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## Electrons in Atoms

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Title: Electrons in Atoms

1
Chapter 9
• Electrons in Atoms
• and the
• Periodic Table

2
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)

3
• 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
• 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
energy through space and travels in waves at the
speed of light
• Waves are periodic The pattern of peaks and
troughs repeats itself at regular intervals

4
• 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
• Shorter wavelengths are higher in energy
• Longer wavelengths, are lower in energy

5
• Frequency (u) is the number of waves that pass a
fixed point in one unit of 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 the rest position to crest position
or from rest position to trough position

amplitude
6
Wavelength and Frequency
• Because all EM radiation travels at the speed of
light (c), a relationship exists between
wavelength and frequency
• This is an inverse relationship so that if the
wavelength doubles, the frequency is halved. If
the wavelength is halved, the frequency doubles
(and vice-versa)

C ??
C 2 ? ½ ?
C ½ ? 2 ?
7
Waves
C ??
frequency
wavelength
frequency
C speed of light
wavelength
8
The Electromagnetic Spectrum
• Light (radiant) energy is the energy of
electromagnetic waves and it is 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 ranges from
high-energy gamma and X-rays to very low-energy

9
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

10
• The electromagnetic spectrum ranges from
high-energy gamma and X-rays to very low-energy
• The visible region of light is a narrow range of
wavelengths between these two extremes

11
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)

12
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

13
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

Each element produces a different discontinuous
spectra
White light produces a continuous spectra
14
Emission Spectra
• Scientists first detected the line spectrum of
hydrogen (mid-1800s) which produced only four
lines

15
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 bright line spectrum of hydrogen,
Bohrs theory of the hydrogen atom began with
this idea and assumed the electrons move in
circular orbits around the nucleus

Light emitted from hydrogen produces only
specific wavelengths of light
16
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

17
The Bohr Model Atoms with Orbits
• The Bohr atom has several orbits with a specific
• 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
• 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

nucleus
18
The Bohr Model Atoms with Orbits
• The different lines in an emission spectrum are
associated with changes in an electrons energy
• Each electron resides in a specific E level
called its principal quantum number (n, where
n1, n2)
• Electrons closer to nucleus have lower energy
(lower n values)
• Electrons farther from the nucleus have higher
energy (higher n values)

19
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
• 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 is called a quantum

A photon (of light)
20
The Bohr Model Excitation and Emission
• The energy of a photon is related by the
equation
• 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)

E h?
c ??
? c/?
E hc/?
21
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 electrons energy
• The electron resides outside the nucleus in one
of seven fixed energy levels
• Energy levels are quantized Only certain energy
values are allowed

22
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)

23
The Bohr Model Electron Energy Levels
• If the EM radiation wavelength is in the visible
spectrum a color is seen

24
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

25
The Quantum-Mechanical ModelFrom Orbits to
Orbitals
• The quantum-mechanical model gives a new way to
view electronic structure
• This model combines 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 electrons
location cannot be described exactly
• The electrons location is described as region of
space (probability) where the electron will be at
any given instant

26
The Quantum-Mechanical ModelFrom Orbits to
Orbitals
• The electron is treated not as a particle but as
a 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 a 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

27
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 various
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 shell electron capacity 2n2
28
The Quantum-Mechanical Model Electron Sublevels
• All electrons in a principal shell (E level) 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

29
The Quantum-Mechanical Model Electron Sublevels
• The number of subshells (sublevels) within a
principle shell (E level), n, varies
• Each principal shell is divided into 1, 2, 3, or
4 subshells
• Subshells are identified by a number and a
letter s, p, d, and f
• Each principal shell contains the same number of
subshells as its own principal shell number

Two electrons per subshell
No. of subshells in a principal shell n
30
The Quantum-Mechanical Model Electron Sublevels
• The order of the increasing energy for subshells
(within an shell)
• The subshells with the lowest to highest energy
• s subshell (holds up to 2 electrons)
• p subshell (holds up to 6 electrons)
• d subshell (holds up to 10 electrons)
• f subshell (holds up to 14 electrons)

s lt p lt d lt f
Lowest energy
Highest energy
31
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

32
Quantum-Mechanical Orbitals
• Each orbital can hold up to 2 electrons
• Each subshell is composed of one or more orbitals
• One orbital in an s-subshell
• Three orbitals in a p-subshell
• Five orbitals in a d-subshell
• Seven orbitals in an f-subshell
• Orbitals within the same subshell differ mainly
in orientation

33
Quantum-Mechanical Orbitals
• The orbitals in each of the four subshells
(sublevels) have characteristic shapes
• Orbitals in an s-subshell do not have the same
shape as orbitals in a p-subshell, etc.
• Orbitals of the same type, but in different
principal shells/E levels (e.g. 1s, 2s, 3s) have
the same general shape, but differ in size
• The nucleus is located at the center of each
orbital

34
Quantum-Mechanical Orbitals s-Orbitals
• There is one s-orbital in each s-subshell
• Every principal shell contains only one
s-orbital within an s-subshell
• S-orbitals are spherical in shape
• The larger the principal shell (energy level),
the larger the sphere
• An s-sublevel can hold a total of two electrons
within the s-orbital

