Chapter 7
Electron Structure of the Atom
Electromagnetic Radiation and Energy
The Bohr Model of the Hydrogen Atom
The Modern Model of the Atom
Periodicity of Electron Configurations
Valence Electrons for Main-Group
Elements
• Electron Configurations for Ions
• Periodic Properties of Atoms
•
•
•
•
•
7-1
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
White Light
•
Is composed of
different colors that can
be separated by a
prism
–
•
Water acts as a prism for
sunlight, giving the effect
of a rainbow
Sources of white light
–
–
Sun
Regular (incandescent)
light bulbs
7-2
Electromagnetic Radiation
•
•
•
•
•
A form of energy
Travels through space at the speed of light (3.0 x
108 m/s) as oscillating waves
Is both an electric and a magnetic field
Also called radiant energy
Some examples of EM radiation
– Light
– X-rays
7-3
1
Differentiating the Kinds of
Electromagnetic Radiation
• Two main characteristics
– Wavelength (λ)
– Frequency (ν)
Figure 7.5
7-4
Differentiating the Kinds of
Electromagnetic Radiation
• Wavelength (λ)
– The distance between two corresponding
points on a wave
– Units are same as length - m, or commonly
nm (10-9 m)
7-5
Differentiating the Kinds of
Electromagnetic Radiation
• Frequency (ν)
– A measure of the number of wave cycles that
move through a point in space in 1 s
– Units are hertz (Hz) which are the same as
inverse seconds (1/s)
7-6
2
Frequency, Wavelength, and the
Electromagnetic Spectrum
• Frequency and wavelength are inversely proportional
– i.e. as one increases the other decreases
c = λν
Where c = speed of light (3.0 x 108 m/s), λ = wavelength (in
meters), and ν = frequency (in Hz)
7-7
Frequency, Wavelength, and the
Electromagnetic Spectrum
• Which has a greater frequency, the red
light or the green light? Which has the
greater wavelength?
Figure 7.7
7-8
•
Duality of Light
Light exists as both waves and particles (or
packets of light called photons)
– Characteristics of waves
• Frequency
• Wavelength
c = λν
– Characteristic of photons
• Energy of a photon
– Is directly proportional to the frequency
and inversely proportional to the
wavelength
Ephoton = hν
Where Ephoton = energy of the photon (in Joules),
h = Planck’s constant (6.626 x 10-34 Js), and ν =
frequency (in Hz)
7-9
3
λ, ν, and Ephoton
c = λν
Ephoton = hν
• Using algebra, we can manipulate these
two equations several ways:
• For c = λν,
– We can solve for λ:
λ=c/ν
or ν:
ν=c/λ
•
For Ephoton = hν
– We can substitute c / λ for ν, giving us the equation:
Ephoton = (hc) / λ
This equation shows the inverse proportionality between
Ephoton and λ (wavelength)
7 - 10
Practice – λ, ν, and Ephoton
• If the wavelength of a microwave
beam is 11.5 cm, then what are
the frequency of the radiation
and the energy of its photons?
7 - 11
Practice Solutions – λ, ν, and Ephoton
• If the wavelength of a microwave
beam is 11.5 cm, then what are the
frequency of the radiation and the
energy of its photons?
λ = 11.5 cm
First, convert to meters:
λ = 11.5 cm x (1 m/100 cm) = 1.15 x 10-3 m
ν=c/λ
8
ν = (3.0 x 10 m/s)/1.15 x 10-3 m = 2.6 x 1011 s-1
7 - 12
4
Line Spectra
•
Continuous spectrum
–
–
•
Contains all the
wavelength of light in
the visible spectrum
Produced by white light
Line Spectrum
–
–
–
Contains a pattern of
distinct colored lines,
each representing a
single wavelength of
light
Produced by an element
that has been heated or
given an electric charge
Each element has a
distinct line spectra,
which is also called
“atomic fingerprint”
7 - 13
Energy is Quantized!
• Max Planck first hypothesized that energy produced by
atoms can only have certain values and is therefore
quantized.
• That’s the reason why only distinct lines are seen in
element line spectras. Energy is quantized and can only
exist at certain wavelengths.
7 - 14
Bohr Model
Niels Bohr hypothesized
that electrons orbit the
nucleus just as the
planets orbit the sun
(planetary model).
• He labeled the electron
orbits with a number,
starting with 1 closest to
the nucleus and
increasing as the orbits
get further away from the
nucleus.
– The number is known
as the Principal
Quantum Number (n).
