Download 3ACh 9 BW Fall 2011

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
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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
 Electromagnetic radiation carries (radiant) 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
Light: Electromagnetic Radiation
 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
Light: Electromagnetic Radiation
 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
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 (and vice-versa)
C=2λ·½ѵ C=½λ·2ѵ
C = λѵ
Waves
frequency
wavelength
frequency
wavelength
C = speed of light
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 highenergy gamma and X-rays to very low-energy
radio and TV waves
The Electromagnetic Spectrum
 EM radiation is 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
 The electromagnetic spectrum ranges from high-energy
gamma and X-rays to very low-energy radio and TV waves
 The visible region of light is a narrow range of wavelengths
between these two extremes
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 element produces a different
discontinuous spectra
 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
White light produces a
continuous spectra
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 bright 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
Light emitted from hydrogen produces
only specific wavelengths of light
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
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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
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)
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 is called a quantum
A photon (of light)
The Bohr Model: Excitation and
Emission
 The energy of a photon is related by the
equation:
E = hѵ
Eѵc==hc/λ
c/λ
λѵ
 “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
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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
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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 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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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
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
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
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
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)
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 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
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
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
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-subshell
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-subshell 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-subshell
 All of equal energy
 The five orbitals first occur in the n=3
shell
 Odd shapes (don’t need to know them)
 d-subshell can hold a total of 10
electrons, 2 electrons in each of five dorbitals
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
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f-subshell
All of equal energy
The seven orbitals first occur in the n=4
shell
Shapes are very difficult, so you don’t need
to know them either
f-subshell can hold a total of 14 electrons,
2 electrons in each of seven f-orbitals
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
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
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.
Writing Orbital Diagrams and
Electron 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
 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 Principle: 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
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
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
 You can “build-up” atoms by reading across the
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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
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
 The specific location of an element in the periodic table can
be used to obtain information about its electron configuration
 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 element’s electron
configuration to differ from the preceding element
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
Principal quantum number (n)=
number of the period
Subshell Filling Order
1
2
3
4
5
6
7
(n-1)d
ns
(n-2) f
np
Writing Electron Configurations
from the
Periodic Table
 Locate the element, the number of electrons is
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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
 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
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
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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
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
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:

 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
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
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