electron Configurations:
How Electrons Occupy Orbitals
 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.
Writing Orbital Diagrams and
Electron Configurations
 The notation illustrates 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
Electron Spin
 Electrons behave as if they are spinning on an axis
 A spinning electron behaves like a small bar magnet with north
and south poles
 Small arrows (pointed up or downward) are used to indicate the
two orientations of spin
 Two electrons in the same orbital must spin in opposite
directions
 Pauli Exclusion Principal: No more than two electrons can be
placed in a single orbital and must be paired (have spins in
opposite directions)
orbital
Orbital Diagrams
 Orbital Diagram Notation:
 Draw a box to represent each orbital
 Use an arrow up or down to represent an
electron
 Two electrons in the same orbital (box)
must have spins in opposite directions:
Only one up and one down arrow is allowed
in a box (paired electrons)
1s
2s
2p
Orbital Diagrams
 In General:
 Begin filling from the lowest to the highest
energy level
 If there is more than one orbital possible
(e.g. px, py, pz), place electrons alone
before pairing up (Hund’s Rule)
 Once each orbital is filled with one electron
they will pair up but must have opposite
spins (Pauli Exclusion Principal)
Orbital Diagrams
 s-orbitals
 Only one per n
 Can hold two electrons for a total of 2
electrons in an s-sublevel
 p-orbitals
 Three per n
 Each can hold two electrons for a total
of 6 electrons in a p-sublevel
Orbital Diagrams
 d-orbitals
 Five per n
 Each can hold two electrons for a total
of 10 electrons in a d-sublevel
 f-orbitals
 Seven per n
 Each can hold two electrons for a total
of 14 electrons in an f-sublevel
Orbital Diagrams
 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
 The elements in the periodic table are arranged in order of
increasing atomic number
 The basic shape and structure of the table is consistent with
and can be explained by the sequence used to build electron
configurations
 The table is divided into sections based on the type of
subshell (s, p, d, or f) that receives the last electron in the
building-up process
Electron Configurations and the
Periodic Table
 No need to memorize the filling order of the




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 specific location of an element in the periodic
table can be used to obtain information about its
electron configuration
 In order to write a complete E config. The order in
which the various subshells are filled can be
obtained by following a path of increasing atomic
number through the table (also taking account of the
various subshells along the path)
 The periodic table can be used to determine the shell
in which the last electron added is located
 It is this last electron added that causes an element’s
E config. to differ from the preceding element
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 Electron Configurations
from the
Periodic Table
 Locate the element, the number of electrons is




equal to the atomic number
Start at hydrogen and move from box to box,
in order of increasing atomic number
The 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
Orbital diagram
Hund’s Rule
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
2
2
Orbital diagram
Hund’s Rule
1s
2s
2p
3s
3p
4s
Electron Configurations
Examples
 May also use the condensed (inner) electron
configuration
 This shorthand notation uses the noble gas
that precedes a particular element and places
it inside square brackets
Noble gas core
2
2
6
2
6
2
[
Ca : 1s 2s 2p 3s 3p ]4s
or [Ar] 4s
2
abbrev. electron
configuration
Electron Configurations and the
Periodic Table
 Per. table graphically represents the
behavior of the elements described by
periodic law
 Elements are arranged by increasing
atomic number
 In the periodic table, elements with
similar properties occur at regular
intervals (in the same vertical column)
 The arrangement of electrons and not
the mass determines chemical properties
of the elements
Valence Electrons
 Valence electrons are those electrons in the
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


outermost (highest) energy level “n” (where n
= 1, 2, 3 …)
Those electrons not in the outermost (highest)
energy level are called core electrons
Valence electrons are the most important
(chemically)
Always found in the outermost s or p sublevels
in the representative elements
For elements in columns 1A-8A, group
number equals the number of valence
electrons
Valence Electrons
 All elements within a column (group) have the same
number of valence electrons and similar outer
electron configurations
 Group IA elements have one valence electron: ns1
 Group IIA elements have two valence electrons: ns2
 Group IIIA elements have three valence electrons:
ns2np1
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




Periodic Trends of the
Elements/Valence Electrons
The periodic table list elements in increasing
atomic number and arranges them is groups
with similar chemical properties
Similar chemical properties arise in a
repeating pattern due to repeating similarity in
electronic configuration (every eighth element
for main group elements)
Across a period, elements become less
metallic and more nonmetallic
Metals tend to lose electrons in chemical
reactions
Periodic Trends of the
Elements/Valence Electrons
 Alkali metals lose their one and only one valence
electron in chemical reactions forming an ion with a
single positive charge and a stable noble gas
electronic configuration
 Group IIA metals lose their two valence electrons in
chemical reactions forming an ion with a 2+ charge
and a stable noble gas electronic configuration
 Group VIIA nonmetals readily gain one electron in
chemical reactions forming an ion with a single
negative charge and obtains the stable electron
configuration of the next higher noble gas
Atomic Size
 Atoms are considered spherical in shape and
their size (atomic radius) is very dependent on
the electronic configuration of the atom
 The electronic configuration gives trends in
atomic size within groups and across periods in
the periodic table (representative elements)
 Within groups, the atomic radius increases with
the period number (increase from top to bottom)
 Across periods, the atomic radius decreases from
left to right with increasing atomic number
(decrease from left to right)
Atomic Size
 Within groups:
 Increase in the period number
 Principal E level (n) increases
 Valence electron is further from the nucleus
 Across periods:
 The atomic radius decreases from L to R with
increasing atomic number
 As atomic number increases, so does the
number of electrons
 The increase in positive charge pulls the
outermost electrons 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 negative 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 (physical state is a gas)
 The more tightly an electron is held,
the higher the ionization energy
 The trend in ionization energy parallels
the metallic to nonmetallic trend in the
chemical properties of the elements in
a period
Ionization Energy
 In the same group (top to bottom)
Ionization Energy decreases




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




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