•
7 - 15
5
Bohr Model
•
•
•
Orbits have a fixed radius.
The orbit with the lowest
energy is closest to the
nucleus. The energy of
each orbit increases as
the orbits get further away
from the nucleus.
When an electron jumps
from one orbit to another,
it absorbs or emits energy
according to the equation:
∆E = Ef – Ei
7 - 16
Modern Model of the Atom
• The modern model of the
atom is based on
Schrodinger’s mathematical
model of waves
• This model describes
electrons as occupying
orbitals, not orbits.
– Orbitals
• Three dimensional
regions in space
where electrons are
likely to be found, not
a circular pathway
– Principal energy level
• Orbitals of similar size
Figure 7.11
7 - 17
Modern Model of the Atom
Figure 7.12
7 - 18
6
Orbitals
Figure 7.13
•
•
Come in different shapes and sizes.
– Lower energy orbitals are smaller.
– Higher energy orbitals are larger and extend further
away from the nucleus.
Four most common types are s, p, d, and f.
– Also known as sublevels
• Consists of just one type of orbital at a specific
energy level
• The number of sublevels is equal to n, the
Principal Quantum Number
7 - 19
s Orbitals
Figure 7.14
7 - 20
p Orbitals
7 - 21
7
d Orbitals
7 - 22
Hydrogen Orbital Diagram
• Orbital diagrams
– Show the sublevels and orbitals that can exist at
each principal energy level
– Each box represents an orbital
– Groups of boxes represent sublevels
• In the hydrogen atom only, the sublevels within a
principal energy level all have the same energy.
7 - 23
Multielectron Orbital Diagram
• In the multielectron atoms, the sublevels within a
principal energy level have different energy levels.
7 - 24
8
Orbital Diagram Rules
•
Two principles and 1 rule determine how the
electrons are filled in the principal energy
levels and sublevels.
– Aufbau principle
•
Electrons fill orbitals starting with the lowest-energy
orbitals.
– Pauli exclusion principle
•
A maximum of two electrons can occupy each orbital, and
they must have opposite spins.
– Hund’s rule
•
•
Electrons are distributed into orbitals of identical energy
(same sublevel) in such a way as to give the maximum
number of unpaired electrons.
Electrons are always filled in their ground
state, or lowest energy state.
7 - 25
Filling Orbital Diagrams
Pg. 251 (Carbon’s orbital diagram)
7 - 26
Orbital Diagrams for the 1st Ten
Elements
7 - 27
9
Electron Configurations
• Shorthand notation which shows the
distribution of electrons among sublevels
• When we write electron configurations, we write
the number of the principal quantum number
followed by a symbol for the sublevel, and then
add a superscript to each sublevel symbol to
designate the number of electrons in that
sublevel.
Carbon has 6 electrons.
Therefore, using the orbital diagram we obtain:
1s22s22p2
7 - 28
Periodicity of Electron Configurations
• Can you tell the patterns among the following groups of
elements?
– Alkali Metals (Group IA (1))
Li
1s22s1
Na
1s22s22p63s1
– Alkali Earth Metals (Group IIA (2))
Mg
1s22s22p63s2
Ca
1s22s22p63s23p64s2
– Halogens (Group VIIA (17))
Cl
1s22s22p63s23p5
Br
1s22s22p63s23p64s23d104p5
– Noble Gases (Group VIIIA (18))
Ne
1s22s22p6
Ar
1s22s22p63s23p5
7 - 29
Using the Periodic Table for
Electron Configurations
• The periodic table can be used to fill orbital diagrams or to
find electron configurations.
– First, we need to separate the periodic table into blocks.
• Blocks contain elements with the same highest-energy
sublevel.
7 - 30
10
Using the Periodic Table for Electron
Configurations
2
• Example: 1s
• Principal quantum number (1)
– Same as the period number for s and p
• Ex. 3s and 3p are both in Period 3
– (Period number – 1) for d
• Ex. 3d is in Period 4
– (Period number – 2) for f
• Ex. 4f is in Period 6
• Sublevel (s) – also called the secondary quantum
number
– Labeled according to the block you’re in
• Number of electrons - superscript (2)
– The number of elements in the block in the period
7 - 31
The Principal Quantum Number
and Sublevel on the Periodic Table
Figure 7.21
7 - 32
An Example of Electron
Configuration
Manganese (Mn) has 25 electrons.
Start at the top left corner of the periodic
table and move from left to right, top to
bottom.