35
Quantum Mechanical Orbitalss-Orbitals
• The spherical s-orbital gets larger as n increases

nucleus
36
Quantum Mechanical Orbitalsp-Orbitals
• The p-orbitals come in sets of three within each
p-subshell
• All of equal energy
• The three p-orbitals first occur in the n2 (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-subshell can hold a total of six electrons, two
electrons in each of the p-orbitals (px, py and
pz)

37
Quantum Mechanical Orbitalsp-Orbitals
p-orbitals have a two-lobe, dumbbell shape. The
nucleus is at the point where the two lobes meet
nucleus
38
Quantum Mechanical Orbitalsd-Orbitals
• d-orbitals come in sets of five within each
d-subshell
• All of equal energy
• The five orbitals first occur in the n3 shell
• Odd shapes (dont need to know them)
• d-subshell can hold a total of 10 electrons, 2
electrons in each of five d-orbitals

39
d-Orbitals
40
f-Orbitals
• f-orbitals come in sets of seven within each
f-subshell
• All of equal energy
• The seven orbitals first occur in the n4 shell
• Shapes are very difficult, so you dont need to
know them either
• f-subshell can hold a total of 14 electrons, 2
electrons in each of seven f-orbitals

41
Electron Configurations How Electrons
Occupy Orbitals
• Two ways to show how the electrons are
distributed in the principal shells within an
atom
• Orbital diagrams
• Electron configurations
• The most stable arrangement of electrons is one
where the electrons are in the lowest energy
subshells possible

42
Electron Configurations How
Electrons Occupy Orbitals
• The most stable arrangement of electrons is
called ground-state electronic configuration
• The most stable, lowest energy arrangement of
the electrons
• The GS configuration for an element with many
electrons is determined by a building-up process

43
Writing Orbital Diagrams and Electron
Configurations
• For the building-up process, begin by adding
electrons to specific principal shells (E levels)
beginning with the 1s subshell
• Continue in the order of increasing subshell
energies

1s?2s ?2p ?3s ?3p ?4s ?3d ?4p ?5s ?4d ?etc.
44
Writing Orbital Diagrams andElectron
Configurations
• The notation illustrates the electron arrangement
in terms of which energy levels (shells) and
sublevels (subshells) are occupied
• The orbital diagram uses the building-up
principal
• Hunds 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

45
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 Principle No more than two
electrons can be placed in a single orbital and
must be paired (have spins in opposite
directions)

orbital
46
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
47
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 (Hunds Rule)
• Once each orbital is filled with one electron
they will pair up but must have opposite spins
(Pauli Exclusion Principal)

48
Orbital Diagrams
• s-orbitals
• Only one per n
• Can hold two electrons for a total of two
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

49
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

50
Orbital Diagrams
• hydrogen
• Only one electron
• Occupies the 1s orbital
• helium
• Two electrons
• Both occupy the 1s orbital
• lithium
• Three electrons
• Two occupy the 1s orbital, one occupies the 2s
orbital

1s
1s
1s
2s
51
Electron Configurationsand 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

52
Electron Configurations and thePeriodic Table
• You can build-up atoms by reading across the
periods from left to right
• It is not necessary to memorize the filling order
of the electron, just use the periodic table
• Follow a path (left to right) across each period
(row) of the table and note the various subshells
as they are encountered
• The atomic numbers are increasing across each
period and this corresponds to increasing
subshell energy
• Since atomic numbers are increasing, each box in
the table (across a period) is also an increase
in one electron

53
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

54
Electron Configurations and thePeriodic Table
• The specific location of an element in the
periodic table can be used to obtain information
• An electron configuration is a statement of how
many electrons an atom has in each of its
subshells
• To write a complete electron configuration
• 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
elements electron configuration to differ from
the preceding element

55
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

56
Subshell Filling Order
Principal quantum number (n) number of the
period
(n-1)d
np
(n-2) f
ns
57
Writing Electron Configurationsfrom 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

58
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
59
Electron Configuration Example 1
(n-1) d
np
(n-2) f
ns
60
Electron Configuration Example 1
Orbital diagram
Hunds Rule
1s
2s
2p
3s
3p
61
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
62
Electron Configuration Example 2
(n-1) d
np
(n-2) f
ns
63
Electron Configuration Example 2
Orbital diagram
Hunds Rule
1s
2s
2p
3s
3p
4s
64
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

abbrev. electron configuration
65
Electron Configurations and the Periodic Table
• The periodic 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

66
Valence Electrons
• Valence electrons are those electrons in the
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

67
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

68
Periodic Trends of the Elements/Valence Electrons
• Write the electron configuration for lithium
• Write the electron configuration for sodium
• Each group 1A element has a single electron in an
s-sublevel. This is the (one) valence electron

Li 1s22s1
Na 1s22s22p63s1
69
Periodic Trends of the Elements/Valence Electrons
• The periodic table list elements by increasing
atomic number and arranges them in groups with
similar chemical properties
• Similar chemical properties arise in every eighth
element due to the similarity in electronic
configurations (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

70
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 obtain the stable electron
configuration of the next higher noble gas

71
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)

72
Atomic Size
• Within groups
• The period number increases downward in a group
• Principal E level (n) increases
• Valence electron is further from the nucleus
• Across periods
• The atomic radius decreases from LEFT to RIGHT
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

73
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) electron causes a
reduction in atomic size
• Positive ions are always smaller than their
parent ions

74
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 electrons - charge
• Increase in size due to repulsion between
electrons

75
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

76
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

77
• End