The period # is 1, groups 1 and 2 are the s
block, and there are two elements in the s
block in Period 1 (when we move He next to H
on the periodic table), so we write 1s2.
7 - 33
11
An Example of Electron
Configuration
Since there are no more elements in period 1,
we move to the next period #, which is 2,
groups 1 and 2 are the s block, and there are
two elements in the s block in period 2, so we
write 2s2.
1s22s2
Moving across the periodic table, the period #
is still 2, groups 13-18 are the p block, and
there are 6 elements in the p block in period
2, so write 2p6.
1s22s22p6
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7 - 34
An Example of Electron
Configuration
There are no more elements on that period,
so let’s go to period 3. In period 3, starting at
the left, groups 1 and 2 are the s block, and
there are 2 elements in the s block on period
3, so we write 3s2.
1s22s22p63s2
In the p block in period 3, there are 6
elements so we write 3p6.
1s22s22p63s23p6
7 - 35
An Example of Electron Configuration
1s22s22p63s23p6
Adding together the superscripts to make sure
what we still need to go on:
2 + 2 + 6 + 2 + 6 = 18 electrons
Remember, Manganese (Mn) has 25 electrons,
so we need to go on.
7 - 36
12
An Example of Electron Configuration
1s22s22p63s23p6
Since there are no more elements in Period 3, we
go to Period 4, where on the left, we are again in
the s block, and there are 2 elements in the s
block in Period 4.
Moving right, we come across the d block, which
is 3d (period # - 1), and there are 5 elements in
3d, including Mn.
1s22s22p63s23p64s23d5
7 - 37
An Example of Electron
Configuration
7 - 38
Practice – Electron Configurations
• Write electron configurations for the
following:
1. Al
2. Sc
3. K
4. Br
5. Zn
6. Hg
7 - 39
13
Practice Solutions – Electron
Configurations
• Write electron configurations for the
following:
1. Al – 1s22s22p63s23p1
2. Sc – 1s22s22p63s23p64s23d1
3. K – 1s22s22p63s23p64s1
4. Br – 1s22s22p63s23p64s23d104p5
5. Zn – 1s22s22p63s23p64s23d10
6. Hg –
1s22s22p63s23p64s23d104p65s24d105p66s2
4f145d10
7 - 40
Valence Electrons for Main-Group
Elements
• Valence level (shell)
– Last-filled principal energy level
– Highest energy level
– Contains orbitals that are larger than
orbitals in lower energy levels
• Valence electron
– An electron that occupies the valence level
– Elements in the same group have the same
number of valence electrons
– Example: Br – 1s22s22p63s23p64s23d104p5
7 - 41
Valence Electrons for Main-Group
Elements
• Valence electrons
– The Roman numeral group number
(which is paired with an A or B).
– The number of valence electrons is also
equal to the number of s and p electrons
in the valence level for any main-group
element.
• Core electron
– An electron in a principal energy level
below the valence level
– Inner electron
7 - 42
14
Valence Electrons for Main-Group
Elements
Figure 7.23
7 - 43
Practice – Valence Electrons
• For each of the following, determine
the number of valence electrons.
1. Magnesium
2. Carbon
3. Boron
4. Chlorine
5. Selenium
7 - 44
Practice Solutions – Valence
Electrons
• For each of the following, determine the
number of valence electrons.
1. Magnesium – In group IIA (or 2): 2 valence
electrons
2. Carbon – In group IVA (or 14): 4 valence
electrons
3. Boron – In group IIIA (or 13): 3 valence
electrons
4. Chlorine – In group VIIA (or 17): 7 valence
electrons
5. Selenium – In group VIA (or 16): 6 valence
electrons
7 - 45
15
Abbreviated Electron
Configuration
•
•
•
Starts the electron configuration at the last noble gas
before the element in question.
Write down the noble gas in brackets, then fill in the
rest of the electron configuration until you reach the
element in question.
Example:
Phosphorus (P) has 15 electrons.
In the long notation, the electron configuration for P
would be:
1s22s22p63s23p3
The abbreviated electron configuration for P would be:
[Ne] 3s23p3
7 - 46
Practice – Abbreviated Electron
Configuration
• Write the abbreviated electron
configuration for the following:
1. Magnesium
2. Carbon
3. Boron
4. Chlorine
5. Selenium
7 - 47
Practice Solutions – Abbreviated
Electron Configuration
• Write the abbreviated electron
configuration for the following:
1. Magnesium – [Ne] 3s2
2. Carbon – [He] 2s22p2
3. Boron – [He] 2s22p1
4. Chlorine – [Ne] 3s23p5
5. Selenium – [Ar] 4s23d104p4
7 - 48
16
Electron Configurations for Ions
•
Ions form because atoms gain or lose
electrons
– Cations
• Positively charged ions
• Subtract the number of the charge from
the total number of electrons
• Move to the left the number of spaces
equal to the charge on the periodic table
– Anions
• Negatively charged ion
• Add the number of the charge to the total
number of electrons
• Move to the right the number of spaces
equal to the charge on the periodic table
7 - 49
Electron Configurations for Ions
•
Isoelectronic
– Have the same number of electrons
– Ions and elements with the same electron
configuration
– Example: Are the bromine ion and strontium ion
isoelectronic?
Bromine forms an ion with a -1 charge, while
strontium form an ion with a +2 charge. Moving the
correct spaces on the periodic table, we find that
bromine’s ion and strontium’s ion share an electron
configuration with krypton, and are therefore
isoelectronic.
Br- – 1s22s22p63s23p64s23d104p6
Sr2+ - 1s22s22p63s23p64s23d104p6
Kr - 1s22s22p63s23p64s23d104p6
7 - 50
Electron Configurations for Ions
• An example problem:
– Write the electron configuration for
Na+:
Na+ has a positive charge of 1;
therefore, we need to subtract 1
electron from the total number of
electrons, 11. Na+ has 10 electrons
and is isoelectronic with Ne.
1s22s22p6
7 - 51
17
Practice – Electron Configurations
for Ions
• Write the electron configuration in
long and abbreviated notation for
the following ions.
1. Br2. N33. K+
4. Sr2+
5. S26. Ni2+
7 - 52
•
Practice Solutions – Electron
Configurations for Ions
Write the electron configuration in long and
abbreviated notation for the following ions.
1. Br- – 1s22s22p63s23p64s23d104p6 [Kr]
isoelectronic with Kr
2. N3- - 1s22s22p6
[Ne]
isoelectronic with Ne
3. K+ - 1s22s22p63s23p6
[Ar]
isoelectronic with Ar
4. Sr2+ - 1s22s22p63s23p64s23d104p6
[Kr]
isoelectronic with Kr and Br-1
5. S2- - 1s22s22p63s23p6
[Ar]
isoelectronic with Ar and K+1
6. Ni2+ - 1s22s22p63s23p6
[Ar]4s23d6
isoelectronic with Fe
7 - 53
Periodic Trends
•
•
Valence electrons are the electrons that
participate in chemical reactions because they
are the farthest electrons from the nucleus.
Because elements in the same group have the
same number of valence electrons, elements in
the same group have very similar reactivities.
Figure 7.25
7 - 54
18
Ionization Energy
• Ionization Energy
– A measure of the
energy required to
remove a valence
electron from a
gaseous atom to form
a gaseous ion.
– In general, atoms with
low ionization energies
do not bind their
electrons very tightly,
and are therefore, very
reactive.
Figure 7.26
7 - 55
Trends in Ionization Energy
• The general
trend for
ionization
energy is for
ionization
energy to
increase
from bottom
to top and
from left to
right across
the periodic
table.
Figure 7.27
7 - 56
Successive Ionization Energies
• The trend describes the first ionization
energy (IE1), or the amount of energy it
takes to remove 1 electron from an atom.
• The amount of energy it takes to remove
a 2nd electron is known as the second
ionization energy (IE2) and is larger than
IE1.
• In general,
IE3 > IE2 > IE1
7 - 57
19
Successive Ionization Energies
7 - 58
Atomic Size
• Atomic size is often
described in terms
of atomic radius.
Figure 7.29
– Atomic radius is the
distance from the
center of the nucleus
to the outer edge of
the atom.
7 - 59
Trends in Atomic Size
7 - 60
20
Trends in Atomic Size
•
The general trend for atomic size (or radius) is
for atomic size to increase from top to bottom
and from right to left across the periodic table.
Figure 7.31
7 - 61
Ionic Size
• Ionic size
– Radius of an
ion
– Atoms
change
radius when
they become
ions
7 - 62
Ionic Size
• For any
isoelectronic
series, as the
number of
protons
increases,
the ion size
decreases.
7 - 63
21
Trends in Ionic Size
•
The general trend for ionic size (or radius) is for ionic
size to increase from top to bottom.
– For cations, as the charge increases, the ionic size
decreases.
– For anions, as the charge increases in negativity, the
ionic size increases.
7 - 64
22